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Reactions of metal-ammonia solutions — VI The alkali metal-ammonia reaction as catalyzed by glass
J. inorg,nucl.Chem.,1968,Vol.30. pp. 2047to 2057. PergamonPress. Printedin Great Britain
REACTIONS OF METAL-AMMONIA SOLUTIONS-VI THE ALKALI METAL-AMMONIA REACTION AS CATALYZED BY GLASS*? D O N A L D C. J A C K M A N and C. W. K E E N A N Department of Chemistry, The University of Tennessee, Knoxville, Tennessee 37916 (Received 5 January 1968) A b s t r a c t - T h e room-temperature alkali metal-liquid ammonia reaction, for lithium, potassium and cesium, is shown to be catalyzed by Pyrex or Kimax glass. The catalytic activity, which is probably due to silanol groups, can be greatly decreased by treating the glass with ammonium fluoride or baking with hot potassium vapour. The lithium and potassium reactions are first-order with respect to metal concentration; the overall rates decrease in the order K > Cs > Li. The half-life of the potassium reaction is several hundred hours in specially treated cells, which makes it the most stable such solution at room temperature reported to date.
ALKALI metal-liquid ammonia solutions have been of interest to chemists for many years, but only recently [1] have rates of decomposition of these solutions been determined quantitatively, and then only for the potassium-ammonia solution. Because a study [2] of the reactions of lithium, sodium, and potassium with ethanol in liquid ammonia indicated that different alkali metals react with ammonia at different rates, we have investigated the rate of the metal-ammonia reaction in glass at room temperature for lithium, potassium, and cesium. That impurities, particularly transition metals, and water physically adsorbed on the glass walls decrease the stability of metal-ammonia solutions is well known [3]. Although relatively stable solutions have been prepared at low temperatures with carefully purified reagents, it is only since the room temperature work described in Reference [ 1] that investigators have suggested that the homogeneous reaction to yield amide ion and hydrogen may not occur at a measurable rate [4]. One of *For paper V see J. Am. Chem. Soc. 8"/, 5799 (1965). ?This paper is based on the dissertation by D. C. J. presented to the graduate school of the University of Tennessee in August, 1966, in partial fulfillment of the P h . D . Presented at the 153rd National Meeting of the American Chemical Society, Miami, Florida, April, 1967. 1. (a) D. Y. P. Chou, M. J. Pribble, D. C. Jackman and C. W. Keenan, J. Am. chem. Soc. 85, 3530 (1963); (b) J. Corset and G. Lepoutre, Metal Ammonia Solutions, (Edited by G. Lepoutre and M. J. Sienko), p. 186. Benjamin, New York (1964); (c) P. Pajot, A. Demotier and G. Lepoutre, ibid. p. 206. 2. E.J. Kelly, H. V. Secor, C. W. Keenan and J. F. Eastham, J. A m. chem. Soc. 84, 3611 (1962). 3. (a) W. M. Burgess and H. L. Kahler, J. Am. chem. Soc. 60, 189 (1938); (b) G. W. Watt, G. D. Barnett and L. Vaska, Ind. Engng Chem. 46, 1022 (1954); (c) I. Warshawsky, J. inorg, nucl. Chem. 25, 601 (1963); (d) I. Warshawsky, J. Catalysis 3, 291 (1964); (e) It should be noted that the reaction reported on in this paper is not the principal reaction studied by Warshawsky, which was the reaction of metal and metal-ammonia solutions with water adsorbed on glass. 4. (a) W. L. Jolly and C. J. Hallada, Non-Aqueous Solvent Systems, (Edited by T. C. Waddington). p. 38 Academic Press, New York (1965): Ib) D. F. Burow and J. J. Lagowski,Adv. Chem. Ser. 50, 131 (1965). 2047
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D . C . J A C K M A N and C. W. K E E N A N
t h e a i m s o f t h i s w o r k w a s to t e s t t h i s h y p o t h e s i s u n d e r t h e s o m e w h a t e x t r e m e condition of room temperature. T o s t u d y t h e r e a c t i o n s o f alkali m e t a l s w i t h l i q u i d a m m o n i a at r o o m t e m p e r a ture, we used a Pyrex or Kimax reaction vessel that could be rigorously cleaned and dried, and that provided for the separation of the reactants until they could b e r a p i d l y m i x e d b y s h a t t e r i n g a g l a s s d i a p h r a g m . T h i s p a p e r r e p o r t s (i) t h e r a t e s of reaction of lithium, potassium and cesium with liquid ammonia at room t e m p e r a t u r e in P y r e x o r K i m a x r e a c t i o n v e s s e l s t r e a t e d t o r e m o v e s u r f a c e i m p u r i t i e s a n d p h y s i c a l l y a d s o r b e d w a t e r , (ii) t h a t t h e p o t a s s i u m - a m m o n i a r e a c t i o n is c a t a l y z e d b y g l a s s , a n d (iii) t h a t t h e g l a s s s u r f a c e m a y b e t r e a t e d t o d e c r e a s e its catalytic activity. EXPERIMENTAL Reagents. High purity potassium and cesium were obtained in ampoules from Americal Potash and Chemical Corporation. The metals were further purified by distilling in vacuo in baked-out, all glass systems, into capillary tubes which were then sealed. The lithium was obtained from Oak Ridge National Laboratory[5]. Because lithium can not be distilled in glass, it was handled in an argonfilled glove box[6]; samples were gouged out of the solid metal with melting point tubes and transferred in these tubes to the reaction cell. Oxides of the metals were semi-quantitatively analyzed [7] spectroscopically with a d-c arc and Stallwood jet. Upper limits in p.p.m, found for impurities were as follows: Cesium: AI, 25; Ca, 25; and Si, 10. Lithium: AI, 10; Ca, 100; Cu, 1; Mg, 25; Si, 50; and Sr, 50. Potassium: AI, 25; Ca, 50: Na, 550, and Si, 25. The absence of transition elements is noteworthy. None of the metals showed any sign of reacting in the sealed cells, although the cesium had a slight golden colour at all stages. This is the colour of the pure metal according to some authors [8(a)], whereas others state that this colour indicates oxide impurity [8(b)]. The anhydrous ammonia (purchased from either E. I. duPont or Olin Matheson) was of 99.99+ per cent purity. Individual samples of ammonia were doubly distilled from sodium-ammonia solutions just prior to use. From a sample of ammonia similarly purified an nmr spectrum with three sharp peaks was obtained. As interpreted by Ogg [9], such a spectrum is evidence of super-dry ammonia containing less than 0.1 p.p.m, of water. Apparatus. A sealed reaction cell is shown in its steel protective jacket J in Fig. 1. The over-all length of the glass cell is about 20 cm. For loading with reagents, the cell is initially connected to vacuum manifolds at A and E. Alkali metal is put into the compartment above the diaphragm D, which also holds the heavy glass breaker B; potassium or cesium is introduced by distillation, or lithium in an open capillary under an argon atmosphere, then the compartment is sealed in vacuo at A. Liquid ammonia is held in the lower compartment, the bottom of which is a Pyrex optical cuvette C which has a one cm light path (Beckman 75152). The bodies of the cells, other than the cuvette, were of either Pyrex or Kimax glass; no differences in the reaction were observed due to such a difference in the
glass. For introducing ammonia, the cell is turned so the cuvette is uppermost, then cooled in a dry ice-alcohol bath in order to distill in ammonia via E from a graduated reservoir containing a little sodium. The volume of liquid at --78° is measured; the cell is sealed at E, put into the steel jacket, and allowed to warm to room temperature, 23° _+ I °. A counter pressure of about 140 psi of nitrogen was sometimes applied via the needle valve N; alternatively a little liquid ammonia on tissue paper was put in the jacket as it was assembled. Heavy silica windows W provided for the spectroscopic monitoring of the cuvette. 5. The authors are indebted to Mr. Harvey Bronstein, Oak Ridge National Laboratory, for a sample of specially purified lithium-7. 6. J. L. Bloomer and C. W. Keenan, U. S. Patent No. 2,862,307, December 2, 1958. 7. The authors are grateful to Dr. I. H. Tipton and Miss Peggy Stewart of the University of Tennessee Physics Department for these analyses. 8. (a) F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, p. 316. lnterscience, New York (1962); (b) J. J. Kennedy, Chem. Rev. 23, 157 (1938). 9. R.A. Ogg andJ. D. Ray, J. chem. Phys. 26, 1515 (1957).
