DISSOCIATION OF LITHIUM AND SODIUM SALTS IN ETHEREAL SOLVENTS
1021
Dissociation of Lithium and Sodium Salts in Ethereal Solvents by D. Nicholls, C. Sutphen, and M. Szwarc Department of Chemistry, State University College of Forestry, Syracuse Unisersity, Syracuse, New York 16220 (Receined September 8, 1967)
The density, viscosity, and dielectric constant of 2-methyltetrahydrofuran (MeTHF) and tetrahydropyran (THP) were determined over a wide temperature range. The significance of the data is discussed. The conductance of the lithium salts of biphenyl, naphthalene, anthracene, and perylene in THF and the conductance of sodium and lithium tetraphenyl borides in THF and MeTHF have also been investigated over a wide temperature range. The conductance data indicate that in each solvent the Stokes radii of solventcoordinated free Lif and Na* ions are nearly the same; however, both are larger in MeTHF than in THF. The dissociation constants of the lithium salts of the aromatic radical anions were found to be essentially independent of the nature of Ar - and their heats of dissociation were low (-3 to 0 kcal/mol). These observations indicate that the salts form solvent-separated pairs in THF over the whole temperature range. The lithium and sodium tetraphenyl borides were also found to exist as solvent-separated pairs in THF, while in h!teTHF it was shown that the proportion of contact ion pairs is significant,particularly in the case of the sodium salt. A comparison of the solvent-separated sodium salts of aromatic radical ions with those of the corresponding lithium salts shows that the latter are 3 4 times less dissociated than the former. This might indicate a higher degree of compressibility of the solvated lithium ion when compared with the sodium ion.
The conductance of various sodium salts has been investigated previously in tetrahydrofuran (THF) and in dimethoxyethane (DAIE). l The present study was initiated with the following aims: (1) to determine the mobilities of alkali cations in THF, 2-methyltetrahydrofuran (MeTHF), and tetrahydropyran (THP), and (2) to determine the heat and entropy of dissociation of sodium and lithium salts in the above solvents, thus gaining insight into the structure of the respective ion pairs in these media. We investigated, therefore, the conductance of the lithium salts of perylene, anthracene, biphenyl, and naphthalene in THF, and lithium and sodium tetraphenyl boride in T H F and MeTHF. The attempt to measure conductance in T H P was unsuccessful in that the data were irreproducible and could1 not be used to determine the relevant dissociation constants which seem to be of the order M . The present data for Na+, BPh4of lo-* to in T H F agree well with those previously reported.lb Extensive knowledge of solvent properties is needed for such investigations. The data characterizing T H F and DME have been reported elsewherelb and the respective values for RleTHF and T H P are reported here.
Experimental Section The purification of the solvents2 and the techniques used in studying their properties have been fully described.'b The lithium salts of perylene, anthracene, biphenyl, and naphthalene were prepared from the recrystallized hydrocarbons which, with the exception of perylene, were sublimed in vacuuo before use. Small pieces of lithium were introduced into the reaction vessel under argon and the system was evacuated and sealed off. The surface of the metal was
scaled by treating it with a T H F solution of the appropriate hydrocarbon. The reacted solution was decanted from the shiny metal into a side ampoule and the pieces washed by condensing on them the solvent distilled from this ampoule. Thereafter, a fresh T H F solution of the same hydrocarbon was contacted with the clean metal surface for 10-15 min. Only a fraction of hydrocarbon reacted and, therefore, the decanted solution contained the radical anions together with an excess of hydrocarbon, but free of dianions. The conductance was measured using the technique described in ref 1. The measurements were performed in the temperature range -75 to 25". All these operations, as well as the previously described preparations, were performed under high vacuum in all-glass equipment with break-seals instead of stopcocks. The concentrations of the lithium salts of the aromatic hydrocarbons were determined spectrophotometrically.1c The respective extinction coefficients were assumed to be the same for the lithium and sodium salts, and the pertinent data justifying this assumption are listed in Table I where similar data of Hoijtinks are also included. Consideration of these data suggests that the extinction coefficients of the lithium salts do not differ by more than 5-6% from those of the respec(1) (a) D. N. Bhattacharyya, C. L. Lee, J . Smid, and .M. Szwarc, J . Phys. Chem., 69, 608 (1965); (b) C. Carvajal, K. J. Tolle, J. Smid, and M . Szwarc, J . Amer. Chem. SOC.,87, 5648 (1965); (0) P. Chang, R. V. Slates, and M. Sawarc, J. Phys. Chem., 70, 3180 (1966). (2) C. Geacintov, J. Smid, and .M.Sawarc, J . Amer. Chem. Soc., 84, 2508 (1962). (3). (a) P. Balk, G. J. Hoijtink, and J. W. H. Schreurs, Rec. Trav. Chzm., 76, 813 (1957); (b) G. J. Hoijtink and P. J. Zandstra, Mol. Phys., 3, 376 (1960); (c) K . H . J. Buschow, J. Dieleman, and G . J. Hoijtink, ibid., 7, 1 (1963).
