Free Radicals Generated by Radiolysis of Aqueous Solutions Harold A. Schwarz Brookhaven National Laboratory, Upton, NY 11973 The free radicals produced in the radiolvsis of aqueous solutions span the range of redox potentials from -2.9 to +2.65 volts. If one also considers the fact that these radicals can be produced in several micromolar concentrations in time periods of a few nanoseconds, then it is not strange that several hundred papers appear each year on the suhject. This article will discuss the identity and nature of these radicals. Only a few rate studies will he mentioned. Excellent rate constant comp:lations are maintained hy rhv Radiation Chemistry Data Center at the i'n~versityof Notre Dame Radiation Laboratiwv ( I ) . Most of the discuss~onwill Iwahout results obtained with low LET radiation sources ("'(:o aamma radiation or electron accelerators). Identities of the Reducing Radicals The types of products found upon irradiating certain aqueous solutions offer clear evidence that two different reducing radicals are produced. For instance, the radiolysis of neutral chloroacetate solutions produces chloride ion as a product, while the radiolysis of acid chloroacetate solutions produces hydrogen (2). The explanation is that hydrated electrons are formed which react with organic chlorides to release chloride ions but which can react also with hydrogen ions to produce hydrogen atoms.
-
cSq + H30t H + H 2 0 (1) T h e hydrogen atoms abstract hydrogen from chloroacetate ions to give Hz as a product in acid solution. T h e hydrated electron yield can be equated to the yield of chloride ion from a dilute neutral solution of either chloroacetate ion or methyl chloride (3),2.65 ions formed per 100 eV absorbed. Some hydrogen atoms are also produced in neutral solution (4). Their yield can be estimated from measurements of Hz produced from solutions of a good hydrogen donor (ethanol) and hydrated electron scavenger. The ohserved hydrogen yield, 1.1, also contains 0.45 molecules/100 eV of "molecular" hydrogen, that is, H2 formed directly from water, so the H atom yield, a t low concentration, is 0.65 atomsl100 eV. The most convincing identification of the acidic reducing species as H atoms is the observation of their ESR spectrum in 0.1 M HC104 solutions irradiated with an electron beam while flowing through an ESR cavity (5).The two line signal with 506.2 gauss separation is ahsolutely characteristic of hydrogen atoms. T h e signal is much reduced, though still present, in neutral solution. An unusual aspect of the signal is that the low field line appears in emission while the high
J
0 5 I
I
I
600
I
800
1000
WAVELENGTH, nrn Figwe 1. Absorption spectra of intermediate$in water rsdiolysis. Hand OH, ref. 77.
ref. 35:
field line amears in absorption. an effect l a r ~ e l vdue to selective reac'tfon of spin statks. his effect has deei developed into a ~owerfultool for measurina- relative rate constants for hydrogen atom reactions. In neutral or sliahtlv basic solution a new ESR spectrum appears (6). It is asingie line, less than 0.15 gauss wide, with agfactor of 2.00033, and it has been identified as the spectrum of the hydrated electron. This identification is certainly consistent with what is known about the hydrated electron hut is somewhat less satisfactory than the identification of the hydrogen atom, and alternate methods of establishing its identity are helpful. The hydrated electron has a very intense and broad optical absorption spectrum (7) with a maximum a t 710 nm, hut i t extends from the UV to the near-infrared, as shown in Figure 1. In slightly basic oxygen-free water a t very low radiation doses (radical concentration helow 5 X 1 0 P M ) , the absorption will last for 10-3 sec. The reaction which ultimately limits the lifetime is e-.,
+ H20
F,
H
+ OH-
(2)
for which the rate constant (pseudo-first-order) is 900 sec-' (8). Reaction rate constants for the reaction of hydrated electrnnr; with various solutes can he measured by fullowing the exponential decay of the visible absorption with time as a function of the solute concentration. The looition and width
Volume 58 Number 2
February 1981
101
of the spectrum insure thar o m r wavelength can be found ar u,hich the reaction con he ltdlowed free or interference frum other ahsorbing species. 111this way reactions of the hydrated electron with charged 2nd neutral solutes have been studied as functions of ionic strength. From such measurements the charge on the common reactant (the hydrated electron) was found to be -1.0 (9). The reverse rate constant for reaction (2) a t about 0.01 M NaOH is 2 X 10" M-' sec-' (10). The eauilibrium constant of reaction (2) can he calculated as the ratib of the forward and reverse rate constants, or 4 X 10-5M (with HzO(1iq) as standard state). This equilibrium constant can be combined with the ion product of water to give the pK of the hydrogen atom as 9.6 (reaction (1) written right to left). The measured unit negative charge and the observation of reaction (1)combine to form a convincing proof of identity for the hydrated electron. The hydrogen atom also has an absorption spectrum associated with it, as seen in the lower left corner of Figure 1.The hvdroeen atom itself should not have a sDectrum in this region. The observed spectrum can be explained as either a perturbation of the first H 2 0 hand, normally beginning at 180 nm, by a nearest neighbor H atom (II), or as a charge transfer band, solvent to solute (12). In any event, this spectrum is of limited use for kinetic studies. A reasonable number of absolute rate constants have been measured bv pulse radiolysis experiments following the growth of prod& which have more convenient spectra. Many relative rate constants have been measured by the ESR technique mentioned above and by studying the efficiency of solutes in decreasing the hydrogen yield from the reaction H + RCH20H
-
H,
+ RCHOH
where RCHzOH is usually ethanol or 2-propanol. 60Co radiation sources are most useful for the latter studies.