Reactions of metal-ammonia solutions - V I
I
2049
I
~
'--N
A~
~D
/ - -
.
// W Fig. 1. Two compartment glass reaction cell inside steel protective jacket. Analyses were recorded with a Cary Model 14 spectrophotometer, with a sealed cuvette filled with ammonia in the reference beam. The base line absorption for ammonia was obtained, then the celljacket assembly was shaken to shatter the diaphragm [ 10] and to mix the reactants. As the assembly was returned to the vertical, the ring seal at R prevented diaphragm fragments from falling into the cuvette. Within about a minute after mixing, the spectrum of the solution was recorded from 70003000 ,~. Tumble mixing of cell contents and rescanning was then repeated periodically. Whether due to the initial fracture or to tumbling with the heavy glass breaker during mixing, the diaphragm fragments were usually reduced almost to the state of ground glass. Many runs were carried out in cells in which the initial positions of the metal and ammonia were reversed, with the diaphragm being near the cuvette; in these cases the base line absorption for ammonia at 6500/~ was determined after disappearance of the blue colour. Runs which involved special treatment of the glass (See next section) were carried out in cells without diaphragms; these reactions were relatively slow, so initial readings were of less interest. Ammonia was doubly distilled directly at -78 ° into cells which contained metal, and the first spectra were recorded after about two-hours, when the cell had warmed to room temperature. Cell treatments. Standard procedure was to treat cells with aqua regia for 2 hr, rinse copiously with doubly distilled water, and bake in vacuo at > 400° and < 10-5 mm for 24 hr prior to loading with reactants. Some reaction cells were treated with a 30 per cent ammonium fluoride solution, following the
10. The breakable glass diaphragms were designed and manufactured by Labglass of Tennessee, Kingsport, Tennessee. The diaphragm must withstand a pressure differential of about 10 arm, but be very thin in the centre so that when it is broken lateral cracks do not spread toward the outer wall of the cell. The glass cells usually could withstand the vapour pressure of ammonia even without the external counter pressure. Incidentally, we never observed a cell failure which could be traced to the fracture of the planar faces of the Pyrex cuvette (in contrast to the report of Reference [l(b)]. On the other hand, our attempts to use fused silica cells have been thwarted by the failure of rectangular cuvettes.
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and C. W. K E E N A N
procedure described by Elmer, C h a p m a n , and N o r d b e r g for r e p l a c i n g - - - O H groups on silica[l 1], and then b a k e d in vacuo at > 500 ° for 24 hr. T w o p o t a s s i u m reactions were c a r d e d out in cells which were first subjected to the standard procedure a n d then baked at about 400 ° with p o t a s s i u m vapour. In the first instance the glass was baked until it was a light yellow. In the second instance, the baking was continued until the glass was a dark amber, t h e n the cuvette alone was heated to about 500 ° to drive s o m e of the p o t a s s i u m out of the optical P y r e x windows. Analysis o f metal concentration. A t the end of each reaction, except for t h o s e b a k e d with potassium vapour, the cells were broken open, the a m m o n i a allowed to evaporate, and the residue was dissolved in very dilute HC1 and analyzed for alkali metal with a flame photometer. F r o m the m e a s ured v o l u m e of a m m o n i a in the reaction cells at --78 ° a n d the k n o w n densities of a m m o n i a at - 7 8 ° and 23 °, the concentrations of the solutions at 23 ° were calculated. The spectra. F o u r spectra for a typical p o t a s s i u m - a m m o n i a reaction are s h o w n in Fig. 2. F o r all kinetic studies, the a b s o r b a n c e A at 6500 ,~ w a s taken as a m e a s u r e of unreacted metal present, that is, the concentration o f solvated electrons. This wavelength was selected b e c a u s e the samples c h o s e n
2.0
1.6 z o
1.2
TIME
ZERO
0.8
0.4
I 3500
I
I 4500
I
I 5500
51 HR I
I 6500
WAVELENGTH (~) Fig. 2. Spectra taken at four different times during a potassium-liquid a m m o n i a reaction at 23 ± 1°. for study, about 1 × 10-3 M initially and of 1 cm light path, had a b s o r b a n c e s less t h a n 2 at 6500 A. F o r the p o t a s s i u m s y s t e m , the a b s o r b a n c e due to the amide ion at a n y time was calculated as follows: [ A 6500A.t ] A NHi (3520,~),t = A 3520,g, .t- A 3520/~,t=0 X A e,500/~,t=~,J where A:~sz.A.t_o and A6,~ooA.r-oare the a b s o r b a n c e readings due to metal at those w a v e l e n g t h s before any reaction had t a k e n place, A.~zo&~ a n d A6nooAa are the absorbance readings at any time t, a n d ANn~:~32o~,tis the calculated a b s o r b a n c e due to amide alone at time t [12]. A similar calculation was m a d e 11. T. H. Elmer, I. D. C h a p m a n and M. E. Nordberg, J. phys. Chem. 67, 2219 (1963). 12. T h e value of 3520,~ for the amide peak at 23 ° is to be compared with an expected value o f 3510/~ calculated from the value of 3350,~ at - 4 8 ° reported by M. Ottolenglhi and H. Linschitz, Adv. Chem. Ser. 50, 149 (1965).
Reactions of metal-ammonia solutions - V 1
2051
for the lithium system, but for cesium a further correction had to be made because even the initial spectrum for each cesium run showed a small but definite amide peak (as well or better developed than the 40 min (potassium) amide peak in Fig. 2). The initial absorption due to (cesium) amide, estimated by measuring the height of the peak above the extrapolated metal curve, was subtracted from all amide peak readings before making calculations and plots for the cesium reaction. The initial amide found in the cesium solutions could have been formed by rapid initial reaction or might have resulted from the presence of a trace of cesium oxide in the distilled metal. The amide peak occurred at 3240 ,~, in the lithium solution and at 3600 ,~ in the cesium solution. In all three systems, the threshold for amide absorption was far below 6500 ,~. Isosbestic points at about 4000/~ were observed for either potassium or cesium spectral curves and at about 3650 .~ for the lithium curves. RESULTS AND DISCUSSION
The reaction being followed. The reaction of interest is eh-m+NH3 ~ NH~-+ 1/2H2.
