4124
R. V. SLATESAND M. SZWARC
Dissociative Equilibria in the Systems
+ Na+
Aromatic HydrocarbonT,Na+1_ Radical Anion'
by R. V. Slates and M. Szwarc Department of Chemistry, State University College of Forestry at Syracuse University, Syracuse, New York 13,910 (Received June 9, 1966)
The conductance of the sodium salts of aromatic radical ions has been investigated in THF. The dissociation constants of the respective ion pairs were determined and found to increase with the size of the anion, being 1.5 X M for biphenyl-,Na+ and 23 X ill for perylene-,Na+. An exceptionally low value was found for naphthalene-,Na+, viz., -0.2 X M . The A0 and A,- values were determined, and their significance is discussed. The Ao- values derived from the conductance studies are slightly smaller than those calculated from the diffusion constants of the respective aromatic hydrocarbons, and thus the diffusion of an aromatic radical ion appears to be only slightly slower than that of the parent aromatic hydrocarbon. Knowledge of K d i s s allows us to correct the values of the electron affinities of aromatic hydrocarbons which were previously determined potentionietrically and spectrophotometrically. The relation between the equilibrium constants derived by these two methods is discussed in terms of the equilibria B & BA (K,t) established between ion pairs and those involving free ions, viz.,Aas compared with AT,?tI+ €3 a B-,M+ A (I&,). The two constants are related, namely, Ks,/KPt = K d i s s , * - ,>I + / K d i a s , ~ ?, M +. A similar problem arises when one considers the disproportionations of radical anions into dianions since the following equilibria have A A2-; AT A-,Na+ & A2-,Na+ A; and 2A-,Na+ a to be considered: 2AA2-,2Xa+ A.
+
+
The chemistry of radical anions derived from aromatic hydrocarbons and related compounds is now rapidly developing, I n solution these species exist in a t least two distinct forms' as ion pairs and as free ions. Other associates such as contact ion pairs and solvent-separated pairs,'b triple ions, dimeric ion pairs,2 etc., may also be encountered. Since each of these species has its own characteristics and usually differs from the others in its reactivity, it is necessary to learn more about the equilibria established between them in order to interpret correctly their behavior in various processes. I n this work we shall deal with the dissociative equilibria, ArT,Na+ ArT Na+, established at 25" in tetrahydrofuran. The following aromatic hydrocarbons were investigated : biphenyl, naphthalene, triphenylene, pyrene, anthracene, tctracene, and perylene. The results reported here were obtained from conductance studies of their sodium
+
The Journal of Physical Chemistry
+
+
+
+
+
+
salts. However, before considering the experimental details of this work, the significance of such studies will be discussed in relation to other experiments involving radical ions. Two examples have been chosen : (1) electron-transfer equilibria involving aromatic hydrocarbons and (2) disproportionation equilibria of radical anions.
Electron-Transfer Equilibria Following the pioneering work of Paul, Lipkin, and ~~~~~
(1) (a) H. V. Carter, B. J. McClelland, and E. Warhurst, Trans. Faraday SOC., 56, 465 (1960); (b) A. C. Aten, J. Dieleman, and G. J. Hoijtink, Discus,a'ons Faraday Soc., 29, 182 (1960); (c) B. J. McClelland, Trans. Faraday SOC.,57, 1458 (1961); (d) N. M. Atherson and S. I. Weissman, J. Am. Chem. SOC.,83, 1330 (1961); (e) E.de Boer and E. L. Mackor, %%id., 86, 1613 (1964); (f) J. Bolton and G. K. Fraenkel, J. Chem. Phys., 40, 3307 (1964). (2) (a) K. H. J. Buschlow, J. Dieleman, and G. J. Hoijtink, ibid., 42, 1993 (1965); (b) K.H. J. Buschlow, Ph.D. Thesis, Amsterdam, 1963.
