NOTES
3193
t
. .
I
250
200
300 T OK.
350
40 0
Figure 1. Heat capacity of U0p.s r\ compared with data on the same sample by Osborne, et al., -0-,2 and with that of UOZ- -v- -,T - - -,8 a-UO2.aa -* . * - ',* pu02.81- * 4- -,* UOZ.66 - -A- -,E - -,IO and uo8 4. * ..*
-
- -
---
-
. ..
dimensions compared to UOn according to Belbeoch, et al.," and Willis,16 while Andresen16 indexed it as tetragonal with A = 2 a d and C = 2a. Placement of the interstitial oxygen atoms in the cubic interstices with eight oxygen atoms as nearest neighbors is clearly objectionable from a chemical point of view, and the thorough neutron diffraction study of single crystals of UIOg by Willis15 has shown that incorporation of the interstitial oxygen atoms is accompanied by ejection of about 10% of the oxygen atoms from their positions in the fluoritetype structure. The interstitial and ejected oxygen atoms are located in the vicinity of the center of the cubic interstices, 60% of them 0.86 8. away along the (110) directions and 40% 1.05 A. away along the (111) directions. Although the detailed ordering of the interstitials has not yet been worked out, it is apparent that about half of the uranium atoms are formally present as Us+. Thus, a formula 2U02.U20s represents the situation better than 3uo2 * UOa. At low temperatures an ordered arrangement of interstitial and displaced oxygen atoms is expected, which again results in an ordered distribution of the cations. It seems tempting to relate the A-type thermal anomaly with an orderdisorder process involving the interstitial and displaced oxygen atoms although no changes in the structure of U409have yet been reported in this temperature region.
In the neroth approximation, the extra oxygen atom in U40gmay either be in a fixed position or randomly distributed over four positions, which, for U O Z . ~ , would lead to an entropy increment of I/& In 4 or 0.68 cal./(g.f.m. OK.). The actual situation is apparently much more complicated, in that different possible modes of ejecting the normal oxygen atoms tending to increase the entropy may exist, while the requirement of even distribution of U4+ and US+will appreciably reduce the entropy. Detailed delineation of the structure is an obvious desideratum. The thermal effect may be somewhat broadened by radiation damage to the sample accumulated over a 10-year period at room temperature. Although repetition of the measurements after conditioningthe sample at high temperature would test this point, a more thorough elimination of this possibility will be achieved by heat capacity measurements to 800°K. on a new sample at the University of Oslo.
Acknowledgment. The partial financial support of the U. S. Atomic Energy Commission is appreciated. (14) B. Belbeoch, C. Piekarski, and P. Perio, Acta Cryst., 14, 837 (1961). (15) B. T. M. Willis, J. phys. radium, 25, 431 (1964). (16) A. F. Andresen, Enlarged Symposium on Reactor Materials, Stockholm, Sweden, Oct. 1959.
Dissociation and Homoconjugation of Certain Phenols in Acetonitrile
by J. F. Coetzee' and G. R. Padmanabhan2 Department of C h k t r y , Unioersity of Pittsburgh, Pttsburgh, Penneglwnia 16.913 (Received March $6,1966)
As compared to water, acetonitrile provides striking differentiation in the dissociation of Brgmsted acids, mainly as a result of the following three factors. (1) The proton-acceptor power of acetonitrile is smaller than that of water by five powers of ten, as given by its reaction with a reference acid of type BH+.a Hence, the primary reaction of Br#nsted acids with the solvent, which is the first step in the over-all dissociation reaction ~~
(1) Address all correspondenceto this author.