lead to the following conclusions: (1)the stability (1.6 eV) is largely due to the presence of the continuum, (2) the water dipoles are oriented toward the center of the cluster, but the structure is not rigid, and (3) the minimum energy configuration is with the four oxygen atoms on a 2.7 A radius sphere (which agrees qualitatively with the observation of a 10 cc/ mole cavity). Approximately 80% of the excess electron density is within 4.2 A radius, which includes the four water molecules. The ESR spectrum of electrons in strongly alkaline ice a t liquid N2 temperatures is most consistent with 6 protons (3 water molecules) surrounding the electron (18). The single line spectrum in liquid water gives no information on this point. The Hydroxyl Radical No ESR spectrum attributable to the OH radical has been seen or is expected to be seen in irradiated water (though it is seen in irradiated ice (19)). Spin relaxation is extremely rapid in small radicals such as OH in which the electron is not in a oure s state. and so ESR lines would be verv broad. The reactivity of the'radical limits the concentration obtainable, and so lines would have to he less than 1rauss wide to be seen (20). The evidence for hydroxyl radical formation is somewhat indirect, then. For instance, ESR spectra of OH radical adducts have been observed, such as HOCHzNOz- formed by reaction with the enol form of nitromethane (20):
-
-
CHaNO1+ OH- CHz=N02- + Hz0 HOCH2NOsOH + CH2=N02Another line of evidence involves kinetic comparison of supposed hydroxyl radicals from several sources. Information, sometimes scanty, is available on the photolysis of hydrogen peroxide
The Nature of the Hydrated Electron
the Fenton reaction,
The electron, as first formed in water radiolysis, will have anywhere from several tenths to several hundred electron volts of kinetic energy. The larger energies are rapidly lost to electronic excitation of the water, and the lower energies to vibrational and rotational excitation, until at some low energy the electron enters the conduction band of water. This band, thoueh not directlv observed. is assumed to exist bv analorv -. with hydrocarbons and liquefied rare gases. The principal characteristic of electrons in conduction hands is high mobility, perhaps lo5 times larger than that of the hydrated electron. In water, the electron leaves the conduction band to become trapped but not fully hydrated in less than 2 ps and fullv hvdrated in about 4 DS (13). several studies have been made of possible chemical consequences of the highly mobile electron in water, sometimes called the "dry"e1ectrou (14). The maximum lifetime of the species (2 ps) is less than present instrumental time resolution, so direct observation is not nossible. The effect of molar concentrations of solutes on'the initial yield of hydrated electrons (at about 30 us) surrests that there is some reaction of dry electrons with s o l ~ t ~ , ~the h u tdry electron probably moves only a few molecular diameters before being trapped. The effect of high pressure on reaction rate constants, particularly on reactions (2) and (-2), leads to the conclusion that a cavity of about 10 cc per mole is associated withe-, (15). . ~The -~~ molar . volume of water is 18 cc Der mole, so there 1s an expansion t&ui\,alent tuabwr a half m'uleculea>ound each e-.. . The ionic cot~ductivitvof the hvdrawd electron has been measured to be 190 mhojcm (16),"which corresponds to a diffusion coefficient of 4.9 X 10-5 em2sec-l. almost the same as OH-. Thus the rate-limiting step in diffusion of ca9 is most likelv motion of the hvdmtine. waters. AL initio calculations (17) on an excess electron in a cluster of four water molecules surrounded by a dielectric continuum
Fe(l1) + H202 Fe(l1l) +OH- + OH the reaction of hydrated electrons with hydrogen peroxide
-
~~
~~
~~~
~
~
102
Journal of Chemical Education
e-.,
-
+ Hz02
-
O H + OH
(3)
and with nitrous oxide,
