ROBERTR. DEWALD AND RICHARD V. TSINA
4520
A Kinetic Study of the Reaction of Water and t-Butyl Alcohol with Sodium in Liquid Ammonia1 by Robert R. Dewald and Richard V. Tsina2 Department of Chemistry, Tufts University, Medford, Massachusetts 02166
(Received June 3,1968)
The reaction of two weak acids, water and t-butyl alcohol, with sodium in liquid ammonia has been studied IC1 conductometricallyat -33.9’. The kinetic data are consistent withthe general mechanism: ROH NH,
+
+
+
2
+
ka
NH4+ RO- and NH*+ earn- NHI 0.5Hz, where R is t-butyl and H. For &butyl alcohol, kl was found to be (2.5 f 0.7) X IOFe M-’ sec-l and 6 / k 3 = 1.3 =k 0.3. The values found for water were kl = (7 i= 2) X M-’ sec-l and kz/ka = 0.55 =k 0.16. The bimolecular rate constant, ks, is estimated to be (4 =t2) x 108 M-l sec-l at -33.9O. -+
Introduction Alcohols react readily with solutions of alkali metals in liquid ammonia with the evolution of hydrogen and the formation of a l k o x i d e ~ . ~Kraus !~ and White,4 in work of a qualitative nature, reported that the reaction between sodium and ethanol in liquid ammonia was vigorous at first but rapidly slowed down, with the result that after 2.5 hr the reaction had gone to only 70% completion. They noted that extensive precipitation occurred. Later, Kelly, et aZ.,5 reported kinetic data for the ethyl alcohol-sodium ammonia system. Using ammonia labeled with tritium, these workers followed the progress of the reaction by monitoring the volume and activity of the evolved gas (Hz and HT). They found that the reaction was initially first order in EtOH and zero order in Na. Between 25 and 50% completion, the reaction order was found to be indefinite, but during the last SO%, the reaction was slower and followed second-order kinetics, first order in both metal and alcohol. Only one set of characteristic data, however, was reported. Recently, JollyaP7 evaluated these data using a Powell plot based on the presumed mechanism
__
f
= 0.10-0.40, where f is the fraction of the reaction, the experimental points deviated unaccountably from the theoretical curve, but the fit was later7 considerably improved when the formation of a complex between ethanol and the ethoxide ion was introduced. Jolly7 pointed out that this assumption is quite arbitrary. One might further note that Kelly, et al., employed relatively large concentrations of ethyl alcohol (0.2-0.4 M ) and sodium (0.2-1.0 M ) , since fairly large amounts of gas were needed in order to perform an accurate analysis. The observations of Chablaya and Kraus4 lead one to suspect that some precipitation of sodium ethoxide probably o ~ c u r r e dthe , ~ effect of which on the observed reaction kinetics is not a priori predictable. We have studied the reaction kinetics of sodium with the two weak acids, water and t-butyl alcohol, with the aim of obtaining kinetic data in a system in which precipitation was not occurring. Water has been shown to react readily with sodium-ammonia solutions.8 Pleskovs reported that solutions of sodium in liquid ammonia may be employed as a rapid method for the determination of small quantities of water. The ob-
kl
+ NH3 NH4+ + RONH4+ + earn-“8, NH3 + 0.5Hz
ROH
ka
(1)
(2)
where R is ethyl. If the ammonium ion concentration is at a low steady-state value, the rate law
is obtained. Jolly found that the mechanism, (1) and (2), was at least qualitatively compatible with the data reported. The best Powell plot fit to the data was obsec-‘ and lc2/k3 = 6 X tained by setting IC1 = 8 X loa. It was found initially,8 however, that in the region The Journal of Physical Chemistry
(1) Presented in part before the Physical Chemistry Division a t the 154th National Meeting of the American Chemical Society, Chicago, Ill., Sept 1967. (2) This paper is part of a thesis submitted to Tufts University in partial fulfillment of the requirements for the degree of Doctor of Philosophy. (3) E. Chablay, Ann. Chim. (Rome), 8 , 145 (1917). (4) C. A. Kraus and G. F. White, J . Amer. Chem. Soc., 45, 768 (1923). (5) E. J. Kelly, H. V. Secor, C. W. Keenan, and J. F. Eastham, ibid., 84, 3611 (1962). (6) W. L. Jolly, “Non-Aqueous Solvent Systems,” T . C. Waddington, Ed., Academic Press, New York, N. Y., 1965, p 39. (7) W. L. Jolly, Advances in Chemistry Series, No. 50, American Chemical Society, Washington, D. C., 1965, p 27. (8) W. C. Fernelius and G. W. Watt, Chem. Rev., 29, 195 (1937). (9) V. A. Pleskov, Zavodskaya Lab., 6, 177 (1937); Chem. Abstr., 31, 61351 (1937).
