2535
2,6-Dichloro-4-nitrophenyl-Amine System
Our observation time is limited to the duration of the fluorescence, which means, practically, that we cannot observe the variation of the rotational temperature beyond 400 to 500 nsec. With 22 Torr of helium the Boltzmann equilibrium is not reached before 300 nsec. Figure 4 shows that in all cases, at least in the time region where the Boltzmann equilibrium is observed, the highest CH (A2A)rotational temperature observed is 2780 f 140’K around 100 nsec. The energy of dissociation according to one of the mechanisms 1-4, appearing as rotational energy of the fragment CH (A2A),is thus equal to 5.5 f 0.1 kcal mol-’, in good agreement with the value measured by Beenakker (2700 f 150’K; i.e., 5.4 kcal mol-l). Acknowledgments. We wish to thank Drs. De Heer and
Beenakker for communication of their results before publication, and helpful corrections of our first manuscript. References and Notes (1) (2) (3) (4) (5) (6) (7) (8) (9)
M. Clerc and M. Schmidt, Faraday Discuss. Chem. SOC.,53, 217 (1972). T. Carrington, J. Chem. fhys., 35, 807 (1961). C. I.M. Beenakker, Thesis, Heiden, 1974, F.O.M. Institute, Amsterdam. C. I. M. Beenakker, P. J. F. Verbeek, G. R. Mohlmann. and F. J. De Heer. J. Quant. Spectrosc. Radiat. Transfer, submitted for publication. M. Clerc, These de Doctorat, Paris, Universite d’Orsay, 1972. R. S. Mulliken, Rev. Mod. fhys., 3, 89 (1931). C. I. M. Beenakker and F. J. De Heer, Chem. fhys., submitted for publication. R. C. Baas and C. i. M. Beenakker, Cornput. fhys. Common., submitted for publication. I. Kovacs, “Rotational Structure in the Spectra of Diatomic Molecules”, Adam Hilger, Ltd., London, 1969.
Ultraviolet Spectra and Proton-Transfer Equilibria in 2,6- Dichioro-4-nitraphenol-Amine Systems H. Romanowski and L. Sobczyk’ Institute of Chemistry, University of Wroclaw, 50-383 Wroclaw, Poland (Received December 13, 1974)
Hydrogen-bonded complexes formed between 2,6-dichloro-4-nitroghenol and amines, of 1:l and 1:2 composition, were studied using uv spectroscopy. Parallel to the proton-transfer equilibrium within the hydrogen bonding, a proton detachment from the phenol molecule due to formation of homoconjugated cations is distinctly observed. This process is particularly favored by increasing the electric permittivity of the solvent. A linear correlation between the logarithm of the proton-transfer equilibrium constant and the ionization potential of the amine has been found for systems without steric hindrances. Triethylamine and sym- collidine do not form homoconjugated cations.
Introduction The problem of proton-transfer equilibria in the phenolamine systems has been already dealt with in a number of papers. Particular attention was paid to the electronic and to the dipole moment^^-^^ of the complexes formed. Though the conditions employed in both methods are somewhat different, the results obtained indicated definitely that, for a certain ApK, value of the reacting components and at a specific solvent activity, there is a protontransfer equilibrium in the complex A H . . . B =+ A- ...H-B+ The proton abstraction from phenol is most evident in the uv spectra; there is a strong bathochromic shift of the lowfrequency c c* band and its intensity increases. Apart of the formation of the simple AH-B adduct, some additional complexation equilibria may appear in a reacting mixture of phenols and amines. A t excess phenol concentrations (AH),-B complexes may be formed and this always favors the proton transfer
-
(AH),
.-B+=
(A,H,
, )... HB+
The symptoms of such phenomena are quite evident in the dielectric proper tie^.'^ In the electronic spectra these equilibria remain usually unnoticed since the measurements are usually carried out at excess amine concentrations. How-
ever, then some other important equilibria are manifested and these may be expressed by means of AH
+ B,
-T
AH-B,,
i A--[HBJ+
Also in this case a stimulated proton transfer can take place, the more so as for primary and secondary amines additional stabilization with the A- anion by “coordination” of a great number of amine molecules occurs. The results obtained so far by means of the ir, conductometric, and potentiometric methods suggest that the most stable species is a complex [BHB]+ ion with a dynamically symmetric hydrogen bonding. It is, however, still unclear which factors affect the stability of this bonding and, hence, the stability of the complex ion. An interesting and important problem is also the effect of the molecular environment on the equilibria in question. It is certain that this is one of the dominant factors on which the proton transfer depends.6 The solvent activity is itself a decisive effect on the formation of the [BHB]+ ion since this process requires, in fact, the ionization to take place. A number of spectroscopic studies have been made using 2,6-dichloro-4-nitrophenol (DCNP), pK, = 3.68, as an acid agent.l* However, the equilibria have not been fully defined. In the present paper we decided to investigate the complexation of DCNP with various aliphatic and aromatic The Joumalgf fhysicaf Chemistry, Vol. 79, No. 23, 1975
2536
H. Romanowskiand L. Sobczyk
Figure 1. Absorption spectrum of DCNP in water at pH 1 (WI) and pH 12 (W12), in pure pyridine (A,,), and in CC14 (cpknol5.871 X the presence of pyridine at various concentrations: 0.0 (Ph); 0.124 M(1); 0.99 M ( 2 ) ; 2.48 M(3); 6.2 M(4).