The radicals from these sources react with solutes with the same rate constants (or rate constant ratios) as do the oxidizine radicals of water radiolvsis. It would seem most reasonagle to conclude that they are all hydroxyl radicals. The OH radical yield may he estimated by equating it, with care, to the oxidation yield of Fe(CN)6-4 at low concentration, M as measured by pulse radiolysis technique. At about Fe(CN)6-4, the yield is 2.75 (21). The hydroxyl radical also has a UV spectrum associated with i t with a maximum around 230 nm. as seen in Fieure 1. The gas phase radical has absorption bahds around 3& mm. The enerev shift between the two bands. 1.3 eV. is rather too large to be accounted fur i ~ yhylruytn honding effects. It is rxAl,le that the 230 nm ~ e o kis charce rmnsfrr from solvent to solute, as was suggested for H andfor e,,-. The OH s ~ e c t r u mlike . the H atom spectrum, is of limited use for kinetic studies. In fact, the number of such studies for each of the three species whose sDectra are shown in Figure 1are in about the iame ratios as ihe areas under the curves. Manv rate constants have been measured by following the gnwih of product, particularly for uddit icm reactions. More commonly, pulse radiolytic aimpetition stud~esare used (22) (with I.'e(CNJ6-' to produre I.'e(CNJ6-", with 11CO~-to proi~ Competition ducr (?&-, and with SCN- to p n ~ l u c(S('NK). methods with Co8O radiation ;,re ulsu ovailahle (23) (using paranitrosodimethylaniline, for instance)
~..
0
-
10-10
10--3
10-8 10-7 TIME. sec
10-6
Figure 2. Time dependence of e-m (36)and OH (37)following production in a 30 ~6 electron pulse.
The pK of the OH radical has heen measured to be 11.9 by utilizing the fact that OH and 0- react with some solutes at different rates (24). For instance, OH oxidizes Fe(CN)a-4 and 0- does not. The apparent rate constant as a function of pH is determined by the pK of the radical. Molecular Products Molecular hydrogen, 0.45 molecules per 100 eV, and hydrogen peroxide, 0.70 molecules per 100 eV, are also formed in water radiolysis. These yields have been observed at high dilution of many solutes and are independent of radiation intensity. Product Yields The yields mentioned so far in this article are for solutions M concentration. These yields of reactive solutes a t about are assembled in Table 1where the total oxidation yield (2 GH~+ o ~GOH) and reduction yield ( ~ G H , GH Gem)are seen to be in excellent agreement. This means that any other intermediates are likely to he insignificant, unless they occur fortuitously in pairs. Specific searches for HOz and 0 atoms as orimarv ~ r o d u c t shave been made hut only vanishingly .. . small yields>ound. Exc~trdstates ufwaterwoulh not show up in the material halance.'l'hrrv i- Inosr likely a low-l\ina triplet state of water (251, but no chemical evidence forit Gas ever been found, even though i t has been looked for. The concept of radical yields becomes blurred a t much higher solute concentrations, because the radicals produced in water radiolysis are being intercepted a t earlier times after formation when more of them are still present. The free radical concentrations surviving as a function of time following a 30 ps pulse of electrons are shown in Figure 2. The decrease in radical concentration with time is mirrored in an increase with solute concentration of product yield from reaction with the solute, shown in Figure 3. These effectsare in qualitative agreement with predictions based on the spur and track model, a basic mechanism underlying all of radiation chemistry (26). The free radicals are formed inhomopeneousl~in spurs which are further organized into tracks. '?he rndicul concentrations are sufficiently high that radical-radical rivictirm ran compete u.ith diffusion oui of the spur. Some of the reactions expected to occur in the spurs are the sources of the molecular hydrogen:
+
.
e-,
+ e-,, (+ 2HzO) -Hz + 20He-, + H (+H20) -Hz + OH-
HtH-HP the source of molecular hydrogen peroxide:
+
,
I
lo-4
1
IO-~ S O L U T E CONCENTRATION
Figure 3. Fe(CN)sC3yield from OH oxidation as function of Fe(CN)sC4 in NIO saturated solution 121).and Nz yield horn e-- reduction as functim of N20 concentration (38).