Volume 72, Number 6 March 1968
D. NICHOLLS, C. SUTPHEN, AND M. SZWARC
1022 Table I Our dataa
Radical anion Na
-
Perylene . Biphenyl Naphthalene
mp-----.
--Xmax,
(259
Anthracene.
(decimal) for the sodium salt e
in T H F
9
-
+
Li +
658
657
7 ,000
577
580
59 ,500
400 323
407 325
38,000 16,000
Hoijtink's data
7
Xmar
e
662(Xa+, 25') 667*(Na+, - 180') 581(Na+, 25") 590b(Na+, - 180') 400b(Li+,25') 324(Na+, 25') 329'(Na+, - 180")
(decimal)
7,250 6,870 43,000 (?) 57,500 36,500 15,500 15,800
Ref
3a 3b 3a 3b 3c 3a 3b
a The extinction coefficients of sodium and lithium perylenee- were redetermined by introducing a known amount of stearic acid to their solutions and measuring the decrease in the respective optical densities. The results show that within 6% the extinction coefficients were the same. At Xmaa = 577 mp, E = 60,000 for perylene . -, Na+ and at Xmax = 580 mp, E = 63,000 for perylene . -, Li +. In 2-methyltetrahydrofuran solution.
tive sodium salts. If this difference is real the corresponding values of A0 would be higher by 5 4 % . Solutions of lithium and sodium tetraphenyl borides were prepared by the gravimetric rather than the volumetric dilution technique. Some experimental difficulties were encountered in determining the conductance of these salts in the lowest concentration range to 10-5 M ) . Apparently adsorption of the salts on glass, particularly that of lithium, led to some errors. Consequently, all the conductance measurements were repeated several times, rinsing the glassware with the appropriate dilute solutions to ascertain the reliability of the calculated salt concentrations in the prepared samples. It was also essential to exclude oxygen from solutions of Li+, BPh4-, since this compound seemed to oxidize readily. Purified helium, therefore, was bubbled through the solvent prior to preparing the solution. The method used in determining conductance is given in ref 1. Fuoss plots (vix., F / A us. c A f 2 / F ) involving the values of A. determined from Ostwald plots were used to find initial values of Ao. The values o of A, thus obtained were plotted in the form of ~ A us. 1/T and the "smoothed" values, derived from such plots, were used in the calculations of the corrected FUOSS plot and of K d i s s .
Table I1 : Physical Properties of 2-Methyltetrahydrofuran (MeTHF)"
25.0 10.0 0.0 -10.0 -20.0 -30.0 -40.0 -50.0 -60.0 -70.0 -75.0 a
d In e/d In T = -1.125; 0.00109 deg-1; log
= a =
The Journal of Physical Chemistry
4.57 5.36 6.01 6.80 7.77 8.98 10.51 12.47 15.04 18.47 20.63
E = - 1.14 7 = -3.635
6.24 6.63 6.92 7.22 7.55 7.91 8.30 8.72 9.19 9.70 9.97
+ (2200/T); d I n v/dT
+ (386/T).