From the Ph.D. thesis of this author, University of Pittsburgh, 1963. (3) J. F. Coetzee and D. K. McGuire, J. Phzls. Chem., 67, 1810 (1963). (2)
Volume 69, Number 9 Beptember 1966
NOTES
3194
HB"+
+ S E SH+B("-1)+ E SH+ + Bb-1I-I(1)
where n can be positive, negative, or zero, is less complete in acetonitrile than in water. (2) The dielectric constant of acetonitrile (36.0) is smaller than that of water (78.5). Hence, for purely electrostatic reasons, as indicated by the Born equation, the second (dissociation) step of reaction 1 will be more sensitive to variations in the effective radii of the ions in acetonitrile than in water, particularly for acids with n values other than +l. (3) A critical factor for acids with n values that are negative or zero is the ljmited capacity of acetonitrile to stabilize anions by hydrogen bonding. Consequently, subtle variations in the polarizability or other properties of the anion affect the equilibrium constant of reaction 1 much more markedly in acetonitrile than in water. In fact, certain anions derive stability by resorting to hydrogen bonding with undissociated acid, with the result that dissociation reactions such M the following occur. 2HBO
+ S E SH+BHB- eSH+ + BHB-
(2)
Considerable qualitative information is available on the differentiation of uncharged or anionic acids in acetonitrile, mainly from the results of empirical potentiometric titratiom2 However, quantitative information is lacking. The main exception is provided by the results of a conductometric and spectrophotometric study by Kolthoff, Bruckenstein, and Chantooni on hydrochloric, hydrobromic, nitric, sulfuric, and picric acids.4 In this communication we report the results of an exploratory potentiometric study of five phenols in acetonitrile. We have shown before that under properly controlled conditions the glass electrode responds reversibly to hydrogen ion activity in anhydrous acetonitrile! The measurements reported here were carried out in a similar manner.
Experimental Sohio acetonitrile was purified as described beforee6 Allied Chemical Co. C.P. phenol was used without further puriiication. The remaining phenols were recrystallized and then dried in vacuo ~ E I indicated: Eastman White Label o-nitrophenol (from methanol, at 25") and p-nitrophenol (water, 50"); Fisher reagent grade 2,4-dinitrophenol (water, 40"); and Baker and Adamson picric acid (twice from acetone, 80'). Tetraethylammonium perchlorate and picrate were prepared m described before.6 Tetraethylammonium o-nitrophenoxide and p-nitrophenoxide, and tetrabutylammonium phenoxide and 2,4dinitrophenoxide, were prepared by titrating the phenol with tetraalkylammo-
nium hydroxide, using a glass electrode to detect the equivalence point. The solutions were evaporated and then cooled in ice, and finally the salts wererecrystallized from ethyl acetate and dried in vacuo at 40-50". Measurements. For each phenol the potential of a Beckman General Purpose No. 1190-80 glass electrode was measured in a series of 10 t o 15 bufTer solutions conM ) of taining a constant concentpation (5.0 X phenoxide and varying concentrations (typically 2.5 X loq4 to 5 X lo-' M ) of the corresponding phenol, using the instrumentation described before! Absorbance measurements were made with a Beckman Model DB spectrophotometer, using a matched pair of silica cells (Beckman No. 46007). Further details are given in ref. 2.
Results and Discussion The results obtained at 25" are presented in Tables I and I1 and Figure 1. Equilibrium constants were calculated as described elsewhere for amines.' The results in Table I1 lead to the following conclusions. (1) All phenols tested, except picric acid, form relatively stable homoconjugate complexes, BH . B-. In addition, unsubstituted phenol also forms a 2: 1 complex, (BH)J3-. The existence of such complexes is well established, even in the solid state: and their formation has been inferred from the shapes of titration
Table I : Spectrophotometric Determination of Formation Constant of Homoconjugate Complex of +Nitrophenol in Acetonitrile CEB4
5.0 X 6.0 X 5.0 x 5.0 X
10-4 10-1
io-a 10-
C B ~
[B-]masdc
2.0 X 10-6 1.9 X 2 . 0 X 10-6 1.3 X 2.0 x 10-6 3.1 x 5.0 X 10-4 8.0 X
[BHB-ld
10-6 1.2 X 10-6 7.5 X 10-6 1.7 x 10-6 4.2 X
10-6 10-6 10-6 10-4
Kf,/
[HBI'
5.0 X 10-4 5.0 X 10-8 5.0
x
10-1
5.0 X 10-2
130 119 109 105
Total (analytical) molar concentration of o-nitrophenol. molar concentration of tetraethylammonium o-nitrophenoxide. Measured by absorbance a t 452 mp. BHB- has absorption maximum at lower wave length. Validity of Beer's law established. [BHB-] = CB[B-Imewd. e [HB] = CEB - [BHB-1. Kt,l = [BHB-]/[B-] [HB], assuming that ~ B H B - / ~ B - ~N E B1. Average value of 1.1 X loaassumed. (I
' Total
'
-
(4) I. M. Kolthoff, 9. Bruckenstein, and M. K. Chantooni, Jr., J . Am. Chem. SOC.,83, 3927 (1961). (5) J. F.Coetzee and G. R. Padmanabhan, J . Phys. Chem., 66, 1708 (1962). (6) J. F. Coetzee, G. P. Cunningham, D. K. McGuire, and G. R. Padmanabhan, Anal. Chem., 34, 1139 (1962). (7) J. F. Coetzee and G. R. Padmanabhan, J. Am. Chem. Soc., in
press. (8) D. Hadii, A. Novak, and J. E. Gordon, J. Phys. Chem., 67, 1118 (1963).