+ Na+
EQUILIBRIA I N SYSTEMS O F AROUATIC HYDROCARRONT,Na+ & RADICAL ANIOF-
4125
Table I
(1) Xaphthalene (2) Terphenylene
O4.85 e2
(1) Pyrene (2) Anthracene
6.8 4.55
( I ) Tetracene (2) Perylene
3
0.029
0. 047b
0.095
0.000
0.041
0.29
111
0.125
0.119
0.114
0.124
1.5
1.6
}
52
0.102
0.108
0.113
0.103
0.65
0.62
(2.2)"
91
0.117
0.102
23 l5
(1) Pyrene (2) 9,lbDimethylanthracene
1
}
...
(0.087)~
3.2
-
The values given in the last column are calculated from log (Kdiss.l/Kdiaa.z) = ( A a l , ~ . ~ 4~ ~ l , ~ , ~ ~ ~ . d ) / 0* .This 0 3 . va.lue is the least reliable. It involves a substantial correction (see ref. 4). These values are calculated from the I(di88,l/Kdiea,2 given in the last column.
Weissman, the electron-transfer equilibria between aromatic hydrocarbons and their radical anions were investigated spectrophotometrically in tetrahydrofuran4 (THF). The investigated processes involved essentially ion pairs, viz. ArlT,Na+
+ Arz Z Arl + ArzT,Na+
(Ksp)
because in the to M THF solutions used in those studies the concentrations of free ions were negligible. The pairs naphthalene-triphenylene, pyand rene-anthracene, pyrene-dimethylanthracene, tetracene-perylene were chosen since their spectra do not overlap, and the respective equilibrium constants are not too large. The results are listed in Table I, and the respective spectrophotometric equilibrium constants are denoted by KEp. I n the same paper,4 the results of potentiometric titrations of aromatic hydrocarbons were reported, being used. the technique developed by Hoijtink, et Thus, the differences in the reduction potentials, A E O I , ~ , were determined. These are related to the respective equilibrium constants Kpt Arr
+ Arl
Arz
+ Arz-
(K,t)
by the conventional equation AE'~,z = (RT/23,000) In Kpt. It should be noticed that K,, and KPt differ from each other since one refers to a process involving ion pairs and the other to free ions. They are, however, related through the equation K,,/K,t = Kdiss,l/ Kdiss,z, where Kdiss,I and Kdiss,Z are the respective dissociation constants of Ar17,Na+ I_ firl-
+ Na+
&ies,l
(1)
and ArzT,Na+
Arz-
+ Na+
Kdiss,z
(2)
Moreover, the directly measured potentials, Le., those established between t,he electrodes when half of the titrated hydrocarbon is reduced, differ from the standard potentials Aeo. As was shown previ~usly,~ the latter are related to the former through the equation AE0i,z =
A€l,Z,rnessd
- ('/zRT/23,000) In (Kdiss,l/Kdiss,d
Hence, the following relations apply a t 25" Ael,Z,sp
=
0.06 log Ksp =
Ael,z,rneasd
0.03 log (Kdiss,l/Kdiss,Z) AE'I,Z =
AEl,Z,meesd
- 0.03 log (Kdiss,l/Kdiss.Z)
These potentials, as well as the respective Kdiss, are also given in Table I. The agreement between the experimental data obtained by the two different techniques becomes better when the dissociation processes are accounted for. This is clearly seen from Table I when the observed and the calculated ratios of the dissociation constants (the last two columns) are compared or when A E ~ , =~ ,0.06 ~ ~ log Ksp (the fourth column) is compared with Ael,z,sp (the seventh column), the latter being derived from the A ~ l , z , m e s s dby means of the above given equation. (3) D. E. Paul, D. Lipkin, and S. I. Weissman, J . A n . Chem. SOC., 78, 116 (1956). (4) J. Jagur-Grodainski, M. Feld, S. L. Yang, and M. Sswarc, J . Phys. Chem., 69, 628 (1965).
(5) G. J. Hoijtink, E. de Boer, P. H. van der Meij, and W. P. Weiiland, Rec. truv. chim., 75, 487 (1956).