Table I11 : Physical Properties of Tetrahydropyran (THP)"
25 20 10 0 - 10 20 - 30 -40
-
Characterization of the Solvents Properties of MeTHF and T H P are given in Tables I1 and 111. The latter solvent freezes a t about -45" and, therefore, the respective data could not be extended to lower temperatures. The lower density of MeTHF as compared with T H F indicates a poorer packing of RleTHF molecules, probably because of their lower symmetry. Yevertheless, the thermal expansion coefficients of both liquids are remarkably similar. As expected, the dielectric constant and its temperature coefficient ( b In e/b In T ) are lower for MeTHF than for THF.
0.848 0.862 0.871 0.880 0.889 0.898 0.908 0,917 0.926 0.935 0,940
- 45 a
0.878 0.883 0.892 0.901 0.910 0.919 0.928 0.938 0.942
d In e/d In T = -0.97; 0.00104 deg-1; log 7
a =
E
=
= 0.11 -4.10
7.64 8.26 9.73 11.61 14.03 17,21 21.48 27.31
5.61 5.71 5.90 6.12 6.35 6.59 6.86 7.15
31.05
7.30
+ (1640/T); d In v/dT + (59l/T).
=
-~ ~
The density of T H P is higher than that of the isomeric MeTHF; in fact, the former is comparable to that of THF. In its flow properties, T H P resembles dioxane, both liquids being much more viscous than THF, MeTHF, or DME. The nonplanar chair con-
DISSOCIATION OF LITHIUM AND SODIUM SALTS
IN
ETHEREAL SOLVENTS
formation of T H P molecules probably leads to substantial stacking in the liquid which accounts for the observed high viscosity and its high activation energy (2.7 kcal/mol for THP, as compared with 1.8 kcal/mol for T H F and IIeTHF). I n spite of their isomerism, €25 = 5.61 for THP, whereas its value is 6.24 for RleTHF. Even more significant is the low temperature coefficient of the dielectric constant, viz., - b In e/b In T = 0.97. The relevant coefficients for I l e T H F and T H F have higher values, viz., 1.125 and 1.16, respectively. The relevant data for T H F are given in ref lb.
BIPHENYL- No'
-
-5.4
-
-5.6
-
-5.8
-6.0
The plots of 7A0 us. l / T are given in Figure 1 for Li+, naphthalene. --, Li+, anthracene. -, and Li+, perylene.-. The maximum experimental scatter of the initial values of A, is within 10% and it seems likely, therefore, that the smoothed values of AO are reliable within 5% in all cases. The lithium salts decompose slowly in tetrahydrofuran, but it seems that the products of decomposition did not affect our results. To ascertain this point, we determined the conductance of lithium anthracene twice, using independently prepared solutions, and the agreement between these two sets of data was most satisfactory (within 5%). The same conclusion was drawn from the studies of the conductance of lithium perylene. The conductance of a series of solutions was measured consecutively as the concentration was decreased. In the course of these determinations, the lithium salt was partially destroyed by spontaneous reaction with the solvent. Thereafter, the solution was concentra,ted and the conductance redetermined. The results agreed with the previous set of data within 2%, indicating no discernible effect of the products of decomposition on the conductance.
-4.8 -5.0
-
-5.2 -5.4
NAPHTHALENE' LI'
-
-5.6
-
-5.0
-
LOG KDISS - 6 . 2
I
d
-
-4.6
-6.0
Conductance of the Lithium Salts of Radical Ions Derived from Perylene, Anthracene, Naphthalene, and Biphenyl in THF
1023
-7.0
1
I
-
m
-
ANTHRACENE- LI+ -5 4
-4 4
PERYLENE- No'
35
3.0
4.5
4.0
5.0
IO)/T
Figure 2. Plots of log radical ions in THF.