NOTES
3195
~
~~~
~~
Table I1 : Formation Constants of Homoconjugate Complexes and Acid Dissociation Constants of Phenols in Acetonitrile (AN) and Water (W) Acid
( -PKf,dAN
Phenol o-Nitrophen 01 p-Nitrophenol 2,4Dinitrophenol 2,4,6-Trinitrophenol
(-PKf,2)AN5
(-PKf.1dANb
5.76 2.00' 3.15 2.00 None
None 1.78 None
+
4.93
(PK~AN'
26.6 22.0 20.7 16.0 ll.og
Wdwd
ApKa'
10.00 7.23 7.15 4.11 0.71
16.6 14.8 13.5 11.9 10.3
+
'
Referring to over-all reaction: B2BH (BH)aB-. From measureReferring t o reaction: BHBBH $ (BH)tB-. ments on solutions containing 1.0 X M EtaNClOc as well, added to maintain constant junction potential Calculations as described elsewhere for amines (ref. 7). From compilation in R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," Butterworth and (pK,)w. Value from Table I: 2.04. Estimated uncertainty in pK values: f O . l . Co. Ltd., London, 1959. ' ApK. = (pK& Private communication from I. M. Kolthoff and M. K. Chantooni, Jr., University of Minnesota. On this basis, autoprotolysis constant of acetonitrile (ref. 5) becomes 3 X Calibration of scale of acidities is discussed further in ref. 7.
-
-0.3
0
'
+0.3 +Q6 +a9 +1.8 t1.5 +IS +2l +2.4 +2.7 t3.0 +3.3 109 (C"*/C.')
Figure 1. Response of glass electrode in phenol-phenoxide buffers in acetonitrile at a constant salt concentration of 5.0 X M: I, phenol; 11, o-nitrophenol; 111,picric acid. In all caws, potentials refer to arbitrary zero a t C ~ = B CB-.
curves in several. solvents; for a literature survey, see ref. 2. (2) Substitution of either phenol or p-nitrophenol with a nitro group in the ortho position to the hydroxy group markedly reduces homoconjugation, and substitution of the latter phenol in both ortho positions entirely eliminates complexation. This effect is attributed mainly to stabilization of the acid form by intramolecular hydrogen bonding. That such stabilization occurs is indicated by the fact that whereas 0- and pnitrophenol have similar dissociation constants in water, the former acid is considerably weaker than the latter in acetonitrile. This conclusion is supported by the infrared data of Brooks and Moman: which show that in acetonitrile some hydrogen bonding between hydroxy and nitro groups persists even in compounds such as 2-hydroxy-3-nitrobenzaldehyde,in which strong competitive hydrogen bonding to the carbonyl group
occurs. On the other hand, in aqueous solution onitrophenol shows little or no intramolecular hydrogen bonding,'O undoubtedly owing to effective competition from hydration. (3) The results show no evidence of dimerization of phenols in acetonitrile, in agreement with the behavior of other acids in this solvent.ll (4) In acetonitrile, differentiation in the dissociation of the phenols listed is increased by 6 pK units (8 kcal. mole-') over that observed in water. The stabilization of the acid form caused by ortho substitution must be more than offset by stabilization of the anion, undoubtedly in part as a result of the increasing polarizability. (5) A serious limitation of acetonitrile as a medium for acid-base reactions is the unavailability of a relatively strong base for this solvent. Any attempt to generate the lyate ion, CHzCN-, in sufficiently high concentration (e.g., by treating acetonitrile with lithium or sodium metal) leads to polymerization of the solvent. Alkali metal hydroxides are virtually insoluble (the absolute activity of the high charge density hydroxyl ion is extremely high in a nonhydrogen-bonding solvent), and anhydrous solutions of tetraalkylammonium hydroxides are unstable. Phenoxide ion is the strongest stable base that we have encountered in acetonitrile (pKb = 28.6 - 26.6 = 2.0). However, the homoconjugation undergone by this ion would detract from its usefulness as a base. It will be worthwhile to investigate the applicability of a hindered phenoxide, such as the tetraalkylammonium salt of 2,6-di-t-butylphenol. (9) C. J. W.Brooks and J. F.Morman, J . Chem. SOC.,3372 (1961). (10) L. B. Magnusson, C. A. Craig, and C. Postmus, Jr., J . Am. Chem. Soc., 86,3958 (1964). (11) J. F. Coetzee and R. Lok, J . Phys. Chen., 69, 2690 (1965).