and water reforming reactions
-
+ OH OHH+OH-Hz0 These are all known reactions with experimentally determined rate constants, diffusion-limited at about 10'O M-1 sec-' ( I ) . Examination of the above reactions will suggest that Hz yields are expected to decrease with increasing concentration of solutes which react with eaq- or H, and Hz02 yields should decrease with increasing concentration of solutes which react with OH, and such effectsare observed. Also, products of reactions with e-.q and OH with the added solute should increase with solute concentration, as shown in Figure 3. Also, a t high LET a t which the spurs are densely packed along the tracks, molecular hydrogen and hydrogen peroxide are expected to increase a t the expense of radical yields, as is exnerimentallv observed. ~ousideribleeffort has gone into calculations of spur effects (27). The unknown oarameters are the numher of radicals in spur and the dimensions of the spurs. Estimates of the numher distribution of radicals in the spurs have been made, and spur dimensions (10 to 20 A) can then be fixed by comparison with experiment. Good agreement is found between calculated effects and curves such as Figure 3, molecular yields, etc., hut the predicted time scale of the events is about an order of magnitude shorter than that observed in Figure 1. Agreement can be obtained by assumine. spurs con. larger . . taini;ly more radicals. Currently it w(n11d appear that the spur diff~tsionmodel must be essentially correct, but it has poorly understood areas, probably in the knowledge of spur sizes and secondary electron effects. The yield of a product of a solute-radical reaction at high solute concentrations is not easy to estimate o priori because effects are not additive. For instance, the yield of Np from reaction (4) in N20-saturated solution is 3.2, and since one OH radical is produced for each Nzand 2.75 are produced directly from water (Table 1)one might expect 5.95 radicals per 100 e-,
a
Table 1. Product Yleld In Water Radlolysis G H OH HI H202
2.65 0.65 2.75 0.45 0.70
Reducing Equiv.
Oxidizing Equiv.
2.65 0.65 2.75 0.90
a ?n
1.40 A is
Volume 58 Number 2 February 1981
103
Table 2. Heats and Free Energies of Formation
eV would he available for reaction with solutes. At low solute concentration tFig. 3 ) the ixp~rimentalnumber is 5.3, 12% !ess. The reason fur th~idifter~.nce ic that while N? is inert to further spur reactions, 0- and OH are not. The radical yields given in Table 1can he used to estimate product yields with about 10%accuracy as long as the product of the rate constant for the radical-solute reaction and the solute concentration is less than about 5 X lo7 sec-', which is less than 0.002 M if the reaction is diffusion limited. Measurements of radical vields as functions of temnerature suggesr that varintim-. lwtrv; t.n (1°C and 100°C are adnut luo; ,28,. Such retnperut..re ctlelfirient work is difficult to do Iw cause solutes n.hiam diw riminarr easily between radicals at X-01are much less discriminntinp. ruum temwrutun I - U V I ~ iii i at high temperatures. Secondary Radicals Nearly every reaction of e-.,, H, or OH produces a secondary radical of some sort, and hundreds have been studied. A few are of general interest for various reasons; some because they have lower redox potentials and cause more specific reactions than the primary radicals (such as (SCN)2-, Clp-, Brz-, 12-, TI2+,COs-, C02-, etc.), and some because they can be made by all three primary radicals and are thus obtained in pure form (HOz, 02-, C02-). In the latter category, HOZand 0 2 - are of special interest because of their obvious involvement in biological processes. 0 2 - can be prepared from all radicals by use of air- or oxygen-saturated formate solutions:
-- - ++ +
H+02+H02 ec.,
+ 0%
OH + HCOO-
01
Hz0 COzco2-+ 0 2 COz 0 2 HOz 3 5 H+ 0 2 All of these reactions arewmplcte inlessthan 10-6secinnirsaturated, 10-:'M formatc solution. The pK of HOz is 4 . 8 . ~ in neutral and I ~ n s i c s ~ ~ I uit texist~ w ~ as 02-. The second-order interesting (29).as the rate constant for decav oI'0,. is uuiw . the reaction HOP+ HOz- H202 + 0 2 is 8 X 10" M-' sec-', and 9 X lo7 M-' sec-' for HOz + 0 2 - -Hop- + 0 2 while O2- Oz- does not react at all (k less than 0.3 M-' sec-1). As a result, w4M O2- at pH 10 and above will have a half-life of more than 10 sec (with care to prevent trace metal contamination of the sample). Consequently, Oz- can be generated by water photolysis (30), as well as by radiolysis, and subsequent reactions studied in a stopped-flow appmatus. It is very simple to work with, but it is very sluggish in its reactions. In fact, it is quite an event to find a compound with which it will react. The absorption spectra of HOz and 0 2 - are in the UV (240 and 260 nm) and so are sometimes not useful for kinetic studies. In such cases product growth must be followed.