Volume 69,Number 12 December 1966
R. V. SLATES AND 13. SZWARC
4126
It should be stressed that the values of Aeol,z, given in the sixth column of Table I, represent the potentials which should be considered when theoretical electron affinities of aromatic hydrocarbons in THF are discussed. For example, the difference of Aeo1,2 values of anthracene and dimethylanthracene, viz., 0.027 v., is more plausible than the small difference derived from the spectrophotometric data (0.008 v.) or even from a direct titration. The interaction of the anion with its counterion is not given by a constant term, and therefore it distorts the results obtained by either spectrophotometric or potentiometric methods. It is also interesting to notice that the dissociation constant of dimethylanthracene-,Na+ (calculated from the dif, ~ ~A e l , ~ , ~ist ) lower than that ference between A e l , ~ and of anthracene‘,Na+. This is to be expected since the electron-donating power of the methyl group should increase the attraction between the anion and its counterion, and apparently the steric effect of the CH3 groups is not large. Finally, the exceptionally low value of Kdiss for naphthalene-,Na+, which is deduced from the conductance studies, is, at least partially, justified by the large ~ ~ Ae01,z. difference between A e 1 , 2 , and Disproportionation of Radical Anions The disproportionation processes Ar-
+ Ar-
Ar2-
+ Ar
(3)
have been discussed by many workers, initially by Hush,6 and later by Hoijtink2J’ and others. At higher dilution, the process ArT
+ Ar-,M+
Ar2-,M+
+ Ar
(4)
may become significant, whereas reaction 3-which is the most interesting theoretically-may hardly be observed. Under normal experimental conditions the reaction involves ion pairs, and hence the observed spectrophotometric equilibria refer to Ar-,M+
+ Ar-,M+
Ar2-,2RI+
+ Ar
(5)
The equilibrium constants Ks, Kq, and Ks are correlated with the dissociation constants ArT,M+ Ar2-,2M+
ArT f R/I+
Kdlss,i
2Ar2-,RI+ + M +
Ar2-,M+
Ar2- -I-JI+
namely
KqIK.5
K‘diss,2 K”diss,Z
Spectrophotometric studies are usually concerned with Ar- = Arrtotal, and Ar2-,2M+ the sums Ar-,Na+ Ar2-,M+ Ar2- = Ar2- total. Hence, the apparent spectrophotometric equilibrium constant for disproportionation, K,,, is given by
+
+
Ksp
=
Kb(1
+
K’dis,.p/M+)/(1
+
Kdiea,l/M+)2
where M+ denotes the concentration of the counterions in the investigated solution. For a very low degree of dissociation K,, = &. The potentiometric titration gives the difference between the first and second reduction potential, (e’ - e”),t, which is related to the potential of reaction 4, (E‘ - e”)4 = -0.06 log K4, viz. (E’
-
d’)4
=
(El
- E’l)pt
0.06 log
+
(Kdiss,1/K’disa,2)(
[M+]Z/[M+II)
Similarly (e’
- d’)3
= (e’
0.06 IOg
- eII),t
+
(K2dise,l/K1diss,2K”dis*,2) ([M+]Z/[M+II)
[M+]I and [M+]2 denote the concentrations of the counterion in the titrated solutions. It is obvious, therefore, that the theoretically interesting difference, (e’ - e”)Q, may substantially differ from that obtained in a potentiometric titration. [(e’ - e’’)pt is the potential established between two electrodes, one immersed in a solution containing an equimolar amount of Ar and Ar-,M+, the other maintained in a solution of equimolar amounts of Ar-,M+ and Ar2-,2M+.]