REACTION OF WATERAND &BUTYL ALCOHOL WITH SODIUM served reaction, which is almost certainly due to solvated electrons, earn-, can be written
+ HzO -+
earn-
0.5Hz
+ OH-
(4)
in which NaOH precipitates readily.lO Chablay* noted that tertiary alkoxides are more soluble than secondary or primary alkoxides. Consequently it was felt that precipitation might best be avoided by using t-butyl alcohol as one of the weak acids in this study. By employing low reactant concentrations, we have successfully avoided precipitation during the course of the reaction.
Experimental Section Ammonia (Matheson) was condensed in vacuo in a trap containing sodium metal. The blue solution was stored in this trap in contact with an iron magnet until the extent of the autodecomposition reaction Na
+ NHa +NaNHz + 0.5&
(5)
was considerable. About half of the ammonia was distilled into a second trap which had been evacuated and flamed until the pressure stabilized at less than 2 X 10-6 torr. The ammonia was then distilled back into the trap containing the sodium. The blue solution was next frozen with liquid nitrogen and degassed. The second trap was again evacuated and flamed. The cycle of distilling back and forth, freezing, degassing, and flaming was repeated at least five times. The ammonia was then stored in contact with sodium, distilling enough for one or two experiments into the second trap just prior to use. Sodium (United Mineral and Chemical Co.) was distilled twice in vacuo and was stored in Pyrex capillaries. Ammonium bromide (Fisher reagent) samples were prepared by placing the salt in break-seal tubes which were sealed off under vacuum after evacuation to about lov6torr. The reaction of the two weak acids with sodium was followed conductometrically in an apparatus similar to that described elsewhere." The reaction vessel was constructed of Pyrex and had a calibrated bulb for volume determination. The conductance cell" and break-seal tubes containing samples of the reactant and solid NH4Br were sealed onto two side arms. The electrodes were either gold-plated platinum or in a few cases gold-plated tungsten. The cells were calibrated with standard KC1 solutions using the data of Jones and Bradshaw.lZ Either a high-precision ac bridge similar to that described elsewherel3 or a Wayne Kerr Universal bridge B 221A was used. The procedure followed in measuring the resistance of the solution is also described elsewhere." The t-butyl alcohol (Fisher Certified) was first distilled through a 3-ft packed column, and the middle fraction was introduced into a trap. The alcohol was next degassed by repeated freezing and evacuation. It was then distilled in vacuo into a second trap into
4521
which sodium metal had been distilled. After the sodium had completely reacted, the alcohol was degassed by repeated freezing and pumping. The alcohol was distilled into tared break-seal tubes which were then sealed off under high vacuum and weighed. Water samples (starting with doubly distilled conductance water) were also prepared as described above, except the sodium metal was omitted. During the course of our investigation it became apparent that very small quantities of water would be necessary for some experiments (around 5 X 10-6 mol or less). These samples were prepared by equilibrating the water reservoir at an appropriate, thermostated temperature with evacuated bulbs of known volume on the vacuum line. The contents of the bulbs were then transferred to breakseal tubes, held at liquid nitrogen temperature, and then sealed under high vacuum. The procedure followed for the sodium-ammonia solution preparation in the Pyrex apparatus is described e1sewhere.l' We feel that this technique is instrumental in removing traces of water from the walls of the apparatus as well as dissolved Hz gas from the solution, After the blue solution had been thoroughly degassed, the apparatus was disconnected from the vacuum line and immersed in a refrigerated bath at -33.9 f 0.2" (Dow Corning No. 200 silicone fluid was used as the bath liquid.) The resistance of the sodium-ammonia solution was subsequently monitored as a function of time for approximately 1 hr to ensure stability. Next, the break-seal containing the t-butyl alcohol or water sample was broken, the solution was mixed, and the resistance of the mixture was followed as a function of time. Bright platinum electrodes were found to be unsuitable for making measurements after either the alcohol or water was added to the sodium-ammonia solution, probably because of contamination of the electrodes by hydrogen evolved during the reaction. The goldplated electrodes, however, were found to be quite suitable for taking conductance measurements in the reacting systems. This observation is consistent with the findings of Windwer and Sundheim14 in a study of the solutions of the alkali metals in ethylenediamine. I n the experiments with 2-butyl alcohol, after the reaction rate became negligible (unbleached solution) , the evolved Hz gas was pumped through three liquid nitrogen traps using mercury leveling bulbs and was collected in a calibrated gas buret in which the volume and pressure of the gas were measured. During the (10) M. Skossarewsky and 11 (1916). (11) R.