+
Thermodynamics Estimates of free energies of formation and some heats of formation for era,, H, OH, 0 - , HOz, and 0 2 - a t 1M , 25'C standard state in aqueous solution are given in Table 2. Only HOz and Oz- have been measured directly (31). 0- can be related to OH via the pK of OH (11.9), and e - , can be related to H via the pK of H, 9.6. The heats and free energies of formation of H and OH are known in the gas phase, so free energies and heats of solution are required and have been estimated by comparison with stable analogues. The free energy of solution of He and Hz are 4.6 and 4.2 kcal/mole (that is, the concentrations in equilibrium with one atmosphere of the gas are 4 X lo-' and 8 X 10-'M, respectively), and the heats of solution are -0.4, and -1.0 kcallmole. Solution of H 104
Journal of Chemical Education
AQU~OUS Species
AH,"
H
AGO
53
51
66
Bag
-1
OH 0-
4.5 23
3
Table 3. Half-Cell Potentialss
+ +
OH H+ e- = H20 H+H++e-=Hp OH e- = OH0 ~ ' 2Ht e- = H,02 H HzO e- = H, OHH02 H+ e- = H& 02- H+ e- = H0.H02 e- = H0.Hf+02+e-=HO. O2 e- = 0.Hf+e-=H e- = e- 4
+
+
+
+
+ + + + + + +
a
+
2.65
2.30 1.83
1.73
1.47 1.44 1.02
0.73 -0.04 -0.33 -2.30 -2.87
H20(lig),H2 lg, 2S0. 1 atrn), 0% (g. 2S'. 1 atrn) as standard stales.
should be similar to these (32), so the free energy of solution is assumed to be 4.4 kcal/mole, and the heat of solution to be -0.7 kcallmole. The free energies of solution (from the gas) for CHsOH, CzHsOH, and Hz0 are -3.2, -3.2, and -4.4 kcallmole. The average, -3.7, is probably reasonable for OH (33). and the heat of solution should be about -10 kcal. The validity of this method may be checked once by estimating AGn for aaueous HO? from the eas ~ h a s bv e usine the free energy of sblution of HZOZas a midel: he result is within 0.5 kcal of the exnerimental value. The data o i ~ a b l 2e can be used to construct the half-cell ~otentinlsof'l'able3.It isseen that thr rnnrrr k v e w wide (2.65 to -2.87). It is not surprising to find OH as the sGongest oxidant. What is surprising is that hydrogen atoms are nearly as good oxidizing agents as OH. This potential is realized in hydrogen abstraction reactions, but otherwise oxidation by H is limited by available mechanisms. Termolecular reactions involving Hs0+ would be too slow, Hz+ is not stable relative to H and H30+, and the half cell potential for hydride ion formation is negative. The only oxidation reactions of H other than abstraction are those forming hydride intermediates which then react with water or acid to produce HZ(34): reaction with Fez+ eives FeH2+. Cr2+ eives CrH2+. and Ti(II1) . . gives TiH(II1). The full OH ~otentialcan be realized onlv in H abstraction to form water,'but the oxidation to form OH- is strongly positive (1.83 v), and also in the oxidation of inorganic ions there is often the possibility of forming OH coordination compounds. This route is difficult to demonstrate because subsequent acid-base reactions of the OH complex are usually very rapid. The question of possible amphoteric behavior for OH and HOz has been raised several times. The systematics of acidbase behavior of free radicals which have fewer H atoms than some parent demonstrate that the pK of the radical is always several units lower than the parent (i.e., the pK of OH, 11.9, is about 4 units lower than HzO, 15.7).This dependence of pK on the H I 0 ratio is well known for all inoreanic acids. A o ~ l i + 4 cation to OH would suggest that the pK O ~ H Z Oshouldbe units lower than H@+ (-1.7) or -5 to -6. The pK of HzOz+ should be less than -10 by comparison with HsOz+. Consequently the acidic forms, HzO+ and H202+. are very unlikely participants in reactions in aqueous solution. In conclusion, water radiolysis provides a starting point in the synthesis of many radicals and radical ions in aqueous
-
solution. The primary radicals, ecaq, H, OH, are well characterized. The radical population can he made to he 90% pure OH (or 0-1 if N 2 0 solutions are irradiated, the remaining 10% being H atoms. 55% of the radicals can be converted to H radicals) are difficult toconvert to H by reaction wiih Hz, due to the slow rate of the reaction. About 100 atmospheres of H2 are required to do the conversion in less than loc6 sec. Usually some thought on the part of the investigator can lead to a system with a known fate for the OH which will prevent interference with the radical of interest. Acknowledgment
This article was written a t Brookhaven National Laboratory under contract with the U S . Department of Energy and supported by its Division of Basic Energy Sciences. Literature Cited
S., Matheron, M.. Rahani, J.. and Thomas, J. K.. J A m i r Chem Snt, 85, 1375 il%li ,.""",.