MJ in
I
C
m : 3
3
h
idloo
22bm
.9[C2+
Figure 2. Absorption spectrum of DCNP in water at pH 12 (WI~),in pure pyridine (A,,,), and in a mixture of 1:2 CC14-C2H4C12 MJ in the presence of pyridine at various concentrations: 0.0 (Ph); 0.062 M(1); 1.24 M ( 2 ) ; 2.48 M(3); 6.2 M ( 4 ) . The Journal of Physical Chemistry, Vol. 79. No. 23, 1975
(Cphenol
5.36
x
2,6-Dichioro-4-nitrophenyl-Amine System
2537
4
' ',
1
I .
%?IO0
0--.--0
3 m
26.000
22.000
"Ti-
,ij[cm
MJ in the presence of 4Figure 3. Absorption spectrum of DCNP in water at pH 1 (WI) and pH 12 (Wi2) and in C2H4CI2(Cphenol 5.19 X tert-butylpyridine at various Concentrations: 0.0 (Ph); 0.001 M(1); 0.01 M ( 2 ) ; 0.1 M(3); 0.68 M(4); 2.04 M ( 5 ) ; 3.4 M(6).
Figure 4. Absorption spectrum of DCNP in water at pH 1 (Wi) and pH 12 (Wlz), in pure n-butylamine (A,,,), and in C2H4CI2(Cphenol 5.03 X MJ in the presence of n-butylamine at various concentrations: 0.0 (Ph), 3.04 X M (1); 1.01 X M (2); 3.04 X M (3); 1.01 X M(4); 0.0101 M ( 5 ) ;0.101 M ( 6 ) ;5.07 M(7). The Journal of Physical Chemistry, Vol. 79,No. 23, 1975
2538
amines in order to determine the effect of the amine structure and their basicity on the equilibria involved. C c 4 ( E D 2.2), 172-dichloroethane ( E D 10.4), and mixtures of these compounds at a 1:2 ratio ( E D 8.2) were used as solvents. Application of more polar solvents is, as a rule, equivalent to a distinct increase of electron donor-acceptor properties of the medium and should lead to additional complications arising from solvent-solute specific interactions.
H. Romanowski and L. Sobczyk
Scheme I Ph-OH f NR,
Experimental Section 2,6-Dichloro-4-nitrophenol (Fluka) was purified by crystallizing it from ethanol and dried in a vacuum desiccator over anhydrous MgS04. The amines were distilled before each series of measurements over KOH. Solvents cc14 and C2H4C12 (FOCh Gliwice, Analar grade) were dried at first over P205 and then distilled, collecting the fractions boiling at 77 and 86OC, respectively, and dried and stored over freshly prepared molecular sieves. The measurements were carried out on a Zeiss-Specord spectrophotometer in the 38000-20000-~m-~range, in silica cells 1 cm thick. Phenol solutions (5 x lW5 M ) and amine solutions of concentrations 5-1 X 10-5 M were prepared immediately before measurement. Only in such conditions are repeatable results obtained. Some irreversible changes which occur in the solution after storing it for a longer time are presumably caused by photochemical reactions.