Results of the Conductance Studies The conductances of the monosodium salts of the following hydrocarbons were investigated : biphenyl, naphthalene, triphenylene, pyrene, anthracene, tetracene, and perylene. I n addition, the conductance of the disodium salt of anthracene was also studied. All of the investigations were carried out in THF, a t 25”, using an apparatus described in a previous communication.’ The concentrations of the radical ions were determined spectrophotometrically, the optical cells being part of the apparatus. The technique of dilution is also described in ref. 7, and it suffices to say that this was accomplished by removal of the solute and not by addition of the solvent. Thus, salt sohtions as dilute as M could be prepared without introducing any impurities or foreign materials. All of the investigated hydrocarbons were twice
= K’dias,z/Kdiss,l
and
The J O U T of ~ LPhysical Chemistry
(6) N. J. Hush and
J. Blackledge, J . Chem. Phys., 23, 514 (1955). (7) D. N. Bhattacharyya, C. L. Lee, J. Smid, and M. Szwarc, J . Phys. Chem., 69,612 (1966).
EQUILIBRIA I N SYSTEMS O F AROMATIC HYDROCARBON-,Na+
crystallized and then sublimed under high vacuum by gentle heating. To prevent any contamination the sublimed material was never exposed to air. The solvent (THF) was purified by a standard procedure. The purified material was kept on a high vacuum line over sodium-potassium alloy to which some anthracene or benzophenone was added, and eventually it was distilled directly, whenever required, to an appropriate ampoule. The solutions of the radical anions were prepared by leaving overnight a T H F solution of the investigated hydrocarbon in contact with a sodium mirror. I n this way biphenyl, naphthalene, and triphenylene were converted into Ar',Na+ (the first one being only partially converted). The remaining hydrocarbons, however, yield under these conditions the disodium salts, Ar2-,2Na+. The following procedure was therefore adopted when dealing with the latter compounds. Their solution was divided into two nearly equal portions; the smaller one reacted with sodium and then mixed with the other. The Ar2-,2Na+ was converted then into Ar-,Na-, leaving in the solution a small residue of the unreacted hydrocarbon. The reliability of the dilution technique was checked for sodium tetracene. The solution contained some excess of tetracene, and its concentration was determined by the optical density at Amax 444 and 474 mp, whereas the concentration of tetracene' was determined by the optical density a t the Amax 711 mp. The results in Table I1 were obtained. These data clearly Table I1
Concn., M
Diln. on basia of tetracene 7
Diln. on baaia of tetracene
3.54
3.50
7.90
7.97
Solution left overnight after being diluted 5 . 7 X lod6 (Previous 4.27 (Killeddur- 2.67 day) ing the night)
0.66
x
10-6 lO-7
2.04
2.03
indicate the correctness of the spectrophotometric determination of the tetraceneT concentration and show that a negligible killing takes place within a few hours. However, the extremely dilute solutions were partially
e RADICALANION^ + Na+
4127
destroyed by the solvent if left for 12 hr. or more. Therefore, only a concentrated solution was kept overnight, and it was stored in a refrigerator. Killing produces a new type of ion pair, e.g., X(CHZ)dO-,Na+, which may contribute to the conductivity of the solution; however, usually its contribution is negligible. The equivalent conductances and t,he concentrations of the investigated solutions are given in Table 111,and the calculated Fuoss lines are shown in Figures 1 and 2. Their slopes, which give l/AO2&iaa, are well determined and reliable within 3-575,. However, the intercepts are less reliable, particularly for the less dissociated radical ions. Since the latter give the values of l/&,the accuracy of &iss is affected by the reliability of the intercept. The final results are collected in Table IV.
Table
IIIa
cx
106, M
C X lOS, M
A
A
Biphenyl+,Na + 1120 4.85 398 7.69 128 12.38 41.3 20.48 15.5 30.99 6.28 44.33 0.541 104.25
Tdphenyl',Na + 1850 7.200 467 12.696 187 18.286 30.9 38.556 16.8 49.871 11.14 58.870 3.87 60.238
Pyrene',Na + 1540 14.05 290 26.0 68.3 44.3 20.5 65.8 12.4 78.5 2.51 120.6
Perylene',Na+ 865 20.22 263 29.00 84.8 42.26 23.1 61.40 8.46 76.01 3.90 85.28 0.95 89.95
Anthracene?,Na + 8150 6.01 1630 9.24 336 16.18 317 17.40 75.6 29.36 15.8 54.55 13.2 56.38
TetraceneT,Na + 1600 13.631 451 20.99 194 29.21 57.2 42.98 13.4 64.75 6.57 75.73 3.00 85.73
Naphthalene-,Na + 554 2.27 225 3.32 85 5.17 35 8.10 15.6 11.38 7.7 15.60 4.7 17.84
Anthracene2-,2Na + 206 2.5 51.5 5.4 25.2 7.6 15.4 11.6 6.55 15.74 6.11 17.00 4.11 14.58
Solvent, T H F ; temperature, 25".