&iss
us. 1 / T for salts of
Comparison of the smoothed Aots obtained for the lithium and sodiumlCsalts of perylene. -, anthracene. -, and naphthalene.- show that they are virtually identical (see Table IV), indicating that in T H F Ao+(Sa+) = Ao+(Li+),contrary to the earlier findings from this laboratory.'" This identity was shown to be valid in the temperature range -60 to 25'. ~~
o.7
"
ANTHRACENE^ NO+
~
Table IV : Comparison of Ao ( cm2/ohm) for Sodium and Lithium Salts in THF at 25'
t 0 0
Radical ion
-
Perylene . Anthracene Naphthalene * -
ha(Ns+ A r * - ) a
106 120 128
ha(Li+ Are-)
107 Yj 121 f 6 129 f 6
' Obtained fiom ref IC. 0.6
0.5
PERYLENE- L i t
0.4
3.0
3.5
4.0
4.5
i 031 T
Figure 1. Plots of the Walden product of lithium salts of radical ions in TlKF v5. 1/T.
1.0
The large degree of destruction of lithium biphenyl in solutions more dilute than 5 x 10-6 M made it impossible to obtain points a t lower concentrations. Consequently, the lack of data prevented an accurate determinat'ion of the relevant ho's. I n the computation of Fuoss plots, it was assumed, therefore, that Ao(B.-, Volume 76,Number 9 March 1068
D. NICHOLLS, C. SUTPHEN, AND 14. SZWARC
1024 Table V : Dissociation Constants of Lithium Salts of Radical Ions in THFavb
T,OC
25 I5 5 -5 - Ifj - 25 - 35 - 45 - 55 - 65 -75
-Biphenyl Slope of Fuoss line
*
-, Li
M
-Naphthalene Slope of Fuoss line
4.2 4.7 5.3 5.6 6.5 6.8 7.0 7.6 8.0 7 . 6 (?) 9.0
19.1 19.7 21.3 23.2 25.5 30.6 37.3 40.7 62.1 87.7 127
+ -
1O5Kdim,
13.5 14.8 16.8 19.8 22.4 28.4 37.0 48.5 66.2 103 135
-, Lit-lO'Rdiss,
M
3.1 3.8 4.3 5.2 6.0 6.7 7.4 8.2 9.0 9.4 10.3
-Anthracene Slope of Fuoss line
15.7 17.0 19.8 23.7 27.5 37.3 43.7 61.0
*
-, Lit---
19.3 21.3 24.9 29.9 36.4 46.3 57.8 82.0 115 189 263
4.4 4.95 5.3 5.6 6.3 6.2 7.2 7.2
...
. I .
133
7.1
*..
...
a The supporting data necessary in calculation of Fuoss plots are given in ref 4. within 10-1570.
---Perylene Slope of Fuoss line
10'Kd i 98s M
*
-, Li '-10'Kdiss,
M
4.5 5.0 5.4 5.7 6.1 6.4 7.0 6.8 7.0 6 . 4 (?) 7.1
' The calculated dissociation constants are reliable
-~
Table V I : Heats (kcal/mole) and Entropies (cal/deg) of Dissociation of Lithium and Sodium Salts of Radical Ions in THF" -------Biphenyl. Li t
AH(l5') A S ( 15 " ) AH(-65') A S ( -65') a
-2.5
- 33
-0.4
- 26
Na i-
-Naphthalene* Lii-
-7.3
- 52
-1.6
-28
-2.8
- 35
-0.5
-26
Na +
-8.2
- 58
-1.8
- 30
--Anthracene Li +
*
---
Na
-1.9 0.0
- 24
Na i-
f
-6.1
- 31
----
YPerylene Li
+
-1.4
-45
- 29
0.0 -21
- 24
-2.2
-29
0.0
0.0 -21
AH values are reliable within 0.3-0.5 kcal/mol, the corresponding A S values within 1-2 eu.