Kl
A Ph-OH ‘..NRJ
secondary amines. Equilibrium scheme I is preserved, but for amines with active proton-donor groups, the formation of complexes with more amine molecules occurs. This applies in particular to the B and C states. Now, let us examine the behavior of DCNP toward particular amines in the three selected environments, i.e., C c 4 , C2H4C12, and a 1:2 mixture of these solvents. The spectral characteristics of the DCNP-py systems are illustrated in Figures 1 and 2. In the case of nonpolar carbon tetrachloride, the A * B equilibrium is merely very weakly marked and the C state does not appear even at high excess amine concentrations. The sensitivity of the equilibrium of an intramolecular proton transfer on the polarity of the environment is manifested by an evident effect of the polar pyridine concentration on the concentration of the B state. Results and Discussion This is probably not a question of the specific solvation of interaction products since this is very similar for both the The band of free 2,6-dichloro-4-nitrophenol in an inacA and B states. A determining factor seems to be the ditive solvent is located near 34000 cm-l. In complexes with electric permittivity of the medium. The solvent effect is “ordinary” hydrogen bonds this band, according to the still more clearly manifested in the B C equilibrium. The data obtained for other phenols, is bathochromically shiftC state corresponds, in fact, to the “pure” ions and does ed to the 31000-32000-~m-~region and corresponds to the DCNP spectrum in water a t pH 1. The DCNP anion in not appear in pure CCL. It is enough to use a mixture of CC14 CzH4C12 to observe this state very clearly as a band water at pH 12 absorbs at 25000 cm-l; it may be then sugat about 23500 cm-l (Figure 2). gested that the absorption in this region may be ascribed to By replacing pyridine with 4-tert-butylpyridine, Bupy, phenolate anions. If, however, we are dealing with the ionic the equilibrium of the intramolecular proton transfer is pairs A--HB+, in which a strong coupling through the hyfound to be evidently shifted toward the B state. The redrogen bond occurs, then this band becomes shifted hyposults for Bupy shown in Figure 3 characterize lucidly and chromically with respect to a more free anion in aqueous univocally all of the equilibria composed in eq i. There are solution and lies between 25000 and 27000 cm-l, dependtwo isosbestic points I and I1 which correspond to the foring on the type of base and solvent. With increasing solvent mation of 1:l and 1:2 complexes. Since the band intensity polarity this band is, as a rule, shifted bathochromically.’ ratio a t 31200 and 25800 cm-l remains constant, this is disApart, however, from the band typical of the phenolate tinct evidence of an intramolecular A + B equilibrium. At anion in aqueous solution, a new low-frequency band is high excess amine concentrations, a low-frequency band found at high excess amine concentrations. Considering the evidence supporting the formation of [BHB]+ ions1p4>6p15-20 appears at 23700 cm-I which corresponds to the undisturbed PhO- anion formed by binding the [HBupy]” catone should expect that in this case we are dealing with fairion to yield a complex [BuppH-.pyBu]+ cation. ly free A- anions. Thus, one may suggest that the weaker In case of stronger bases, the interaction with DCNP the interaction of this anion with its environment, the leads to the formation of a 1:l complex with a distinct prelower the frequency of this band. This fact becomes comdominance of the tautomeric B form, and the contribution prehensible when the band observed is assumed to correof the A form may be neglected in quantitative considerspond to the intramolecular electron transfer. Stabilization ations. As shown in the spectra of the n-butylamine (BA)of the state with the charge localized on the oxygen atom DCNP system in C2H4C12 (Figure 4),the isosbestic point I (through specific interaction) should shift the band in is well preserved