Volume 69,Number 18 December 1966
4128
R. V. SLATESAND M. SZWARC
Table IV bo,
Concn. range,
cm.Z/ohm
1v
Radical anion
0.5 x 4x 4x 2.5 x 1.3 X 3X 1x
Rip henylT NaphthaleneTrip henylerie' Pyrene' AnthraceneTctracene' Per ylene'
Anthracene*-
equiv.
io-~x i 10-3 i o - ~ - jx 10-4 10-6-2 x 10-3 10-61.5 x 10-3 10-'-8 X lo-' 10-6-1.6 X IO-' 10-6-9 x 10-4 1o -0-1 0 -4
510p0
125
55.0 460.0 17.0 7.0 14.0 6.75 4.90 -600.
(120)
110 148(128)" 125 100 95
(?I
The results given in parentheses are based on the assumed value of 128 for high. a
A0
of pyrene',Na+.
1.15 0.15 4.85 6.75 (9.0)" 4.55 15.0 23.0 4 . 1
The value 148 seems to be unduly
PYRENE, SLOPE, 7.0 he= 148 M
/IO-'
NAPHTHALENE, SLOPE.460 A.=IIO
Cp'
/
1.0
'
0.5
" I
I
-
'
'
I04xCAf2/F
20
10
-
K ~ l s eO.2 s x IO-'
/
IO' F/A
25 . r
2.0
1.5
1.95 K oiss = 23.3 X 16'M
1.0 05
BIPHENYL, SLOPE.55.0 Ao=125 KOlss =l.15.10-6M
0
10
20
30
40
R4XcAf2/F
TRIPHENYLENE, SLOPE * 17.1
Figure 1.
A'.*125
Kmss =4.85x ld'M
The Significance of A. Values A brief discussion of 110 is not out of place. Inspection of Figures 1 and 2 indicates that A. is reliableat least within 10% for perylene,- anthracene-, and tetracene- arid is good within 15% for pyrene-. The observed gradation is plausible. A0 should increase with decreasing size of the ions, and indeed A" values are 95 and 100 cm.2/ohm equiv., respectively, for the largest anions, perylene- and tetracene-, and still larger for anthracene- (125) and for pyrcne' (148). Hence, the The Journal of I'hysieal Chemislry
I
10
1
20 I04x CAfz/F
1
30
1 40
Figure 2.
values of A0 = 110 for triphenylene- and A 0 = 125 for biphenyl- appear to be reasonable although they might be slightly too low. For naphthalene- we assume, rather arbitrarily, A. = 120. Nevertheless, it will be seen that the true A. value for naphthalene-,Naf cannot be much different from the chosen one.
EQUILIBRIA I N SYSTEMS O F AROMATIC HYDROCARBONT,N&+e RADICAL ANION-
There is an alternative method to check the reliability of the proposed A. values. The mobility of a salt is given by the sum of Aof and AO-, viz., the mobilities of the cations and anions. For the Na+ ion in T H F a t 25", Ao+ was determineds to be 48 cm.2/ohm equiv., and hence the data given in Table I V lead to the values for A(,- listed in Table V. Table V : Ao- of Aromatic Radical Ions in T H F a t 25'" Ao- calcd. Aromatic radical ion
B i p heny lT Naphthalene' Triphenylene' Pyrene' AnthraceneTetracene' Perylene'
cm.a/sec.
from D, cm.P/ohrn equiv.