Lif) = Ao(B.-, Na+). The final calculations led us to the values of the respective dissociation constants given in Table V and presented graphically in Figure 2. The calculated heats and entropies of dissociation are given in Table VI, together with the corresponding data for the sodium salts. The striking equality of the Ads derived for the respective lithium and sodium salts is most significant. It proves that the mobilities of the Li+ and Xa+ ions in tetrahydrofuran are almost identical, Le., the solvated ions have almost the same Stokes radii. Experimental determination of A0 in ethereal solvents is difficult and an error of 5% is not uncommon. Hence the errors in lo-and lo+ may be of the order of 10-20%. However, our present work permits us to claim that Xo+(Li+) = Ao+(Na+) within 5%) since we compare a series of lithium and sodium salts with common anions. The lithium salts in solvents of low dielectric constant might be agglomerated. This appears not to be the case for the salts of the investigated aromatic radical ions in tetrahydrofuran. A typical Fuoss plot shown in Figure 3 demonstrates that the linear relation between P / A and cAf2/F is well obeyed in the concentration range to 7 X iM. This should not be the case if the dissociation is represented by the equation (Ars-, Li+),
J-nAr.-
The Journal of Physical Chemistry
+ nLi+
J 0
IO
20
30
40
f 2 c A / Fx IO'
Figure 3. A typical Fuoss plot; Li+, anthracene*- in THF at + 5 " .
Alternatively, one could argue that. the dissociation of agglomerates follows the equation (Ar . -, Li+), 1_ Li+
+ (Ar. -,
Li+,-l)
but then the value of A. for AS-, Li+ should be substantially smaller than the respective AO of A * -t Naf, contrary to our observation. There is no evidence whatsoever for any aggregation of sodium salts, and the remarkable agreement between the A0 values of the lithium and the corresponding sodium salts shown in Table IV implies that the transport in both systems is due to single ions. The dissociation constants of the sodium salts of radical anions were shown to be virtually independent
DISSOCIATION OF LITHIUM AND SODIUM SALTSIN ETHEREAL SOLVENTS of the nature of the anion if the pair is of the solventseparated type.lO This is illustrated by the data obtained in dimethoxyethane and by those in tetrahydrofuran at low temperatures. The heat of dissociation of such pairs is low, viz., between 0 and -2 kcal/mol. Our present results show the same behavior for the lithium salts in tetrahydrofuran. As may be seen from the data given in Tables V and VI, the dissociation constants of all the lithium salts are virtually identical and their heats of dissociation are all low. We conclude, therefore, that the lithium salts in tetrahydrofuran are virtually solvent separated, at least at temperatures below 25". Our present data rationalize the observations of Hoijtink, et CLZ.,~ who found the conductance of a M solution of lithium anthracene to decrease monotonically with decreasing temperature. The temperature coefficient of conductance is given by E, '/ZAHdiss, and for anthracene. -, Li+ in THF, its value increases from 0.9 kcal/mol at 15" to about 1.8 kcal/mol at -65". The relation between the dissociation constants of the sodium and lithium salts in tetrahydrofuran may be seen from inspection of Figure 2. The sodium salts of naphthalene - and biphenyl. - form contact pairs at higher temperatures ( - A H is high) and, consequently, the respective dissociation constants are lower than those of the corresponding lithium salts. A smaller fraction of contact pairs present at temperatures above 25' may account for the behavior of anthracene Na+. Under conditions when the solvent-separated pairs dominate, the lithium salts are less dissociated than the sodium salts. Since the solvated free Ka+ and Li+ ions are approximately of the same size, our observation tentatively indicates that the solvation shell of the lithium ion is more compressible than that of the sodium ion. It is also possible that the solvated lithium ion is slightly smaller than the sodium cation, since a difference of, say, 0.2 in the Stokes radii might not be detectable in values of A, which are uncertain within 5%. This is plausible, since the unsolvated Li+ cation is smaller than the Na+ cation. It is obvious, however, from the dissociation constants and the respective AX values that the solvated lithium ion cannot be larger than the sodium cation. The conductance data for Are-, Li+ are in good agreement with those reported for lithium fluorenyl in tetrahydrofuran.6 This may be seen from Table VII. The slightly higher -AHdiss found for the fluorenyl salt could be expected; the more localized negative charge in fluorenyl carbanion when compared with anthracene - should lead to a higher proportion of the contact pairs in the former case. Hence, the dissociation constant at 25" of fluorenyl lithium should be lower than that of lithium anthracene, in agreement with the observations. These results suggest, therefore, that fluorenyl lithium is not agglomerated in tetrahydrofuran, contrary to recent claims.8
+
-
-)
T,'C
1025
Li+. anthracene. -
Li+, fluorenyl-
25 0 - 30 60
3.9" 6.6a 12" 19"
-
1.7' 2.gb 5.2' 8. tib
4.5 5.5 6.7 7.2
a The original results, calculated on the basis of the assumed & = 88.4 at 25'. * The recalculated values, based on the A0 = Ao(anthracene*-, Li+) = 121 at 25'. In view of similar sizes of A . - and fluorenyl, this identity is plausible.