(1)
The appearance of the isosbestic points indicates that only two species, e;m and NH~, are absorbing in the spectral region monitored. As a check on this conclusion, A65ook,t was plotted vs. ANHr.tfor all reactions (as in Fig. 3) and a direct proportionality found in all cases, although this behaviour was limited in the case '
I
'
I
'
I
'
I
'
1.6
-
o~ 1.2 0
z
7
0.8
en
q
0o')
(
(
0.4
0
0
t
I 0.2
t
I 0.4
t
I
,
0.6
I 0.8
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ABSORBANCE OF NH2"Fig. 3. Decrease in e~-mversus increase in NH~- concentration for reactions of three alkali metals with liquid ammonia.
of lithium solutions by the slight solubility of lithium amide. T h e direct proportionality was expected, because both sodium solutions [ 1 3 (a)] and amide solutions [1 3(b)] have been s h o w n to follow Beer's law. 13. (a) W. L. Jolly, C. J. Hallada, and M. Gold, Metal-Ammonia Solutions, (Edited by G. Lepoutre and M. J. Sienko) p. 174. Benjamin, New York (1964); (b)J. J. Lagowski, Private communication.
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D . C . J A C K M A N and C. W. K E E N A N
It is noted that reaction (1) has been shown to be reversed measurably by increasing the hydrogen pressure[14]. However, at the hydrogen pressures developed in our cells, estimated as about 0.01 atm, the calculated equilibrium concentration of e~m is negligible for our concentration ranges. The short vertical segment at the beginning of the lithium run plotted here was not observed for other lithium runs. It indicates that in this case some of the lithium was used initially in a reaction which produced no amide ion. (Two unusual sources of contamination were present in the lithium experiments: The metal-containing compartment of the cell was exposed in the gloved box to argon after being baked out, and the fragment of melting point tube used to transfer the lithium sample into the cell was not baked.) The latter part of the lithium curve reveals that after a slight degree of supersaturation, lithium amide precipitates. Shortly after the spectrophotometer indicated the NH~ concentration was remaining constant, a thin transparent film was seen to be forming on the walls of the cuvette. The deviation of the lithium plot from linearity just prior to precipitation may be due to ion-pairing. The dotted extrapolated point above the lithium curve describes a saturated solution at 23 °. The estimated concentration of the NH~- in the saturated solution is about 2.6 × 10-4 M; the total Li ÷ concentration for this reaction is 1.2 × 10-3 M, assuming complete dissociation of the metal into ions and solvated electrons. For runs Li-2 and Li-3, the estimated concentrations of NH~- in the saturated solutions were 2.9 and 3.0 × 10-4 M, respectively. Rates of reactions in standard cells. The potassium and lithium reactions are first-order with respect to the concentration of unreacted metal, as shown by the linearity of logA 6500~,vs. time (as in Fig. 4). The cesium reaction is not first-order until about half the cesium has reacted, then its plot becomes linear, also. Data for runs in cells with broken diaphragms are listed in Table 1. The three curves shown in Fig. 4 are representative. In the case of the potassium reactions no Table 1 Run K-1 I(-2 K-3 K-4 K-5 Li- 1 Li-2 Li-3 Cs-1 Cs-2 Cs-3
mg metal
Initial molarity metal × 10-4
ttn, hr*
Molar absorbancy at NH~- peakt
1"08 0"55 0.29 1"64 0"32 0-18 0"49 0"38 1.82 2.30 1.65
7"5 3"7 2"0 11"3 2"4 12'0 17'0 14.0 4.0 5.5 3.4
10 13 9 11 8 140 48 72 44 36 36
3500 2700 2700 3500 2800 2400 1400 -2580 2610 2810
*Half-times for the lithium reactions are for the reaction after LiNH~ began to precipitate; for the cesium reactions, half-times are for the latter part of the reaction. tMolar absorbancy indices were calculated using concentrations based on weight of metal found by flame photometric analysis. 14. E.J. Kirschke and W. L. Jolly, Science 147, 45 (1965).
Reactions of m e t a l - a m m o n i a solutions - V i TIME
2.0
'
50 I
t HR I
I00 I
'
'
2053
Li
150 I
'
200 I
'
L.i-I 1.0
0.7 I-
W U
_
0.5
z
m n- 0 . 3 0 v) a3 <: 0 . 2
0.1
I
I I0
, TIME,
I 20 HR,
I K
I 30
I
AND
C=
I 40
Fig. 4. Log concentration e~m vs. time for reactions of three alkali metals with liquid a m m o n i a at 23 --- 1°.
deviation from linearity near time zero was observed for any reaction; this contrasts with the data of Corset and Lepoutre [l(b)]. In the case of lithium, two linear sections intersect about 60 hr after initiation; the change to the slower rate occurs at about the time LiNHz begins to precipitate. First-order plots for potassium were linear through about 85 per cent, for lithium through 70 per cent reaction. Although the cesium reaction is relatively rapid at first and is not first-order, there is little doubt that reaction (1) is indeed occurring, because of the linearity of the plot in Fig. 3. This kinetic behaviour of the cesium reaction is at variance with the observation that the cesium-ammonia reaction is catalyzed by the presence of cesium amide[15]. It has also been reported that significant decomposition of metal-ammonia solutions occurs in solutions of strongly basic species [4(b)]. It seems clear that, for our reactions, amide in solution is not a catalyst; indeed precipitation of lithium amide actually decreases the rate of the lithium reaction, probably due to the precipitation of amide on the wall of the reaction vessel. Rate o f reactions in special cells. Because the potassium reaction is straightforward first-order and fairly reproducible, it was used to evaluate the effect of variables in cell treatment. Results of several tests are shown in Table 2. To measure the effect on the rate of a change in glass surface area, cells without diaphragms were constructed with perforated glass plates dividing the cells approximately in half. A cuvette was attached to the half below the plate; the upper half was packed with fire-polished lengths of 8 mm Pyrex tubing. After standard cleaning and bake-out, potassium was distilled into the cuvette and 15. J. W. Hodgins, Can. J. Research, 27B, 861 (1949).
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a n d C. W. K E E N A N Table 2
Run
Initial moles K/1
t~2, hr
Comments
K-6
1.04 x 10-~
Ratio of areas, 3.3; ratio of rate constants, 4.4
K-7
0-96 x 10-3
Cuvette end, 19.3 Packed end, 4.4 Cuvette end, 23.9 Packed end, 5'9 370 380 620 4300 (?)
K-8 K-9 K- I 0 K- I 1
3-3 × 6.0 × 1.34 × 1.3 ×
10-4 10-4 10-a 10-4
Ratio of areas, 3.6; ratio of rate constants, 4.0
Fluoride treated Fluoride treated K vapour bake-out K vapour bake-out
ammonia was condensed directly on the potassium at -78°; the cell was warmed in about two hours to room temperature, and then kinetic data were collected. The rate of the potassium-ammonia reaction was determined first in the cuvette end for about 3 hr, then the cell was inverted and the rate determined in the packed end. In the latter case, the solution was poured into the cuvette periodicaUy for spectral monitoring. As the data for runs K-6 and K-7 in Table 2 show, the increase in the rate of the reaction in the packed end is roughly proportional to the increase in the area of the glass in contact with the solution. The longer half-lives in the cuvette ends of these two cells as compared to those for the potassium runs in Table 1 are presumably accounted for by the absence of diaphragm fragments in the two solutions described in Table 2. The half-times of runs K-6 and K-7 in the cuvette e'ffds agree fairly well with the results reported by others [l(b), (c)]. Dr. M. L. Hair called our attention to the catalytic role of silanol groups on glass surfaces and to the presence of both Bronsted and Lewis acid sites on Pyrex, e.g., silanol groups and boron atoms. T o determine the effect of replacing silanol ----OH with - - F , by the technique mentioned under Cell Treatments, two runs were made, K-8 and K-9 in Table 2. These cells had no diaphragms, so the first kinetic readings were made about two hours after the cells were sealed, when the solutions had warmed to room temperature. The effect of the fluoride treatment can be seen by comparing t1/2 for K-8 and K-9 with the cuvette-end times of K-6 and K-7. A bake-out temperature of just over 500 ° in v a c u o after ammonium fluoride treatment does not effect a complete replacement o f - - O H by - - F on silica [ 11 ], but this was the maximum temperature we felt we could use without mechanical damage to our cells. Baking out at about 500 ° probably does eliminate adjacent ----OH groups on silica[11]; this would decrease the likelihood of hydrogen bonding between two silanol groups, a process which is attractive in providing a mechanism for the production of H~ molecules as a result of e~m attack on adsorbed NH3 molecules. In the case of the potassium reaction, the fluoride treatment of the glass presumably decreases the rate of the reaction because the number of catalytic ---OH groups is decreased. Further work is necessary to permit even this simple
Reactions of metal-ammonia solutions - V!