at a not too high excess of BA concentraquestion toward higher frequencies. All the experimental tion, i.e., under conditions when no 2:l complexes are evidence collected until now seems to be in full agreement formed. Isosbestic point I1 is not preserved since before the with such an approach. phenol becomes fully complexed to yield a 1:l complex, the Putting aside additional solvation processes, the reacB C equilibrium appears. tions of DCNP with the tertiary amines may be presented It is interesting to note the behavior of the DCNP specby means of Scheme I. The stability of the C state should trum in the presence of the aliphatic primary and secondepend strongly on the solvent polarity and amine strucdary amines. The maximum band absorption of the “free” ture which is, in fact, confirmed by our results. A more PhO- ion lies at much higher frequencies than, e.g., in the complicated situation occurs in the case of the primary and
+
The Journal of Physical Chemistry, Val, 79, No. 23, 1975
2539
2,6-Dichloro-4-nitrophenyl-Amine System
TABLE I: Positions of the Absorption Bands (cm-I) for Various Prototropic Forms of 2,6-Dichloro-4-nitrophenol-amine Complexes
Amine Pyridine
Solvent CCl, 1:2 CCl4-CzH,C1, C2H4C12
4
- tevt- Butylpyridine
Pure amine CCl, 1:2 CCl,-C,H,Cl, CZH4C12
Collidine
CC1, 1:2 CCl4-C2H4C12 C2H4C12
n- Butylamine
CC1, 1:2 CC14-C,H*C12 C2H4C12
Diethylamine
Triethylamine
Pure amine CCl, 1:2 CC1,-C,H4C1,
PhO- (“free”) in PhOPhOH*..NG
34300 34000 33800
31800 31300 31300
34300 34000 33800 34300 34000 33800 34300 34000 33800
32000 31500 31200 31800 31700 31700 31700 31600
PhO-..*HN’T/ 25800 25700 25700
26000 25800 26700-25400 25600-25200 2 54 00-2 5000 27300 26300 25800
C2H4C1’2
34300 34000 33800
27100 25800 25500
C2H4C12
34300 34000 33800
26750 25450 25040 26700 25000 (pH 12)
Pure amine cc1, 1:2 CCl,-C,H,Cl,
,
[SN.+
PhOH (free)
Pure amine 32100
Water
complexes 23500 23300 23250 2 52 00 24600 23700
25250 24600 24300 2 5150 256007 251501 24800? 25550?
(PH 1) TABLE 11: Equilibrium Constants of Prototropic Reactions Amine and IP” Pyridine I , = 9.23 eV 4
- tevt- Butylpyr idine
I,
8.95 eV
Collidine I, 8.8 e V
n- Butylam ine I , = 8.71 eV Diethylamine I , = 8.01 eV Triethylamine I, = 7.5 eV
Solvent
log K*
CC1, 1:2 CCl,-C2H,C1,
0.17 0.48 1.84 1.78 2.48 2.48 2.23 2.70 3.30 3.40 4.60 4 -48 4.73 5.25 5.58 4.23 5.41 5.50
C2H4C12
CCl, 1:2 CCl,-C,H,ClZ C2H4C12
CCl, 1:2 CC14-C2H4C1, C2H4C12
CCl, 1:2 CCl,-C,H,Cl, C2H4C12
CCl, 1:2 CCl,-C,H,Cl, C2H4C12
cc1, 1:2 CC1,-C2H,C1, C2H4C12
a
log K,*
12
n
-0.79 -0.34 -0.19 0.74
0.91 0.94
1.88 2.07
0.20 0.56 2.40 (?) 1.67 2 -43 2 -69
Ionization potentials from ref 28.
presence of the aromatic amines (cf. Table I). This is presumably the result of interactions between the anion and the amines through the hydrogen bond which stabilizes firmly the ground state / Ph-O-... H-N, or
:. H \N Ph-0-.,, , H
In the presence of triethylamine no “free” anions are formed (Figure 5) which is due to the fact that this compound is not capable of forming homoconjugated cations r\
/1+
LLN...H...xF] /
\
Our results are in agreement with this conclusion. The Journal of Physical Chemistry, Vol. 79, No. 23, 1975
2540
H. Romanowski and L. Sobczyk
31000
3WOO
26000
Figure 5. Absorption spectrum of DCNP in water at pH 12 ( W I ~ ) ,in pure triethylamine (A,,), and in CCI4 (cphenol 5.96 X M ) in the presence of triethylamine at various concentrations: 0.0 (Ph); 3.6 X M ( 1 ) ; 1.08 X lov4 M ( 2 ) ; 1.08 X lov3 M ( 3 ) ; 0.0108 M ( 4 ) ; 3.6 M ( 5 ) .
34*000
30600-
2doa
22600
3I
r
n
O
-
d
2-,
Figure 6. Absorption spectrum of DCNP in water at pH 12 ( W I ~ ) ,in pure diethylamine (Am), and in a 1:2 mixture of solvents CCI4-C2H4Cl2 in the presence of diethylamine at various concentrations: 0.0 (Ph); 4.8 X M (1); 9.6 X M (2); 9.6 x (cphenol 4.93 X M (3); 0.096 M (4); 0.26 M (5); 4.825 M (6). The Journal of Physical Chemistry, Vo/. 79, No. 23, 1975
2,6-Dichloro-4-nitrophenyl-Amine System
2541
/ /-,
rDT I
'
1 P i
r5 1 i i
T
o,J.