N. Tchitchinadze, J . Chem. Phys., 14,
R. Dewald and J. H. Roberts, J . Phys. Chem., 7 2 , 4224
(1968). (12) G. Jones and B. C. Bradshaw, J. Amer. Chem. SOC.,5 5 , 1780 (1933). (13) G. E. Smith, Ph.D. Thesis, Michigan State University, 1963. (14) 8. Windwer and B. R. Sundheim, J . Phys. Chem., 66, 1254 (1962).
Volume 7.8, Number 19 December 1968
ROBERTR. DEWALDAND RICHARD V. TSINA
4522 12
pumping, the blue solution was frozen and kept a t liquid nitrogen temperature. Next, solid ammonium bromide was added, via a break-seal, to the blue solution, and the remainder of the H2 gas was collected. Finally, the total amount of metal could be determined from the stoichiometry of the reactions
+ Na 0.5H2 + NH3 + NaBr ROH + Na -+ 0.5H2 + RONa
NH4Br
(6)
--f
(7)
;
in which R is t-butyl. The initial sodium concentration, as determined from the evolved H2, was in good agreement with the value determined from the initial resistance of the sodium-ammonia solution. I n this work, the conductivity measured for the reacting mixture was taken to be due to sodium alone, and the contribution to the conductance by either of the weak acids was considered to be negligible. Support for this procedure is apparent from the following data: for a lod2 M solution of water in liquid ammonia at -33.9", the specific conductance16 is 2.45 X 10-7 ohm-' cm-l compared with 1.06 X ohm-' M sodium-ammonia solution" at the cm-l for a same temperature. Similarly, the contribution to the total conductivity due to alkoxides and hydroxides'o produced was considered negligible in the initial phase of the reaction from which data were used in calculating rate constants. Conductance-concentration data obtained in this laboratoryll for sodium-ammonia solutions were used in the calculations.
Results a. The Reaction of t-Butyl Alcohol with Sodium in Liquid Ammonia. Table I contains a summary of the results obtained from the reaction of sodium with tbutyl alcohol in liquid ammonia at -33.9". Figure 1 shows three sodium concentration us. time plots. It should be noted that the alcohol is present in excess. As shown in Figure 1, the sodium reacts rapidly at first and then the reaction rate levels off to an almost negligible change with time, even though the solution
:
20
0'
40
60 TIME,min.
100
80
120
Figure 1. Sodium concentration, M , us. time plots for the reaction of sodium with t-butyl alcohol at -33.9' in liquid ammonia: W, run 114A; 0, run 28C; 0 , run 113A.
remained blue. Other experiments using different initial concentrations gave similar results. When sodium amide was present, as determined by following the conductance of a decomposing sodium solution (for 24 hr), no large drop in the initial sodium concentration was observed upon addition of the alcohol. I n this experiment the initial concentrations were about 1.2 X 4.2 X lo-*, and 2.66 X M sodium, sodium amide, and t-butyl alcohol, respectively. No precipitation was observed during the course of the reaction of sodium with t-butyl alcohol in any of the experiments. This observation is consistent with the reported greater solubility of tertiary a l k ~ x i d e s . ~ b. The Reaction of Water with Sodium in Liquid Ammonia. Table I1 lists the kinetic data obtained for the reaction of sodium with water in liquid ammonia. It can be seen that the range of concentrations covered is indeed very large. I n experiments 31C, 45B, and 46B (high concentrations) precipitation was
Table I1 : Kinetic Data for the Reaction of Sodium with Water in Liquid Ammonia at -33.9' Table I: Kinetic Data for the Reaction of Sodium with &Butyl Alcohol in Liquid Ammonia at -33.9' Run no.
11B 28C 30C 102A 104A 113.4 114A 142A Av a
10n[Nal,
102[ (CHa)aCOHI,
kr/ka
M
M
1.1
6.07 0.882 1.34 5.69 5.46 0.657 1.13 0.405 . 3 .~. .
5.14 1.54 4.57 4.43 2.14 1.85 1.74 0.805
loekl,
M-1
sea-1
3.3 1.7 3.8 1.7 1.7 3.1 2.1
1.8 1.5 1.3 0.78 1.6 1.2 2.3 0.97 2.5&0.7" 1 . 3 i ~ 0
Average error.