2055
conclusion, for preliminary results show that in its first-order re#on the cesium reaction is a little faster in fluoride-treated cells[16]. It has been found that partial replacement of ----OH by - - F enhances the acidity of the remaining ---OH groups [ 17]; possibly this effect is more important for the cesium reaction than for the potassium. As mentioned under Cell treatments, two reactions were carried out in cells which were baked out with potassium vapour; the aim was to remove all Bronsted acid sites. The results of two runs, K-10 and K-11, are given in Table 2. These cells had no diaphragms; metal and then ammonia were distilled into the baked out cell via break-seal connections. The tl/2 for K-11 is only approximate, because the reaction was followed for only 160 hr. Miscellaneous observations. The suggestion was made in the Discussion Section of Ref. [l(c)] that the rate of the metal-ammonia reaction might be sensitive to stirring, if the decomposition occurs at the wall. To test the effect of stirring, the cell in Run K-9 was agitated vigorously with a wrist-action shaker after the rate of the reaction, as shown in Table 2, had been determined in the unshaken cell. The ratio of the first-order rate constants for the shaken vs. the unshaken cell was 1.62, but the cell wall area bathed by the solution for the shaken vs. unshaken cell was estimated as 1-6, so it appears that the agitation per se resulted in no increase in the rate of the reaction. The magnitude of the activation energy of the metal-ammonia reaction is consistent with that of a diffusion-controlled reaction[l(c)], so it is of interest to consider such a possibility. Jolly has claimed [18] that solvated electrons in ammonia have an abnormally high rate of diffusion, his view evidently being based on the fact that the anionic species in a metal solution does have a large transference number. We hold, however, that the rate of thermal diffusion of the electron cannot be abnormally high, because of the electroneutrality limitation. In the case of concentration differences in a non-homogeneous metal-ammonia solution, solvated electrons cannot diffuse to establish homogeneity any faster than the metal ions which must accompany them. In the case of solutions such as we are reporting on, which can be assumed to be homogeneous with respect to metal ions, the diffusion rate of the electrons toward the wall will be limited by the rate at which amide ions diffuse away from the wall. It is agreed that electrons can migrate without carrying ammonia molecules with them, but we feel that in the case of thermal diffusion the electron must come within the coulombic field of a positive charge, that is, within a few angstroms, in order that the positive charge be attacked. We have relied on a similar argument elsewhere[19], and feel that Jolly's view is invalid. CONCLUSIONS
At room temperature, 10-4 to 10 -3 M potassium reacts with ammonia in a reaction which is first-order with respect to solvated electrons and which is catalyzed by borosilicate glass. Probably silanol ----OH groups are the main catalytic sites, although Lewis acid sites may be active also. The adsorption by 16. 17. 18. 19.
Studies in progress by Joseph Knauer in this laboratory. I. D. Chapman and M. L. Hair, J. Catalysis 2, 145 (1963). W. L. Jolly, A dr. Chem. Ser. 50, 33 (1965). C.W. Keenan, H. V. Secor, E. J. Kelly and J. F. Eastham, J. A m. chem. Soc. 82, 1831 (1960).
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D.C.
J A C K M A N and C. W. K E E N A N
glass of ammonia on both silanol groups and on non-Bronsted sites, possibly boron atoms, has been described [ 17]. Baking at >400°C and < 10-s mm for 24 hr removes adsorbed water so effectively that our spectroscopic analyses show that the amide produced is equivalent to the initial metal concentration. The results of Lepoutre and coworkers [ 1(b),(c)] indicate that if a little water does remain on the wall (as revealed by the initial curvature of their first-order plots), that potassium solutions of stability comparable to ours result after this water reacts. This is essentially the method of cell treatment reported by Burow and Lagowski[4(b)]; before use, they aged glass vessels 48 hr with 10 -z M potassium solutions. They stated that a 10-5 M solution showed no decomposition in 20 hr. Though their temperature is not given, one of their working temperatures is -70°C; at this temperature the half-life of the reaction in one of our cells without a diaphragm, treated according to our standard procedure, would probably be more than a year, so we would find "no decomposition" in 20 hr, also. They attributed the stability of their solutions to the replacement by ammonia of "adsorbed water and/or hydroxyl species on the surface of the glass . . . . " How ammonia could replace a hydroxyl group is not clear to us; in any event, our results indicate that silanol groups are not removed by dilute metal-ammonia solutions, but indeed remain as catalysts for the metal-ammonia reaction. For the potassium reaction the catalytic activity of the glass can be greatly decreased by treatment with ammonium fluoride prior to baking or by baking with potassium vapour; these findings are consistent with the view that -----OH groups are being removed by these treatments. It is proposed that in our standard cells (Table 1) an ammonia molecule is bound to a surface acid group and that the resulting enhancement of the acidity of the ammonia hydrogens leaves them open to attack by a solvated electron. Possibly attack of adjacent adsorbed NHa molecules facilitates formation of H2. The most effective catalysts for the metal-ammonia reaction are the transition metals. It has been suggested/20] that certain catalytic activity of nickel is due to is possessing "inherent acid properties." Under the conditions described, the over-all rates of metal-ammonia reactions for three metals decrease in the order K > Cs > Li, The lithium reaction is, like the potassium, first-order over the majority of the reaction; there is no evidence that amide produced in solution is a catalyst for the reactions. It seems likely that the lithium reaction would be even slower if the lithium could be transferred into the cell in as clean a state as the other two metals. The very slow rates of the potassium reaction in the cells treated to decrease the catalytic activity of the glass indicate clearly that the homogeneous reaction of solvated electrons with ammonia may be immeasurably slow even at room temperature. We are now trying other treatments of the glass, e.g., baking with chloromethylsilanes[21], and we plan to make cells of other materials. Further 20. H. Pines, M. Shamaiengar, and W. S. Postl, J. Am. chem. Soc. 77, 5099 (1955). 21. We are indebted to Dr. M. L. Hair for this suggestion. See, for example, T. E. White, Soc. Plastics Ind., Proc. 20, 3-B, I (1965). Studies in progress by W. E. Moddeman in this laboratory indicate that baking with dichlorodimethylsilane is very effective in decreasing the catalytic activity of the glass.
Reactions of metal-ammonia solutions - V 1
2057
improvement should make possible long-term experiments at room temperature on the same stable solution. Acknowledgements--The
authors are indebted to Drs. James L. Dye, M. L. Hair, and M. H. Lietzke for advice on several points.