2.3 2.3
85 85
lWD,
...
... 2.2
...
...
...
... 81
...
...
Ao- caled. from Ao,
cm.P/ohm equiv.
77 (70) 62 100 (80) 77 52 47
Diffusion constant D is calculated from the data of ref. 9 using their average values for the product Dq and taking 4.6 mp. for the viscosity of THF at 25".
4129
than the Li+ or Na+ cations, strongly indicating a lack of any specific coordination of the negative ions with the solvent (THF). The aromatic radical anions have apparently the shape of a relatively thin platelet which moves easily, whereas the Li+ or Na+ ions, which are probably coordinated with at least four molecules of THF, form large, slowly moving spheres.
Discussion The values of the electron affinities of aromatic hydrocarbons in THF, which were reported in the previous paper14need some corrections to account for the variable degree of association of the respective radical anions with Na+ counterions. These corrections can now be calculated, and the adjusted electron affinities, together with the original ones, are listed in Table VI. Table VI : Electron Affinities of Aromatic Hydrocarbons in T H F a t 25", Corrected for the Variable Degree of Association with N a + Ions Hydrocarbon
On the other hand, the phenomenological relation between A,- of an ion and its diffusion coefficient, 6.47 X lo6&- = D / k T , permits the calculation of one from the other. It is probable that the diffusion coefficient of the Ar- radical ion is only slightly smaller than that of the corresponding hydrocarbon, and therefore the latter may be used to calculate the relevant Ao-. The diffusion coefficients of some aromatic hydrocarbons have recentlyg been determined, and from the published data, which show the constancy of Dq, the diffusion constants in THF a t 25" were calculated. These are listed in the second column of Table V, the following column giving the Ao- computed from these D values. Comparison of the last two columns of Table V is most gratifying; it shows that our Ao- values are indeed only slightly smaller than those derived from the diffusion constants of the respective hydrocarbons. This agreement indicates The low conductthat Ao- for naphthalene- is -70. ance of its solution made it difficult to determine directly the A0 of this salt from the conductance data, and therefore we estimated its value from the diffusion constant of naphthalene. Having established the degree of reliability of the Ao- values of aromatic radical ions, we may compare them with the Ao+ of Li+, Na+, CS+, and BudN+, viz., 36.6, 48, 68, and 44.5, respectively.s I n spite of their large areas, the radical anions are more mobile
+ T\ia+
Biphenyl Naphthalene Triphenylene Phenanthrene Pyrene 9J0-Dimethylanthracene Anthracene Perylene Tetracene
Original e", v., ref. 4
(0.0) 0.066f.0.02 0.113zt0.01 0.124f0.005 0.505 f 0 . 0 0 5 0.607 f 0.005 0.624 f 0.005 0.917 f 0.005 1.025 ztO.01
Corrected to, V.
(0.0)
0.043 0.128
...
0.529 0.616 0.642 0.956 1.058
On the whole, the corrected E O values are only slightly larger than those reported in ref. 4, the largest increase of about 0.04 v. being found for perylene (0.917 to 0.956). Naphthalene is an exception, its electron affinity being by about 0.02 v. lower than that reported previously. The extremely strong association of naphthalener with Na+ increases, however, the spectrophotometric equilibrium constant biphenyl-,Na+
+ naphthalene JJ biphenyl + naphthalene-, N a
+
The large degree of association between naphthalene- and Na+ calls for some comments. Because of its exceptional behavior, the conductance studies of its (8) D. N.
Bhattacharyya, C . L. Lee,J. Smid, and M. Szwarc, J .
Phys. Chem., 69, 608 (1965). (9) T.
A. Miller, B. Prater, J. K. Lee, and R. N. Adams, J. Am.
Chem. Soc., 87, 121 (1965).