2
Flgure 7. Absorption spectrum of DCNP in water at pH 12 (WI2 and in C2H4Cl2(cphenol 5.53 X M(1);7.56 X IO- M ( 2 ) ; 0.378 M ( 3 ) . various concentrations: 0.0 (Ph); 7.56 X
Dependence of the molar absorption coefficient e for DCNP complexes with collidine on excess of amine concentrations (B 25000 cm-l) Flgure 8.
The behavior of DCNP in the presence of diethylamine (DEA) is not quite clear (Figure 6). The studies of Coetzee, Padmanabhan, and Cunninghamla indicate that DEA may only to a slight extent form the complex cations. The spectral changes occurring at high DEA concentrations could not be explained as due only to a drastic change in the polarity of the environment and to quite different solvation conditions. Also in this case a B ==C equilibrium probably exists but to a much lesser extent. In the case of collidine, the B == C equilibrium is again invisible, though at high excess concentrations of this base a bathochromic shift of the band takes place. This applies particularly to CC4. The weakest effect is found for ethylene chloride (Figure 7). This effect may be ascribed to the polarity changes due to the high dipole moment of colli-
M)
in the presence of symcollidine at
dine. The fact that collidine probably does not form any complex [BHB]+ cations is certainly due to steric reas o n ~ . ~ ~ - ~ ~ The effect of solvent polarity on the intramolecular proton transfer in a DCNP complex with collidine is illustrated in Figure 8. An increase in polarity is manifested both by a drastic change in the equilibrium constant K* and by only a slight increase in the molar extinction coefficient. An attempt was made to determine quantitatively the proton equilibria occurring in the systems under investigation. The results given in Table I1 are the average values of equilibrium constants calculated from the spectra at various wavelengths by the Rose-Drago method.25The values of K* constants (eq 1) given in Table I1 were calculated from the peak band intensity of the form B, and the values of K*2 constants (eq 2), from the C band intensity (for such reactions the Rose-Drago equation assumes a form presented in ref 23). In eq 1 and 2 Coph and C O B are the initial concentrations of the DCNP and amine, respectively, E is the observed absorbance at given wavelength, and tC and Cph are the molar absorption coefficients of the formed complex and of the DCNP, respectively.
E - €phCoph) E cc
- cph
- CphCOph ec
- eph
(2)
The quantitative analysis of the DCNP-Bupy system is The Journal of Physlcal Chemistry, Vol. 79, No. 23, 1975
H. Romanowski and L. Sobczyk
2542
Figure 9. Dependence of log K’ and log K‘ on the ionization potential ID of amine: (i) ionization potential of collidine was assumed as 8.8 eV (for various lutidines 8.85 eVZ8);(ii) ionization otential of Bupy was assumed to be 8.95 eV (for y-picoline 9.01 eV?S ).
the most accurate for reasons given in the description of spectra. From the assumed model it follows that in the first equilibrium one amine molecule participates per phenol molecule, while in the second equilibrium about two amine molecules participate on the average. The estimated error in the log K values given in Table I1 can reach even f0.3. The error is serious, as the equilibria dealt with are fairly complicated. As expected, both of these equilibrium constants increase with the increasing solvent polarity, I t is more difficult to determine the effect of amine basicity on the K* equilibrium constant since, apart from the proton-transfer process alone, the autosolvation effects become evident, particularly for the primary and secondary amines. Nevertheless, as shown in Figure 9, there is an approximately linear correlation between log K* (determining the free energy of the proton transfer) and the amine ionization potential,26which is the most reasonable quantity to express the amine basicity in nonaqueous solution. A distinct abnormal deviation from linearity occurs for the DCNP complexes with TEA. Disregarding the errors resulting from the calculations of the equilibrium constants, the reason for such a deviation is most probably steric hindrances which are drastically manifested by the absence of the B C equilibrium. The effect of the steric structure of the amine on the complex-forming ability with phenols has also been observed by Lin and It seems, however, that the steric effect becomes significant already in the secondary amines. There is also a linear correlation between log K2* and IDbut with a lower inclination.