REACTIONS OF METAL-AMMONIA SOLUTIONS-VI THE ALKALI METAL-AMMONIA REACTION AS CATALYZED BY GLASS*? D O N A L D C. J A C K M A N and C. W. K E E N A N Department of Chemistry, The University of Tennessee, Knoxville, Tennessee 37916 (Received 5 January 1968) A b s t r a c t - T h e room-temperature alkali metal-liquid ammonia reaction, for lithium, potassium and cesium, is shown to be catalyzed by Pyrex or Kimax glass. The catalytic activity, which is probably due to silanol groups, can be greatly decreased by treating the glass with ammonium fluoride or baking with hot potassium vapour. The lithium and potassium reactions are first-order with respect to metal concentration; the overall rates decrease in the order K > Cs > Li. The half-life of the potassium reaction is several hundred hours in specially treated cells, which makes it the most stable such solution at room temperature reported to date.
ALKALI metal-liquid ammonia solutions have been of interest to chemists for many years, but only recently [1] have rates of decomposition of these solutions been determined quantitatively, and then only for the potassium-ammonia solution. Because a study [2] of the reactions of lithium, sodium, and potassium with ethanol in liquid ammonia indicated that different alkali metals react with ammonia at different rates, we have investigated the rate of the metal-ammonia reaction in glass at room temperature for lithium, potassium, and cesium. That impurities, particularly transition metals, and water physically adsorbed on the glass walls decrease the stability of metal-ammonia solutions is well known [3]. Although relatively stable solutions have been prepared at low temperatures with carefully purified reagents, it is only since the room temperature work described in Reference [ 1] that investigators have suggested that the homogeneous reaction to yield amide ion and hydrogen may not occur at a measurable rate [4]. One of *For paper V see J. Am. Chem. Soc. 8"/, 5799 (1965). ?This paper is based on the dissertation by D. C. J. presented to the graduate school of the University of Tennessee in August, 1966, in partial fulfillment of the P h . D . Presented at the 153rd National Meeting of the American Chemical Society, Miami, Florida, April, 1967. 1. (a) D. Y. P. Chou, M. J. Pribble, D. C. Jackman and C. W. Keenan, J. Am. chem. Soc. 85, 3530 (1963); (b) J. Corset and G. Lepoutre, Metal Ammonia Solutions, (Edited by G. Lepoutre and M. J. Sienko), p. 186. Benjamin, New York (1964); (c) P. Pajot, A. Demotier and G. Lepoutre, ibid. p. 206. 2. E.J. Kelly, H. V. Secor, C. W. Keenan and J. F. Eastham, J. A m. chem. Soc. 84, 3611 (1962). 3. (a) W. M. Burgess and H. L. Kahler, J. Am. chem. Soc. 60, 189 (1938); (b) G. W. Watt, G. D. Barnett and L. Vaska, Ind. Engng Chem. 46, 1022 (1954); (c) I. Warshawsky, J. inorg, nucl. Chem. 25, 601 (1963); (d) I. Warshawsky, J. Catalysis 3, 291 (1964); (e) It should be noted that the reaction reported on in this paper is not the principal reaction studied by Warshawsky, which was the reaction of metal and metal-ammonia solutions with water adsorbed on glass. 4. (a) W. L. Jolly and C. J. Hallada, Non-Aqueous Solvent Systems, (Edited by T. C. Waddington). p. 38 Academic Press, New York (1965): Ib) D. F. Burow and J. J. Lagowski,Adv. Chem. Ser. 50, 131 (1965). 2047
2048
D . C . J A C K M A N and C. W. K E E N A N
t h e a i m s o f t h i s w o r k w a s to t e s t t h i s h y p o t h e s i s u n d e r t h e s o m e w h a t e x t r e m e condition of room temperature. T o s t u d y t h e r e a c t i o n s o f alkali m e t a l s w i t h l i q u i d a m m o n i a at r o o m t e m p e r a ture, we used a Pyrex or Kimax reaction vessel that could be rigorously cleaned and dried, and that provided for the separation of the reactants until they could b e r a p i d l y m i x e d b y s h a t t e r i n g a g l a s s d i a p h r a g m . T h i s p a p e r r e p o r t s (i) t h e r a t e s of reaction of lithium, potassium and cesium with liquid ammonia at room t e m p e r a t u r e in P y r e x o r K i m a x r e a c t i o n v e s s e l s t r e a t e d t o r e m o v e s u r f a c e i m p u r i t i e s a n d p h y s i c a l l y a d s o r b e d w a t e r , (ii) t h a t t h e p o t a s s i u m - a m m o n i a r e a c t i o n is c a t a l y z e d b y g l a s s , a n d (iii) t h a t t h e g l a s s s u r f a c e m a y b e t r e a t e d t o d e c r e a s e its catalytic activity. EXPERIMENTAL Reagents. High purity potassium and cesium were obtained in ampoules from Americal Potash and Chemical Corporation. The metals were further purified by distilling in vacuo in baked-out, all glass systems, into capillary tubes which were then sealed. The lithium was obtained from Oak Ridge National Laboratory[5]. Because lithium can not be distilled in glass, it was handled in an argonfilled glove box[6]; samples were gouged out of the solid metal with melting point tubes and transferred in these tubes to the reaction cell. Oxides of the metals were semi-quantitatively analyzed [7] spectroscopically with a d-c arc and Stallwood jet. Upper limits in p.p.m, found for impurities were as follows: Cesium: AI, 25; Ca, 25; and Si, 10. Lithium: AI, 10; Ca, 100; Cu, 1; Mg, 25; Si, 50; and Sr, 50. Potassium: AI, 25; Ca, 50: Na, 550, and Si, 25. The absence of transition elements is noteworthy. None of the metals showed any sign of reacting in the sealed cells, although the cesium had a slight golden colour at all stages. This is the colour of the pure metal according to some authors [8(a)], whereas others state that this colour indicates oxide impurity [8(b)]. The anhydrous ammonia (purchased from either E. I. duPont or Olin Matheson) was of 99.99+ per cent purity. Individual samples of ammonia were doubly distilled from sodium-ammonia solutions just prior to use. From a sample of ammonia similarly purified an nmr spectrum with three sharp peaks was obtained. As interpreted by Ogg [9], such a spectrum is evidence of super-dry ammonia containing less than 0.1 p.p.m, of water. Apparatus. A sealed reaction cell is shown in its steel protective jacket J in Fig. 1. The over-all length of the glass cell is about 20 cm. For loading with reagents, the cell is initially connected to vacuum manifolds at A and E. Alkali metal is put into the compartment above the diaphragm D, which also holds the heavy glass breaker B; potassium or cesium is introduced by distillation, or lithium in an open capillary under an argon atmosphere, then the compartment is sealed in vacuo at A. Liquid ammonia is held in the lower compartment, the bottom of which is a Pyrex optical cuvette C which has a one cm light path (Beckman 75152). The bodies of the cells, other than the cuvette, were of either Pyrex or Kimax glass; no differences in the reaction were observed due to such a difference in the
glass. For introducing ammonia, the cell is turned so the cuvette is uppermost, then cooled in a dry ice-alcohol bath in order to distill in ammonia via E from a graduated reservoir containing a little sodium. The volume of liquid at --78° is measured; the cell is sealed at E, put into the steel jacket, and allowed to warm to room temperature, 23° _+ I °. A counter pressure of about 140 psi of nitrogen was sometimes applied via the needle valve N; alternatively a little liquid ammonia on tissue paper was put in the jacket as it was assembled. Heavy silica windows W provided for the spectroscopic monitoring of the cuvette. 5. The authors are indebted to Mr. Harvey Bronstein, Oak Ridge National Laboratory, for a sample of specially purified lithium-7. 6. J. L. Bloomer and C. W. Keenan, U. S. Patent No. 2,862,307, December 2, 1958. 7. The authors are grateful to Dr. I. H. Tipton and Miss Peggy Stewart of the University of Tennessee Physics Department for these analyses. 8. (a) F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, p. 316. lnterscience, New York (1962); (b) J. J. Kennedy, Chem. Rev. 23, 157 (1938). 9. R.A. Ogg andJ. D. Ray, J. chem. Phys. 26, 1515 (1957).