Volume 69, Number 1.2 December 1966
R. V. SLATESAND M. SZWARC
4130
solutions were repeated several times. There is, therefore, little doubt about its low dissociation constant. The dissociation of naphthalene- ,Na + was investigated also by Atherson and Weissmanld who used an e.s.r. technique. This interesting method distinguishes ion pairs, in which an additional splitting of the em-. signal occurs due to the presence of Na23, from the free ions which do not show further splitting. Experimentally, the method is not too reliable, as far as absolute values are concerned (see the text of ref. Id), and moreover it determines probably only the contact ion pairs, counting the solvent separated pairs as free ions. Hence, the dissociation constant, 1.5 X lo4 M , determined by this technique may be too high. The relative smallness and symmetry of thc anion probably leads to a central location of the Na+ cation in the ion pair, and the proximity of all of the negative charge to Ka+ accounts therefore for the large binding energy. The greater delocalization of the negative charge in the larger radical ions reduces the attraction and probably causes an unsymmetrical placement of the counterion. Perylene-,Na+ represents one of the extreme cases since the counterion can only be associated with one of the two naphthalene rings that form this hydrocarbon. Thus, Na+ is in the proximity of only 0.5 of the negative charge residing on this anion, and hence the high degree of dissociation of the perylene-,Xa+ pair is understandable. It should not be concluded that the counterion is rigidly located with respect to the anion. On the contrary, there is evidenceldII0that the counterion vibrates rapidly between the positions of the highest negative density if two or more of such positions exist in the anion.
Remarks Concerning the Paper by Buschlow, Dieleman, and Hoijtink2* I n a most interesting paper, Hoijtink, et aLlZadiscussed the dissociation of ion pairs-aromatic-, alkali ion+-in tetrahydrofuran (THF) and methyltetrahydrofuran (MTHF). The conductance of each of the investigated ion pairs was determined over a wide temperature range (25 to -75") but a t one concentraM ) only. The degree of dissociation was estition mated from the temperature dependence of the conductance. Although this method may give some qualitative information about the extent of the dissociation, often it could be misleading. For example, Hoijtink considers a salt to be completely dissociated if no maximum is shown in the conductance-temperature curve. This is not a valid conclusion, and indeed it led Hoijtink to believe that in THI? at 25" anthraccne-,Na+ The Journal of Physical Chemistry
is completely dissociated a t the concentration of
lo-*
M , while its dissociation under these conditions amounts to 8% only. At a constant salt concentration, the temperature dependence of the conductance curve is governed by several factors: (1) the temperature dependence of the solvent viscosity, (2) the heat of dissociation of the investigated ion pair, (3) the temperature dependence of Aoq, and (4)the degree of dissociation of the salt. The viscosities of T H F and M T H F have been shown" to be given by a linear, Arrhenius-type relation which is valid over the whole temperature range of 25 to -75". Thus, -3.655
+ 393/T
log ~ M T H F= -3.625
+ 384/T
log qTHF
=
and
and the respective "activation energies" are -1.76 and 1.72 kcal./mole. The heat of dissociation of the Na+, CS+, and NBus(i-Am)+ salts of BPh4- have also been investigated in this temperature range," and in T H F the heat of dissociation of Na+BPh4- was found to increase from the -1.3 kcal./mole a t 25" to about 0 kcal./mole a t -70". The concentration of the ions is proportional to Kdiaa"*if the degree of dissociation is low, and therefore the temperature coefficient of con1/2AHdisa. The conductductance is given by "Eq" ance of Na+BPh4- decreases, therefore, steadily on lowering the temperature, clearly indicating the unreliability of Hoijtink's criterion. There are good reasons to believe that in T H F solutions Li+X- ion pairs may form solvent-separated pairs, and, since the exothermicity of their dissociation is quite low,12their conductance may not show a maximum. Hence, the criterion discussed above might again be misleading. The contribution of triple ions to the conductance must not be neglected, and, since their concentration decreases at lower temperatures, this effect may lead to the disappearance of the maximum in the conductance-tempcrature curve. Finally, if the investigated ion pairs are completely dissociated, the observed A should be close to Ao. It was showns that Ao+ of Li+ and Na+ in THF a t 25" are 36.6 and 48, respectively, whereas the A values of the Li+ and Na+ salts of anthracene- reported in ref. 2a are -18 and 28 only. This proves again
+
(10) E. d e Boer and E:L. Mackor, J . Am. Chem. Soe., 86,1513(1964). (11) C. Carvajd, J. K. Tiille, J. Smid, and M. Szwarc, ibid., in press. (12) T. E. Hogen-Esch and J. Smid, ibid., 87, 669 (1965).