Conclusions In the present work the acid-base equilibria in the 2,6dichloro-4-nitrophenol-aminessystems were investigated. The following points were demonstrated. (1) In the presence of excess amine and a t low phenol concentration, the complexation process occurs in three
The Journal of Physical Chemistry, Vol. 79,No. 23, 1975
stages, the determining steering factor being the formation of the [BHB]+ cations and “release” of the phenolate ion PhO-. (2) The type of the equilibrium is strongly affected by the solvent-an increase in the solvent polarity facilitates the proton transfer and the formation of cations and anions. (3) The appearance of the B C equilibrium depends to a large extent on the steric structure of the amine. ( 4 ) The behavior of the aliphatic and aromatic (cyclic) amines is different-aliphatic amines interact with the “free” phenolate ions to reverse the equilibrium. (5) The presence of particular stages and the stability of the complexes formed depend, above all, on the acidity of phenols and basicity of amines. (6) There is an approximately linear relationship between the logarithm of equilibrium constant and the amine ionization potential, the equilibrium constant K* of the intramolecular proton transfer being increased more rapidly than the K2* constant with increasing amine basicity (decreasing ID). Acknowledgment. This work was financially supported by the Polish Academy of Sciences under the program PAN-111.
References and Notes S. Nagakura and H. Baba, J. Am. Chem. SOC.,74, 5693 (1952). S. Nagakura, J. Am. Chem. SOC.,76, 3070 (1954). H. Baba and S. Suzuki, J. Chem. Phys., 35, 1118 (1961). H. Baba, A. Matsuyama, and H. Kokubun, J. Chem. Phys., 41, 895 (1964). (5) R. Scott, D. De Paima, and S. Vinogradov, J. Phys. Chem., 72, 3192 ( 1968). (6) H. Baba, A. Matsuyama, and H. Kokubun, Spectrochlm. Acta, Part A, 25, 1709 (1969). (7) R. Scott and S. Vinogradov, J. Phys. Chem., 73, 1890 (1969). (8) R. A.-Hudson, R. M. Scott, and S. Vinogradov, J. Phys. Chem., 76, 1969 (1972). (9) J. W. Smith, J. Chlm. Phys. Phys.-Chim. Blol., 58, 182 (1964). (IO) H.Ratajczak and L. Sobczyk, J. Chem. Phys., 50, 556 (1969). (11) J. P. Hawranek, J. Oszust, and L. Sobczyk, J. Phys. Chem., 76, 2112 (1972). (12) R. Nouwen and P.Huyskens,J. Mol. Struct., 16, 459 (1973). (13) L. Sobczyk and 2. Pawelka, ROC.?.Chem., 47, 1523 (1973). (14) T. Jasihski and T. Widernikowa, Rocz. Chem., 43, 1031, 1253, 1877 (1969). (15) D. Had%,J. Chem. SOC.,5128 (1962). (16) D. Cook, Can. J. Chem., 41, 2575(1963). (17) J. F. Coetzee, G. R. Padmanabhan, and G. P. Cunningham, Talanta, 11, 93 (1964). (18) J. F. Coetzee and G. R. Padmanabhan, J. Am. Chem. SOC.,87, 5005 (1965). (19) Z. Dega-Szafran, Bull. Acad. Pol. Sci., Ser. Sci. Chim., 15, 393 (1967). (20) R. Clements, R. L. Dean, and J. L. Wood, J. Chem. SOC.U,1127 (1971). (21) Ch. Debeuf and P. Huyskens, Ann. SOC.Sci. Bruxelles, Ser. 1, 83, 171 (1969). (22) L. Cattalini,M. Nicolini, and A. Orio, lnorg. Chem., 5, 1674 (1966). (23) T. Jasihski, T. Misiak, and T. Skariyhska-Klentak, Rocz. Chem., 42, 875 (1968). (24) T. Widernikowa and T. Jasihski, Zesz. Nauk. Wyzsz. Szk. Pedagog. Gdansku: Mat. Fir. Chem., 10, 115 (1970). (25) N. J. Rose and R. S. Drago, J. Am. Chem. SOC.,81,6138 (1959). (26) S. Nagakura and M.Gouterman, J. Chem. phys., 26, 88 1 (1957). (27) M.-L. Lin and R. M. Scott, J. Phys. Chem., 76, 587 (1972). ( 2 8 ) R. P. Blaunstein and L. G. Christophorou, Radiat. Res. Rev., 3, 89 (1971). (1) (2) (3) (4)