Reactions of metal-ammonia solutions - V I
I
2049
I
~
'--N
A~
~D
/ - -
.
// W Fig. 1. Two compartment glass reaction cell inside steel protective jacket. Analyses were recorded with a Cary Model 14 spectrophotometer, with a sealed cuvette filled with ammonia in the reference beam. The base line absorption for ammonia was obtained, then the celljacket assembly was shaken to shatter the diaphragm [ 10] and to mix the reactants. As the assembly was returned to the vertical, the ring seal at R prevented diaphragm fragments from falling into the cuvette. Within about a minute after mixing, the spectrum of the solution was recorded from 70003000 ,~. Tumble mixing of cell contents and rescanning was then repeated periodically. Whether due to the initial fracture or to tumbling with the heavy glass breaker during mixing, the diaphragm fragments were usually reduced almost to the state of ground glass. Many runs were carried out in cells in which the initial positions of the metal and ammonia were reversed, with the diaphragm being near the cuvette; in these cases the base line absorption for ammonia at 6500/~ was determined after disappearance of the blue colour. Runs which involved special treatment of the glass (See next section) were carried out in cells without diaphragms; these reactions were relatively slow, so initial readings were of less interest. Ammonia was doubly distilled directly at -78 ° into cells which contained metal, and the first spectra were recorded after about two-hours, when the cell had warmed to room temperature. Cell treatments. Standard procedure was to treat cells with aqua regia for 2 hr, rinse copiously with doubly distilled water, and bake in vacuo at > 400° and < 10-5 mm for 24 hr prior to loading with reactants. Some reaction cells were treated with a 30 per cent ammonium fluoride solution, following the
10. The breakable glass diaphragms were designed and manufactured by Labglass of Tennessee, Kingsport, Tennessee. The diaphragm must withstand a pressure differential of about 10 arm, but be very thin in the centre so that when it is broken lateral cracks do not spread toward the outer wall of the cell. The glass cells usually could withstand the vapour pressure of ammonia even without the external counter pressure. Incidentally, we never observed a cell failure which could be traced to the fracture of the planar faces of the Pyrex cuvette (in contrast to the report of Reference [l(b)]. On the other hand, our attempts to use fused silica cells have been thwarted by the failure of rectangular cuvettes.
2050
D.C.
JACKMAN
and C. W. K E E N A N
procedure described by Elmer, C h a p m a n , and N o r d b e r g for r e p l a c i n g - - - O H groups on silica[l 1], and then b a k e d in vacuo at > 500 ° for 24 hr. T w o p o t a s s i u m reactions were c a r d e d out in cells which were first subjected to the standard procedure a n d then baked at about 400 ° with p o t a s s i u m vapour. In the first instance the glass was baked until it was a light yellow. In the second instance, the baking was continued until the glass was a dark amber, t h e n the cuvette alone was heated to about 500 ° to drive s o m e of the p o t a s s i u m out of the optical P y r e x windows. Analysis o f metal concentration. A t the end of each reaction, except for t h o s e b a k e d with potassium vapour, the cells were broken open, the a m m o n i a allowed to evaporate, and the residue was dissolved in very dilute HC1 and analyzed for alkali metal with a flame photometer. F r o m the m e a s ured v o l u m e of a m m o n i a in the reaction cells at --78 ° a n d the k n o w n densities of a m m o n i a at - 7 8 ° and 23 °, the concentrations of the solutions at 23 ° were calculated. The spectra. F o u r spectra for a typical p o t a s s i u m - a m m o n i a reaction are s h o w n in Fig. 2. F o r all kinetic studies, the a b s o r b a n c e A at 6500 ,~ w a s taken as a m e a s u r e of unreacted metal present, that is, the concentration o f solvated electrons. This wavelength was selected b e c a u s e the samples c h o s e n
2.0
1.6 z o
1.2
TIME
ZERO
0.8
0.4
I 3500
I
I 4500
I
I 5500
51 HR I
I 6500
WAVELENGTH (~) Fig. 2. Spectra taken at four different times during a potassium-liquid a m m o n i a reaction at 23 ± 1°. for study, about 1 × 10-3 M initially and of 1 cm light path, had a b s o r b a n c e s less t h a n 2 at 6500 A. F o r the p o t a s s i u m s y s t e m , the a b s o r b a n c e due to the amide ion at a n y time was calculated as follows: [ A 6500A.t ] A NHi (3520,~),t = A 3520,g, .t- A 3520/~,t=0 X A e,500/~,t=~,J where A:~sz.A.t_o and A6,~ooA.r-oare the a b s o r b a n c e readings due to metal at those w a v e l e n g t h s before any reaction had t a k e n place, A.~zo&~ a n d A6nooAa are the absorbance readings at any time t, a n d ANn~:~32o~,tis the calculated a b s o r b a n c e due to amide alone at time t [12]. A similar calculation was m a d e 11. T. H. Elmer, I. D. C h a p m a n and M. E. Nordberg, J. phys. Chem. 67, 2219 (1963). 12. T h e value of 3520,~ for the amide peak at 23 ° is to be compared with an expected value o f 3510/~ calculated from the value of 3350,~ at - 4 8 ° reported by M. Ottolenglhi and H. Linschitz, Adv. Chem. Ser. 50, 149 (1965).
Reactions of metal-ammonia solutions - V 1
2051
for the lithium system, but for cesium a further correction had to be made because even the initial spectrum for each cesium run showed a small but definite amide peak (as well or better developed than the 40 min (potassium) amide peak in Fig. 2). The initial absorption due to (cesium) amide, estimated by measuring the height of the peak above the extrapolated metal curve, was subtracted from all amide peak readings before making calculations and plots for the cesium reaction. The initial amide found in the cesium solutions could have been formed by rapid initial reaction or might have resulted from the presence of a trace of cesium oxide in the distilled metal. The amide peak occurred at 3240 ,~, in the lithium solution and at 3600 ,~ in the cesium solution. In all three systems, the threshold for amide absorption was far below 6500 ,~. Isosbestic points at about 4000/~ were observed for either potassium or cesium spectral curves and at about 3650 .~ for the lithium curves. RESULTS AND DISCUSSION
The reaction being followed. The reaction of interest is eh-m+NH3 ~ NH~-+ 1/2H2.