Conductance of Li+, BPh4-, and of Na+, BPh4in THF and MeTHF The values of A, and K d i s s of Li+, BPh4- in TFIF and MeTHF are given in Table VIII, and the corresponding data for Na+, BPh4- are given in Table IX. I n view of the difficulties discussed previously, some of the quoted A0)s may be uncertain within 15%. It is desirable, therefore, to look for some regularities which could indicate the reliability of these data. Indeed, one notes a monotonic change with temperature ~ ~ for e ~the~ lithium ~ of the ratios (AOV)THF:( A O ~ )both and sodium salts (see the last columns of Tables X and XI. Stokes radii of the solvent-coordinated free Na+ and Li+ ions appear to be similar, confirming our previous observations. I n THF, the Na+ cation seems to be slightly larger (6-7%) than the Li+ ion, whereas the relation appears to be reversed (- 5%) in MeTHF. These deviations may reflect the experimental uncertainty. Both ions appear to be larger in RleTHF than in THF. This is plausible; the coordination shell is expected to be larger when MeTHF replaces THF. Most interesting are the data dealing with the heats of dissociation of the respective ion pairs. The dissociation constants, determined at various temperatures,
-
loa Koua.
No+, THF Li*, THF Na*, MeTHF
Li',
Me'lHF
-5.0
1031~ 3.5
4.0
4.5
5.0
Figure 4. Dissociation of tetraphenyl borides. (4) K. H. J. Buschow, J. Dieleman, and G. J. Hoijtink, J . Chem. Phye., 42, 1993 (1965). (5) T. E. Hogen-Esch and J. Smid, J. Amer. Chem. Soc., 88, 318 (1966). (6) T.E. Hopen-Esch and J. Smid, ibid., 89, 2764 (1967). Volume 78, Number 9
March 1968
D. NICHOLLS, C. SUTPHEN, AND M. SZWARC
1026 Table VIII: Dissociation Constant of Li+, BPha-" -THF
a
MeTHF
r
7
A@,
T,OC
cma/ohm equiv
Slope
10'Kdiasr M
An, cml/ohm equiv
25.0 15.0 5.0 -5.0 -15.0 -25.0 -35.0 -45.0 -55.0 -65.0
90 82 73 64 56 48 41 34 28 22
2.39 2.86 3.34 4.14 4.64 5.80 7.55 10.05 12.0 17.5
5.1 5.3 5.8 6.0 6.9 7.5 8.0 8.8 10.8 11.5
77 68 60 52 45 38 32 26 21 17
A. values reliable within 15%;
&is8
values within 30y0. The variations in
Table IX : Dissociation Constant of Na+, BPh4-
&is8
THF
14.6 15.1 16.5 19.0 21.2 25.8 31.4 41.2 52.8 80.2
1.2 1.4 1.7 1.9 2.3 2.6 3.I 3.5 4.1 4.3
4.4 3.7 3.3 3.0 3.0 2.8 2.6 2.5 2.6 2.7
are more reliable than its absolute value.