(101 Rahani, d., ~n %dvated ~ l o c t r o n . " ~ d o othem. n srr.SO. R. F. C,,,,in mdiinri~ .r , 2, Amer. Chem. Sac, Warhin~tun,l).C.11945). I111 Nielren, S. 0.. Michael, H. I>.. and HarL. E.J..J P h w Chem. 80.2482 119761.
.
~~
,.",",. I191 McMiUan.J.A..Matheron, M.S..andSmaller,B., J. Cham. Phys.. 33,609 (19601. 1201 Behar,D.. and Fessenden,R. W., J. Phyn Chem.. 76,17M (1972). 121) Schuler, R. H.. Hartzell, A. L.. and Rehar, B., submitted iur publication. 1221 Dorfman, L. M.. and Adams. G. E., "Reactivity of the HydrorylRsdiesl in Aqueous Solutions? NSRDS-NBS 46 (see Ref. 1). (23) Kraljic,I..andTrumhnre,C. N..J . Amrr Chrm. Soc.. 87,2547 (1965). (241 Rabani,J.,and Matherun, M. S.,J Amen C h r m Soc., 85.3175 11964). (25) Tlsjmar, S., Williams, W . , a n d Kupgerman. A., J . Chem Phys. 54,2274 (19711. (26) Draganit.!. G.,and Draganit, 2. D..'TheRadisfionChemistryofWafer?Aeademic Pros3,New York, 1971. 127) Kupperrnan, A., in "Radiation Reiesrch," G. Silini. IEditor1, North Holland Publishing Co..Amsterdam, 1912.p. 212. 1281 Hochanade1.C. J.,sndGhormley,J. A.,Rild&at.Reseorch,15.65311962):Michael.B. D.,Hait, E. J.,and Schmidt. K. H . J . Ph'hys.Uhem, 75,2796 (1971). (291 Bie1ski.B H. J.. and Al1en.A. 0..J. Phyr C h m . 81,1048 (19771. 130) Hulroyd,R. A., and Bie1ski.R. H. J., J A m w Chem Sor.. 100,5796 11978). 131) 1lan.Y.A.. Czapski.G.,andMeisel,D.,Binthim Riophys. Acta, 430.209 (1978):Divesek, J., and Kaatening. B., J . Elsrlroanal. Chem., 65,603 (1975). 132) Taffel. 1% and Henglein. A , Forodoy Disc. Chem Soc.. 63,124 (19771. (38) Baxendale, J . H.. Ward, M. D., and Wardman, P., Tions. Foradoy Sac, 57, 2532
,."..,.
,1971,
(34) Clapski. G.,Jorfner, J.,andStein,G.. J . Phys Chsm.. 65,96011961):Cohen, H..snd MeyerstehD., J. C h r m Sor. Dalton, 1974,2539; Behar,D.,and Samini,A..Chrm Phys. Lett., 22,105 119731. (351 Keene, J. P., Radial Research, 22, i 1197%): Gordon, S.,andHan,E.J., J. Amei Chem. S u i , 86,5343 119641. (36) J0nah.C. D., Matheson. M. S..Miller, J. R.,andHart.E.J. J P h y s . Chem., 80,1267 11976). (37) Jonah, C. D., and Miller,J.R., J . Phys Chrm., 81, 1974 (1977). (36) Dainton, F. S., and Logan. S. R.. Trans. Foiaday Sor.. 61,715 (19651.
Volume 58
Number 2
February 1981
105