(1)
The appearance of the isosbestic points indicates that only two species, e;m and NH~, are absorbing in the spectral region monitored. As a check on this conclusion, A65ook,t was plotted vs. ANHr.tfor all reactions (as in Fig. 3) and a direct proportionality found in all cases, although this behaviour was limited in the case '
I
'
I
'
I
'
I
'
1.6
-
o~ 1.2 0
z
7
0.8
en
q
0o')
(
(
0.4
0
0
t
I 0.2
t
I 0.4
t
I
,
0.6
I 0.8
I 1.0
ABSORBANCE OF NH2"Fig. 3. Decrease in e~-mversus increase in NH~- concentration for reactions of three alkali metals with liquid ammonia.
of lithium solutions by the slight solubility of lithium amide. T h e direct proportionality was expected, because both sodium solutions [ 1 3 (a)] and amide solutions [1 3(b)] have been s h o w n to follow Beer's law. 13. (a) W. L. Jolly, C. J. Hallada, and M. Gold, Metal-Ammonia Solutions, (Edited by G. Lepoutre and M. J. Sienko) p. 174. Benjamin, New York (1964); (b)J. J. Lagowski, Private communication.
2052
D . C . J A C K M A N and C. W. K E E N A N
It is noted that reaction (1) has been shown to be reversed measurably by increasing the hydrogen pressure[14]. However, at the hydrogen pressures developed in our cells, estimated as about 0.01 atm, the calculated equilibrium concentration of e~m is negligible for our concentration ranges. The short vertical segment at the beginning of the lithium run plotted here was not observed for other lithium runs. It indicates that in this case some of the lithium was used initially in a reaction which produced no amide ion. (Two unusual sources of contamination were present in the lithium experiments: The metal-containing compartment of the cell was exposed in the gloved box to argon after being baked out, and the fragment of melting point tube used to transfer the lithium sample into the cell was not baked.) The latter part of the lithium curve reveals that after a slight degree of supersaturation, lithium amide precipitates. Shortly after the spectrophotometer indicated the NH~ concentration was remaining constant, a thin transparent film was seen to be forming on the walls of the cuvette. The deviation of the lithium plot from linearity just prior to precipitation may be due to ion-pairing. The dotted extrapolated point above the lithium curve describes a saturated solution at 23 °. The estimated concentration of the NH~- in the saturated solution is about 2.6 × 10-4 M; the total Li ÷ concentration for this reaction is 1.2 × 10-3 M, assuming complete dissociation of the metal into ions and solvated electrons. For runs Li-2 and Li-3, the estimated concentrations of NH~- in the saturated solutions were 2.9 and 3.0 × 10-4 M, respectively. Rates of reactions in standard cells. The potassium and lithium reactions are first-order with respect to the concentration of unreacted metal, as shown by the linearity of logA 6500~,vs. time (as in Fig. 4). The cesium reaction is not first-order until about half the cesium has reacted, then its plot becomes linear, also. Data for runs in cells with broken diaphragms are listed in Table 1. The three curves shown in Fig. 4 are representative. In the case of the potassium reactions no Table 1 Run K-1 I(-2 K-3 K-4 K-5 Li- 1 Li-2 Li-3 Cs-1 Cs-2 Cs-3
mg metal
Initial molarity metal × 10-4
ttn, hr*
Molar absorbancy at NH~- peakt
1"08 0"55 0.29 1"64 0"32 0-18 0"49 0"38 1.82 2.30 1.65
7"5 3"7 2"0 11"3 2"4 12'0 17'0 14.0 4.0 5.5 3.4
10 13 9 11 8 140 48 72 44 36 36
3500 2700 2700 3500 2800 2400 1400 -2580 2610 2810
*Half-times for the lithium reactions are for the reaction after LiNH~ began to precipitate; for the cesium reactions, half-times are for the latter part of the reaction. tMolar absorbancy indices were calculated using concentrations based on weight of metal found by flame photometric analysis. 14. E.J. Kirschke and W. L. Jolly, Science 147, 45 (1965).
Reactions of m e t a l - a m m o n i a solutions - V i TIME
2.0
'
50 I
t HR I
I00 I
'
'
2053
Li
150 I
'
200 I
'
L.i-I 1.0
0.7 I-
W U
_
0.5
z
m n- 0 . 3 0 v) a3 <: 0 . 2
0.1
I
I I0
, TIME,
I 20 HR,
I K
I 30
I
AND
C=
I 40
Fig. 4. Log concentration e~m vs. time for reactions of three alkali metals with liquid a m m o n i a at 23 --- 1°.
deviation from linearity near time zero was observed for any reaction; this contrasts with the data of Corset and Lepoutre [l(b)]. In the case of lithium, two linear sections intersect about 60 hr after initiation; the change to the slower rate occurs at about the time LiNHz begins to precipitate. First-order plots for potassium were linear through about 85 per cent, for lithium through 70 per cent reaction. Although the cesium reaction is relatively rapid at first and is not first-order, there is little doubt that reaction (1) is indeed occurring, because of the linearity of the plot in Fig. 3. This kinetic behaviour of the cesium reaction is at variance with the observation that the cesium-ammonia reaction is catalyzed by the presence of cesium amide[15]. It has also been reported that significant decomposition of metal-ammonia solutions occurs in solutions of strongly basic species [4(b)]. It seems clear that, for our reactions, amide in solution is not a catalyst; indeed precipitation of lithium amide actually decreases the rate of the lithium reaction, probably due to the precipitation of amide on the wall of the reaction vessel. Rate o f reactions in special cells. Because the potassium reaction is straightforward first-order and fairly reproducible, it was used to evaluate the effect of variables in cell treatment. Results of several tests are shown in Table 2. To measure the effect on the rate of a change in glass surface area, cells without diaphragms were constructed with perforated glass plates dividing the cells approximately in half. A cuvette was attached to the half below the plate; the upper half was packed with fire-polished lengths of 8 mm Pyrex tubing. After standard cleaning and bake-out, potassium was distilled into the cuvette and 15. J. W. Hodgins, Can. J. Research, 27B, 861 (1949).
2054
D.C.
JACKMAN
a n d C. W. K E E N A N Table 2
Run
Initial moles K/1
t~2, hr
Comments
K-6
1.04 x 10-~
Ratio of areas, 3.3; ratio of rate constants, 4.4
K-7
0-96 x 10-3
Cuvette end, 19.3 Packed end, 4.4 Cuvette end, 23.9 Packed end, 5'9 370 380 620 4300 (?)
K-8 K-9 K- I 0 K- I 1
3-3 × 6.0 × 1.34 × 1.3 ×
10-4 10-4 10-a 10-4
Ratio of areas, 3.6; ratio of rate constants, 4.0
Fluoride treated Fluoride treated K vapour bake-out K vapour bake-out
ammonia was condensed directly on the potassium at -78°; the cell was warmed in about two hours to room temperature, and then kinetic data were collected. The rate of the potassium-ammonia reaction was determined first in the cuvette end for about 3 hr, then the cell was inverted and the rate determined in the packed end. In the latter case, the solution was poured into the cuvette periodicaUy for spectral monitoring. As the data for runs K-6 and K-7 in Table 2 show, the increase in the rate of the reaction in the packed end is roughly proportional to the increase in the area of the glass in contact with the solution. The longer half-lives in the cuvette ends of these two cells as compared to those for the potassium runs in Table 1 are presumably accounted for by the absence of diaphragm fragments in the two solutions described in Table 2. The half-times of runs K-6 and K-7 in the cuvette e'ffds agree fairly well with the results reported by others [l(b), (c)]. Dr. M. L. Hair called our attention to the catalytic role of silanol groups on glass surfaces and to the presence of both Bronsted and Lewis acid sites on Pyrex, e.g., silanol groups and boron atoms. T o determine the effect of replacing silanol ----OH with - - F , by the technique mentioned under Cell Treatments, two runs were made, K-8 and K-9 in Table 2. These cells had no diaphragms, so the first kinetic readings were made about two hours after the cells were sealed, when the solutions had warmed to room temperature. The effect of the fluoride treatment can be seen by comparing t1/2 for K-8 and K-9 with the cuvette-end times of K-6 and K-7. A bake-out temperature of just over 500 ° in v a c u o after ammonium fluoride treatment does not effect a complete replacement o f - - O H by - - F on silica [ 11 ], but this was the maximum temperature we felt we could use without mechanical damage to our cells. Baking out at about 500 ° probably does eliminate adjacent ----OH groups on silica[11]; this would decrease the likelihood of hydrogen bonding between two silanol groups, a process which is attractive in providing a mechanism for the production of H~ molecules as a result of e~m attack on adsorbed NH3 molecules. In the case of the potassium reaction, the fluoride treatment of the glass presumably decreases the rate of the reaction because the number of catalytic ---OH groups is decreased. Further work is necessary to permit even this simple
Reactions of metal-ammonia solutions - V!