MeTHF
A@ ,If
cml/ohm equiv
Slope
1O6Kdiie, M
Rdias(THF) : KdisdMeTHF)
86 76 67 59 52 44 38 32 26 20
1.52 1.73 2.00 2.38 2.85 3.56 4.47 5.88 11.5 15.4
8.8 10.0 11.0 12.0 13.2 14.4 15.5 16.6 17.0 16.2
80 71 63 55 47 40 34 28 23 18
12.5 12.0 12.0 12.2 12.8 15.2 17.4 22.6 31.0 43.5
1.2 1.6 2.1 2.7 3.5 4.0 5.0 5.6 6.3 7.0
7.1 6.1 5.2 4.4 3.8 3.5 3.1 2.9 2.7 2.3
-
' AO values reliable within
15%;
10'Kdisav
Kdisswithin
Table X : Walden Products, Aoq, of Li+, BPh4-
30%. The variations in
Kdies
THF
MeTHF
Ratio
T," C
25 15 5 -5 - 15 25 35 - 45 55 65
0.417 0.418 0.417 0.415 0.412 0.408 0.403 0.396 0.390 0.382
0.350 0.345 0.339 0.332 0.326 0.319 0.311 0.302 0.294 0,284
1.19 1.21 1.23 1.25 1.27 1.28 1.29 1.31 1.33 1.35
25 15 5 -5 15 -25 -35 -45 - 55 - 65
THF, 1.09; MeTHF, 1.23.
are listed in Tables VI11 and IX, and the plots of log vs. 1/T are shown in Figure 4. The heats of dissociation in T H F are -1.2 f 0.9 kcal/mol for Li+, BPh4- and -1.3 f 0.9 kcal/mol (decreasing, however, a t lower temperatures) for Na+, BPhd-. This implies that both ion pairs are solvent separated, being coordinated with THF, and dissociate into T H F Kdiss
The Journal of Physical Chemistrg
are more reliable than its absolute value.
Table XI: Walden Products,
a
T ,OC
Ratio (25'/-65"):
KdiSs(THF): KdidMeTHF)
Slope
-
a
lo6Kdiss, M
oma/ohm equiv
25 15 5 -5 - 15 25 35 -45 -55 65
-
Slope
'
Aa
T,'C
7
-
a
THF
0.397 0.389 0.387 0.383 0.381 0.376 0.377 0.375 0.365 0.343 ( 1 )
Ratio (25"/-65"):
Aoq,
of Na+, BPB-
'
MeTHF
0.367 0.362 0.356 0.350 0.343 0.336 0.328 0.320 0.310 0.300
Ratio
1.08 1.08 1.09 1.10 1.11 1.12 1.15 1.17 1.18 1.14(?)
THF, 1.16; MeTHF, 1.23.
coordinated alkali ions. Hence, the entropies of dissociation should be relatively small; in fact, they are -24 and -23 eu, respectively. In RIeTHF the heat of dissociation of the lithium salt is also relatively low, viz., -2.2 f 0.9 kcal/mol, the exothermicity being higher than in T H F only by 0.9 kcal/mol. However, a substantial increase of exothermicity is
GAS-PHASE RADIOLYSIS OF BENZENE
1027
observed for the sodium salt, uiz., AH = -3.8 f 0.9 kcal/mol, decreasing to somewhat lower values at - 70". Apparently, the degrees of solvent coordination of the Li+, BPh4- pair and Li+ free ions are still comparable in MeTHF. The larger Na+ ion, which also appears to be coordinated by MeTHF when free, associates into a noncoordinated, contact Na+, BPh4ion pair. The latter becomes gradually more solvated as temperature decreases, this being reflected in the
respective value of AH (see Figure 4). The decrease of the value of - AH at lower temperatures is a general phenomenon observed in many other systems.' Acknowledgment. We wish to acknowledge financial support of this investigation to the National Science Foundation, to the United States Air Force (Contract AF 33(615)-3788))and to the Petroleum Research Fund (Grant 2475-C) administered by the American Chemical Society.