The Journal of Physical Chemistry
Run
lOeki,
no.
M-1 sea-1
kz/ks
29B 38B 34c 35c 27B 46B 45B 31C Av
4.0 5.6 6.5 8.0 12 a
0.63 0.87 0.49 0.32 0.21 0.76 0.58 0.55 0.55 f O.16*
a
a
7 f 2b
' Data insufficient to calculate kl.
104[Nal, M
10'[HzOI,
0.400 0.175 0.714 0.535 1.53 82.3 28.7 84.7
0 268 0.116 0.118 0.0757 0.123 89.3 275 18.8
..
* Average error.
(16) R. R. Dewald and R. V. Tsina, unpublished work.
M I
...
REACTION O F WATER AND t-BUTYL ALCOHOL WITH SODIUM
4523 40ki(a - X)(b - X ) d~ _ -- _ b -x
dt
IOoo
+ (kp/k3)~
(12)
-
in which a = (ROH)i, b = (eam-)i, x = (R0H)i (ROH) = (RONa), and 40 is the molar concentration of the solvent, ammonia, in the pseudo-first-order step, eq 8. Ingold, et al.," have shown that when V z and Va (V, is the velocity of step n) are similar and (RO-)i = 0, eq 12 may be integrated to yield
O.lo
F , 90
40
80
60
100
I20
TIME,min.
Figure 2. Semilogarithmic plots of sodium concentration, M , us. time for the reaction of sodium with water a t -33.9" in liquid ammonia: (3, run 46B; W, run 31C; 0, run 27B; 0 , run 38B.
observed during the course of the reaction. Owing to the rapid initial slope change in the above three experiments, the data were insufficient to calculate kl, although reasonable values of kz/ka could be obtained. The five runs, 25C, 27B, 29B, 34C, and 38B, proceeded without observable precipitation and were used in calculating both kl and k2/k8. Figure 2 shows examples of semilogarithmic plots of the sodium concentration vs. time for the reaction of sodium with water in liquid ammonia at -33.9'.
Discussion For both weak acids, water and t-butyl alcohol, the kinetic data can be best explained in terms of reactions 1 and 2. This mechanism is similar to that suggested for other systems as outlined by Russell.16 For the scheme ki
ROH NH4
loa
RO-
+ earn-
+ NH4+
NHa
+ 0.5Hz
(8) (9)
in which kl pertains to a pseudo-first-order step, if the concentration of NH4+ is small so that at any time -d(ROH) dt
- -d(e,,) dt
2d(Hd dt
the rate expression becomes
Equation 11 may be rewritten in the form"
(10)
The above expression was used to evaluate kl and k2lk.8 for the t-butyl alcohol-sodium and water-sodium runs in Tables I and 11. It might be noted, however, that Ingold, et aZ.,l7 have pointed out that eq 13 is not very sensitive to the choice of k2/k3, a choice of 1.0 in some of their data being about as good as a choice of 1.2. The values of kz/k3 obtainable by this method therefore cannot be expected to be reliable to better than about 20%. Our precision ranges from 23% in the case of t-butyl alcohol experiments to 29% in the case of water experiments. I n the latter case, the low concentrations needed to avoid precipitation are probably reponsible for the somewhat higher imprecision. For the case of the partly decomposed sodium solution in which a substantial quantity of amide ion was initially present, it would be expected that the ammonium ion concentration would be small since NH4+
+
"2-
+2NH3
(14)
is known to proceed far to the right.ls I n this experiment a slow initial rate for the disappearance of sodium was observed upon addition of the alcohol. The slow reaction rate in this case would seem to substantiate the proposed two-step mechanism. The equilibrium constant for the ammonolysis
+
+
H2O 3" NH4+ OH(15) in wet ammonia at -33.9' is unreported. We performed a conductance experimentls in which known quantities of water vapor were added to carefully purified ammonia (as determined from its conductivity), raising the water concentration from 3.56 X 10-6 M to a final value of 8.69 X M . We were able to estimate an equilibrium constant at -33.9" for reaction 15 of about (4 f 2) X 10-l2, which is comparable with the value estimated at 25' by Clutter and Swift.18 (16) G. R. Russell in "Techniques of Organic Chemistry," S. L. Friess, E. 8. Lewis, and A. Weissberger, Ed., Vol. 8 , Interscience Publishers, New York, N. Y.,1961,Part 1, p 383. (17) E. D.Hughes, C. K. Ingold, 9. Patai, and Y . Pocker, J . Chem. Soc., 1230 (1957). (18) H. Smith, "Organic Reactions in Liquid Ammonia," Interscience Publishers, New York, N. Y.,1960,p 39.