2055
conclusion, for preliminary results show that in its first-order re#on the cesium reaction is a little faster in fluoride-treated cells[16]. It has been found that partial replacement of ----OH by - - F enhances the acidity of the remaining ---OH groups [ 17]; possibly this effect is more important for the cesium reaction than for the potassium. As mentioned under Cell treatments, two reactions were carried out in cells which were baked out with potassium vapour; the aim was to remove all Bronsted acid sites. The results of two runs, K-10 and K-11, are given in Table 2. These cells had no diaphragms; metal and then ammonia were distilled into the baked out cell via break-seal connections. The tl/2 for K-11 is only approximate, because the reaction was followed for only 160 hr. Miscellaneous observations. The suggestion was made in the Discussion Section of Ref. [l(c)] that the rate of the metal-ammonia reaction might be sensitive to stirring, if the decomposition occurs at the wall. To test the effect of stirring, the cell in Run K-9 was agitated vigorously with a wrist-action shaker after the rate of the reaction, as shown in Table 2, had been determined in the unshaken cell. The ratio of the first-order rate constants for the shaken vs. the unshaken cell was 1.62, but the cell wall area bathed by the solution for the shaken vs. unshaken cell was estimated as 1-6, so it appears that the agitation per se resulted in no increase in the rate of the reaction. The magnitude of the activation energy of the metal-ammonia reaction is consistent with that of a diffusion-controlled reaction[l(c)], so it is of interest to consider such a possibility. Jolly has claimed [18] that solvated electrons in ammonia have an abnormally high rate of diffusion, his view evidently being based on the fact that the anionic species in a metal solution does have a large transference number. We hold, however, that the rate of thermal diffusion of the electron cannot be abnormally high, because of the electroneutrality limitation. In the case of concentration differences in a non-homogeneous metal-ammonia solution, solvated electrons cannot diffuse to establish homogeneity any faster than the metal ions which must accompany them. In the case of solutions such as we are reporting on, which can be assumed to be homogeneous with respect to metal ions, the diffusion rate of the electrons toward the wall will be limited by the rate at which amide ions diffuse away from the wall. It is agreed that electrons can migrate without carrying ammonia molecules with them, but we feel that in the case of thermal diffusion the electron must come within the coulombic field of a positive charge, that is, within a few angstroms, in order that the positive charge be attacked. We have relied on a similar argument elsewhere[19], and feel that Jolly's view is invalid. CONCLUSIONS
At room temperature, 10-4 to 10 -3 M potassium reacts with ammonia in a reaction which is first-order with respect to solvated electrons and which is catalyzed by borosilicate glass. Probably silanol ----OH groups are the main catalytic sites, although Lewis acid sites may be active also. The adsorption by 16. 17. 18. 19.
Studies in progress by Joseph Knauer in this laboratory. I. D. Chapman and M. L. Hair, J. Catalysis 2, 145 (1963). W. L. Jolly, A dr. Chem. Ser. 50, 33 (1965). C.W. Keenan, H. V. Secor, E. J. Kelly and J. F. Eastham, J. A m. chem. Soc. 82, 1831 (1960).
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J A C K M A N and C. W. K E E N A N
glass of ammonia on both silanol groups and on non-Bronsted sites, possibly boron atoms, has been described [ 17]. Baking at >400°C and < 10-s mm for 24 hr removes adsorbed water so effectively that our spectroscopic analyses show that the amide produced is equivalent to the initial metal concentration. The results of Lepoutre and coworkers [ 1(b),(c)] indicate that if a little water does remain on the wall (as revealed by the initial curvature of their first-order plots), that potassium solutions of stability comparable to ours result after this water reacts. This is essentially the method of cell treatment reported by Burow and Lagowski[4(b)]; before use, they aged glass vessels 48 hr with 10 -z M potassium solutions. They stated that a 10-5 M solution showed no decomposition in 20 hr. Though their temperature is not given, one of their working temperatures is -70°C; at this temperature the half-life of the reaction in one of our cells without a diaphragm, treated according to our standard procedure, would probably be more than a year, so we would find "no decomposition" in 20 hr, also. They attributed the stability of their solutions to the replacement by ammonia of "adsorbed water and/or hydroxyl species on the surface of the glass . . . . " How ammonia could replace a hydroxyl group is not clear to us; in any event, our results indicate that silanol groups are not removed by dilute metal-ammonia solutions, but indeed remain as catalysts for the metal-ammonia reaction. For the potassium reaction the catalytic activity of the glass can be greatly decreased by treatment with ammonium fluoride prior to baking or by baking with potassium vapour; these findings are consistent with the view that -----OH groups are being removed by these treatments. It is proposed that in our standard cells (Table 1) an ammonia molecule is bound to a surface acid group and that the resulting enhancement of the acidity of the ammonia hydrogens leaves them open to attack by a solvated electron. Possibly attack of adjacent adsorbed NHa molecules facilitates formation of H2. The most effective catalysts for the metal-ammonia reaction are the transition metals. It has been suggested/20] that certain catalytic activity of nickel is due to is possessing "inherent acid properties." Under the conditions described, the over-all rates of metal-ammonia reactions for three metals decrease in the order K > Cs > Li, The lithium reaction is, like the potassium, first-order over the majority of the reaction; there is no evidence that amide produced in solution is a catalyst for the reactions. It seems likely that the lithium reaction would be even slower if the lithium could be transferred into the cell in as clean a state as the other two metals. The very slow rates of the potassium reaction in the cells treated to decrease the catalytic activity of the glass indicate clearly that the homogeneous reaction of solvated electrons with ammonia may be immeasurably slow even at room temperature. We are now trying other treatments of the glass, e.g., baking with chloromethylsilanes[21], and we plan to make cells of other materials. Further 20. H. Pines, M. Shamaiengar, and W. S. Postl, J. Am. chem. Soc. 77, 5099 (1955). 21. We are indebted to Dr. M. L. Hair for this suggestion. See, for example, T. E. White, Soc. Plastics Ind., Proc. 20, 3-B, I (1965). Studies in progress by W. E. Moddeman in this laboratory indicate that baking with dichlorodimethylsilane is very effective in decreasing the catalytic activity of the glass.
Reactions of metal-ammonia solutions - V 1
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improvement should make possible long-term experiments at room temperature on the same stable solution. Acknowledgements--The
authors are indebted to Drs. James L. Dye, M. L. Hair, and M. H. Lietzke for advice on several points.