Volums 69,Number 9 September 1966
NOTES
3196
Acknowledgment. We gratefully acknowledge financia1 support by the National Institutes of Health under Grant No. GM-10695.
High Resolution Mass Spectrum of Piperidine
by L. W. Daasch
c h a R ~ ~ ~ X ZandT CDevelopment ~ Labwatarisa, Edoeroaod Arsenal, Edgewood, Mawland (Received March 88,1986) Budsikiewicz, Djerassi, and Willirr.mR’ refer to their cleavage mechanisms for piperidine as tentative and in need of verification by isotope labeling or high resolution mass spectrometry, Such a spectrum2 under a resolution oiabout 1 in 2500 [Ah/H & 0.011 is presented in Table I.a Mass determinations for the fraglpents which produce the single peaks at m/q 42,43,44, 56, 57, and 70 Table I: Maea Spectrum of Piperidine m/n
26 27
Empirical formula
CaHz CzHs CHN
28
cza
29
NZ(background) CzHs
30
C“
CHzN CHsN
Intensity, arbitrary unitsa
16 44 3 18 80 17 18 64 71 37 3 3 30 10 68 58 84
56 57 -. 70
CAN
3 7 21 4 100 100 26
a Since instrumental adjustments were necessary while these data were being obtained, the relative intensities from one grroup of p& to m t h e r (as indicated by the spacings in the table) should be considered aa only a p p r o h t e l y correct.
prove that the ion formulas proposed for these fragments by B, D, and W are correct. This is equally true for the major peak contributing to m/q 28 and 29 whose structures were apparently chosen by B, D, and W on the basis of the ease with which nitrogen assumes a positive charge in ionization reactions. Note, however, that both m/q 28 and 29 have contributions from hydrocarbon fragments. The fragments (CHr NHCH2)+, m/q 44 and (CHsNH)+, m/q 30 require transfer of hydrogen and the moderately intense hydrogen-deficient hydrocarbon fragments CB7, C&a, and C B a probably result from these abstracting reactions. The hydrocarbon ions could be formed concurrently, with the corresponding nitrogen-containing ion, in competing reactions, or from an initial fragmentation without rearrangement and then subse quent loss of hydrogen. CsHuN+ +CH*N+ C4H7
+
CHiN or
Ca11N+ 4C H a
+
+ C&7+ + C4Hs+
CBa+ H CB7+ There is no evidence in these spectra of the hydrocarbon ion immediately above m/q 55 (or m/q 41) as might be expected from the second reaction mechanism. Some indication of the purity of this sample of piperidine was obtained from a gas chromatographic analysis in which only 0.01% impurity was d e t e ~ t e d . ~ (1) H.Budsikiewics, C. Djersssi, and D. Williams,“Interpretationof Mass Spectra of Organic Compounds,” Holden-Day, Inc., San Francisco, Calif., 1964, pp. 98-102. (2) J. Beynon, “Msss Spectrometry and Its Application to Organic Chemistry,” Elsevier Publishing Go., New York, N. Y., pp. 52, 53. (3) Peaks in parent range and several very weak peaks are not included in the data. The low resolution spectrum is in agreement with that given in “Catalog of Mass Spectra,” American Petroleum Institute Ramarch Project 44, Agricultural and Mechanical College of Texas, College Station, Texas, Spectrum No. 618. Spectrometer used here is a Model 21-110 double focusing instrument, manufactured by Consolidated Electrodynamics Corp. (4) Results obtained by Mr. Robert Grula, CRDL, are much appreciated.
The Behavior of Alkali Metal Cations in the Pores of Silica Gel
by Russell W. Maatman Department of Ch-, Dordt College, Swux Center, Iowa (Received April 1,1966)
Tien interpreted the interaction of alkali metal cations with silica gel in terms of ion exchange, physical