Volume 7.9, Number 15 December 1088
DONALD W. RICEAND N. W. GREGORY
4524 This result allows us to calculate a tentative value for the rate constant, kat of reaction 2. We find, using the kl and k2/k3 averages tabulated for the water-sodium reaction in Table 11, a value of k3 = (4 i 2) X lo6 M-l sec-' at -33.9' in liquid ammonia. We may compare this result with the results found for the comparable reaction of the hydrated electron, eaq-, with the ammonium ion in water at zero ionic strength, as summarized by Rabani.zO The rate constants reported vary from 1.1 X lo6 to 1.8 X lo6 M-' sec-' a t ambient (presumably) temperature. Our estimate of kl appears to be in agreement with the work in the aqueous systemz0 and is also consistent with the rate of the reaction between cesium and ethylenediammonium ions in ethylenediamine.21 Our kinetic data for the reaction of the two weak acids with sodium in liquid ammonia support the conclusion that the mechanism for reaction 4 is reactions 1 and 2.22 The average values of kl and k2/k3 determined in the present work are in generally poor agreement with the comparable constants deduced from the data of Kelly, et aL,6 by Jolly.8J We feel that this
discrepancy may be due to unavoidable precipitation occurring when ethyl alcohol and sodium react at the concentrations employed by the former workers. We are now reinvestigating the kinetics of the reaction of ethyl alcohol with sodium in liquid ammonia. Also, the value reported in this work for the rate constant of the ammonium ion-solvated electron reaction is within the range of the stopped-flow method, and an effort is now being made in this laboratory to measure this constant directly.
Acknowledgment. This research was supported by the National Science Foundation under Grant No. GP 6239. (19) D. R. Clutter and T. J. Swift, J . Amer. Chem. Soc., 90, 601 (1968). (20) J. Rabani, Advances in Chemistry Series, No. 50, American Chemical Society, Washington, D. C., 1965,p 242. (21) L. H. Feldman, R. R. Dewald, and J. L. Dye, Advances in Chemistry Series, No. 50, American Chemical Society, Washington, D. C., 1905,p 163. (22) R. R. Dewald and R. V. Tsina, Chem. Commun., 647 (1967).
Vaporization Equilibria in the Sodium Chloride-Zinc Chloride System by Donald W. Rice and N. W. Gregory Department of Chemistry, Universityof Washington,Seattle, Washington 98106
(Received June 6 , 1068)
Partial pressures of ainc chloride and of species of the form Na,ZnClz+, in equilibrium with condensed mixtures of NaCl and ZnClz, and with pure NaCl(s), have been derived from effusion and transpiration data. Evidence is found cor the existence of the solid-state compound NazZnClr, for which values of AH{' = -299 f 3 kcal mol-' and So = 95 f 4.5 cal de$-' mol-' at 625°K are derived. I n the presence of NaCl(s), partial pressures of the complex in the vapor phase, assumed to be NaZnCla(g), are only ca. 1% of those of ZnClz a t 600'.
Relatively little is known about the zinc chloridesodium chloride binary system. A melting point study by Nikonowa, Pawlenko, and Bergman indicates the existence of only one solid-state intermediate compound, NasZnCl4, with an incongruent melting point around 41OO.l Its crystal structure has not been determined. Dijkhuis and Ketelaar studied NaC1-ZnClz melts by emf methods at GOO" and found that deviations from the ideal Temkin model appeared to be a maximum around X N ~ C=I 0.59; they suggested that this dissymmetry may be due to complex ion formation.z Ellis reported evidence from Raman studies for ZnCl3and ZnC12- in KC1-ZnC12 meltsa3 Markov and Volkov observed that the volume change on mixing liquid NaCl The Journal of Physical Chemistry
and ZnClz is negative for mole fractions of NaC1 less than 0.5 and suggested the formation of complex ions.4 We now report a thermodynamic study of this system in which vaporization characteristics between 300 and GOO" have been determined by effusion and transportation experiments. Experimental Section The torsion effusion and Knudsen effusion apparatus, (1). N. Nikonowa, S. P. Pawlenko, and A. G. Bergman, Bull. Acad. Scz. URSS Classe Sei. Chim., 391 (1941).
(2) C. Dijkhuis and J. A. Ketelaar, Electrochim. Acta, 12, 795 (1967). (3) R.B.Ellis, J. Electroehem. SOC.,113, 485 (1966). (4) B. F. Markov and S. V. Volkov, Ukr. Khim. Zh., 29, 946 (1963).