Gas-Phase Radiolysis of Benzene by Robert R. Hentz and Stefan J. Rzad Department of Chemistry and the Radiation Laboratory,l University of Notre Dame, Notre Dame, Indiana 46666 (Received September 11, 1067)
The radiolysis of benzene vapor was studied a t room temperature as a function of pressure (5-80 torr), dose, dose rate, and added NO and NzO. G( C&) and the yields and variety of ring-fragmentation products are considerably greater than in radiolysis of the liquid. At pressures greater than -60 torr, product yields are independent of pressure. At 66 torr, values of G(H2) = 0.11 and G(C2HZ) = 0.63 are obtained which are independent of dose, dose rate, and the presence of NO or N2O. It is concluded that, at >60 torr, HBand CzH2are not formed via free-radical or ion-neutralization reactions, but are formed by molecular elimination from either superexcited states or ion-molecule reactions of CaHe+ with CBH, or both. G values of the following products also were studied : CHI, C2H4, propadiene, butadiyne, cyclohexene, cyclohexadienes, biphenyl, and dihydrobiphenyls. The yields of all such products show, in addition to a pressure dependence below -60 torr, a dependence on certain of the other radiolysis variables; free radicals or secondary ions or both are involved a t some stage in their formation. At the lowest dose studied, except for G(bipheny1) = 0.17, all such products are formed with G < 0.1. At a total pressure of 66 torr, radiolysis of benzene in the presence of NzO yields NZand phenol in a 2: 1 ratio; G(N2) = 13.3 at 10 mol % NzO appears to require some kind of short-chain mechanism.
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Introduction Considerable attention has been devoted to study of the radiation chemistry of benzene in the liquid phase.2 However, no detailed study of the gas-phase radiolysis has been published, although several authors have reported G values for a few products under a specific set of conditions,3-'J Indeed, isopropylbenzene appears to be the only aromatic hydrocarbon whose gas-phase radiolysis has been studied in any detail.' The present study of benzene radiolysis in the gas phase was undertaken for comparison with a recently completed study of the gas-phase photolysis of benzene with the X e and Kr resonance lamps.* Such a comparison is of interest because negligible ionization of benzene occurs with the Xe lamp, while the Kr lamp produces ionization in benzene with a quantum efficiency of -6O%.@ Experjmental Section Materials. The benzene used was Fisher Certified Reagent; it was purified with an Aerograph Autoprep
Model A 700 using a 3/~-in.X 12-ft column which was packed with 20 wt % P,P'-oxydipropionitrile on 60-80 mesh Chromosorb W and operated at 70". Benzene of purity better than 99.99 mol % was obtained. Nitrous oxide and nitric oxide, obtained from Matheson, were purified by low-temperature distillation. (1) The Radiation Laboratory of the University of Notre Dame is operated under contract with the U. S. Atomic Energy Commission. This is AEC Document No. COO-38-542. (2) For a. summary of the resultsof such studies, see E. A. Cherniak, E. Collinson, and F. 9. Dainton, Trans. Faraday SOC.,60, 1408
(1964). (3) J. P. Manion and M. Burton, J. Phys. Chem., 56, 560 (1952). (4) V. P. Henri, C. R. Maxwell, W.C. White, and D. C. Peterson, ibid., 56, 153 (1952). (5) P. Huyskens, P. Claes, and F, Cracco, Bull. SOC.Chim. Belges, 68, 89 (1959). (6) L. M.Theard, J . Phys. Chem., 69, 3292 (1965). (7) R. R. Hentz, ibid., 66, 1622 (1962). (8) R. R. Hentz and S. J. Rzad, ibid., 71, 4096 (1967). (9) J. C. Person, J. Chem. Phys., 43, 2553 (1965). Volume 72, Number 3 March 1068