4131
REACTIONS OF ALANINE WITH REDUCING SPECIESFORMED IN WATERRADIOLYSIS
that under those conditions the dissociation is less than 50Yc.
Conductance of Anthracene2-,2Na+ Several samples of anthracene2-,2Na+ were prepared, and their conductances were investigated. Unfortunately, the scatter is too great to allow a reliable determination of A, or l/Ao2Kdi,,; however, the very low conductance of these solutions proves the low degree of dissociation which we estimate a t about lo-’ M . Our findings agree again with the qualitative observations of Hoijtink, et aL2* Apparently, A2-,2Na+ are contact agglomerates whereas A’,Na+ may be, to a large extent, a solvent-separated ion pair. See in this
connection ref. 13. Discussion of the disproportionaA which resembles ours was tion of 2A- into A2published by Shaten~tein’~ in 1964. See also ref. 13.
+
Acknowledgment. The financial support of this study by the U. S. Army Research Office (Durham), Grant DA-ARO (D)-31-124-G52 1, Texaco, Inc., and the National Science Foundation is gratefully acknowledged.
(13) R. C. Roberts and M. Sawarc, J. Am. Chem. SOC.,in press. (14) A. I. Shatenstein, E. S. Petrov, and M. I. Belusova, “Organic Reactivity,” Vol. 1, Tartu State University, Estonia, U.S.S.R., 1964, p. 191.
Reactions of Alanine with the Reducing Species Formed in Water Radiolysis’
by Boyd M. Weeks, Sibyl A. Cole, and Warren M, Garrison Lawrence Radiation Laboratory, University of California, Berkeley, California (Received June 10,1966)
A detailed study has been made of the effects of pH and of added scavengers on the reactions of alanine with the reducing species formed in water radiolysis. A principal reaction of the hydrated electron (eaq-) leads to degradation of the N-C linkage, e.g., eaqRNH3+ + R “3. The rate of removal of g,- by such reaction is strongly dependent on the ionic form of the a-amino acid; /?-alanine is relatively less reactive toward eaq- under all conditions of pH. A suggested reaction scheme accounts both qualitatively and quantitatively for the observed effects of pH on product yields from oxygen-free solutions of alanine under y rays.
+
+
The principal actions of ionizing radiations on solutes in dilute aqueous solution are initiated by the radiation-induced step2
HzO -m+
HzOz,Hz, OH, H, eaq- H+
(1)
Subsequent reactions of these oxidizing and reducing species with the simpler amino acids such’ as glycine and alanine in oxygen-free solutions lead to both oxidative and reductive deamination with formation of the corresponding keto acid and fatty acid as major degradation products. The relative and absolute yields of both major and minor products from glycine and alanine are strongly dependent on the pH of the
irradiated solution. One pH-dependent reaction that must be considered is of course the conversion of eaqto H by the hydronium ion2 eaq-
+ H30++H + H 2 0
(2)
However, conversion of eaq- to H is not specific to the (1) This work waa done under the auspices of the U. S. Atomic Energy Commission. (2) E. J. Hart, Science, 146, 19 (1964). (3) (a) G. Stein and J. Weiss, J. Chem. SOC.,3256 (1949); (b) N. E. Sharpless, A. E. Blair, and C. R. Maxwell, Radiation Res., 2, 135 (1955); (c) B. M. Weeks and W. M. Garrison, ibid., 9, 291 (1958).
Volume 69,Number 18 December 1966
