The Journal of Physical Chemistry, Vol. 83, No. 4, 1979
Equilibrium Reaction of Electrons with p-Difluorobenzene
have calculated back the rate constant hlb from the initial slope of the NH(32:-)absorption records. The hand values listed in Table I11 were such data obtained by using the absorption coefficient 1.0 X IO7cm:! molP. The Arrhenius plots of hlb are as shown in Figure 8. The least-squares treatment gave the expression klb
(IO) T. Fueno, K. 'Tabayashi, and 0. Kajimoto, J . fhys. Chem., 77, 575 (11) (12) (13) (14) (15) (16) (17)
= 1013.8"*0.35 exp[-(31.2 f
2.0) kcal mol-l/R'T] cm3 mol-l s-l (VII)
(18) (19)
Although the errors involved in the evaluation of the initial slope are relatively large, eq VI1 is in good agreement with the lower-bound hlb (eq VI) obtained from the rakes of HN3 disappearance. The proposed initial decomposition kinetics and the subsequent sequence of reactions are both perfectly compatible with the observed HN3 and NH(32-) absorption profiles.
(20) (21) (22) (23)
Acknowledgment. This work was assisted in part by a grant-in-aid (No. 2417069) from the Ministry of Education, Japan. Reference!$and Notes (1)
435
(24) (25) (26) (27) (28) (29) (30)
H. S.Berry in "Nitrenes", W. Lwowski, Ed., Interscience, New 'fork,
1970, p 13. (2) ti, Okable, J . Chem. fhys., 49, 27261 (1968). (3) F?, J. Paur and E. J. Bair, J . Photochem., 1, 255 (1972/73); Int. J . Chem. Kinet.. 8. 139 (1976). (4) F?. S.Koinar, S.Matsumoto, and B. de B. Darwent, Tcans. Faraday Soc., 67, 1698 (1971). (5) J. R. McDonald, R. G. Miller, and A. P. Biaronavski, Chem. Phys. Lett., 51, 57 (1977). (6) W. C. Richardson and D. W. Setser, Can. J. Chem.,47, 2725 (1969). (7) I, S.Zaslonko, S. M. Kogarko, and E. Y. Mozzhukhin, Kinet. Katal., 13, 289 (1972). (8) C. B. Moore and K. Rosengren, J. Chsm. fhys., 44, 4108 (1966). (9) J. R. McDonald, J. W. Rabalais. and S.P. McGlynn, J. iChem. fhys., 52, 1332 (1970).
(31) (32)
(1973). J. M. Lents, J . Quant. Spectrosc. Radiant Transfer, 13, 297 (1973). K. Dressler and D. A. Ramsay, fhii. Trans. R. SOC., A251,553 (1959). R. N. Dixon, Can. J . Phys., 37, 1171 (1959). G. H. Dieke and R. W. Blue, fhys. Rev., 45, 395 (1934). W. Tsang, S.H. Bauer, and M. Cowperthwaite, J. Chem. Phys., 36, 1768 (1962). H. Kijewski and J. Troe, Int. J . Chem. Kinet., 3, 223 (1971). H. G. Schecker and W. Jost, Ber. Bunsenges. fhys. Chem., 73, 521 (1969). K. W. Michel and H. Gg. Wagner, Symp. (Int.) Combust., [Porc.], loth, 1964, 353 (1965). H. A. Olschewski, J. Troe, and H. Gg.Wagner, 6er. Bunsenges. fhys. Chem., 70, 450 (1966). J. Troe and H. Gg. Wagner, 6er. Bunsenges. fhys. Chem., 71,937 (1967). P. J. Robinson and K. A. Holbrook, "Unimolecular Reactions", Wiley-Interscience, London, 1972. G. Z. Whitten and B. S.Rabinovitch, J. Chem. fhys., 38,2466 (1963). The AHoo values were either taken directly from or calculated on the basis of the data compiled in the following: D. R. Stull and H. Prophet, "JANAF Thermochemical Tables", 2nd ed, National Bureau of Standards, Washington, D.C., 1971; M. W. Chase, J. L. Curnutt, A. T. Hu, H. plophet, A. N. Syverud, and L. C. Walker, J. fhys. Chem. Ref. Data, 3, 311 (1974); M. W. Chase, J. L. Curnutt, H. Prophet, R. A. McDonald, and A. N. Syverud, ibid., 4, 1 (1975); R. F. Harnpson, Jr., and D. Gnrvin, Natl. Bur. Stand. Tech. Note, No. 866 (1975). = 72.66 kcal/mol was calculated from The value of AH,",(HN,) AH,",,(HN,) = 71.66 kcal/mol which was given by P. Gray and T. C. Waddington, R o c . R . SOC. London, Ser. A , 235, 106 (1956). H. S.Johnston and C. Parr, J . Am. Chem. Soc., 85, 2544 (1963). S. W. Bsnsori, "Thermochemical Kinetics", Wiley, New York, 1966. G. M. Meabum and S.Gordon, J . Phys. Chem., 72, 1592 (1968). G. Le Bras and J. Combourieu, Int. J. Chem. Kinet., 5, 559 (1973). M. H. Hanes and E. J. Bair, J . Chem. fhys., 38, 672 (1963). S. Gordon, W. Mulac, and P. Nangia, J. fhys. Chem., 75, 2067 (1971). P. V. Khe, J. C. Soulignac, and R. Lesclaux, J . f h y s . Chern., 81, 210 (1977). D. L. Baulch, ID. D. Drysdale, D. G. Horne, and A. C. Lloyd, "Evi3luated Kinetic Data for High Temperature Reactions", Vol. 1, Butterworths, London, 1972. W. H. Smith, ,I. Brzozowski, and P. Erman, J. Chem. fhys., 641,4628 (1976). w, "Quantitative Molecular Spectra and Gaseous , Addison and Welsley, Reading, Mass., 1959.
Reversible Reaction of Excess Electrons with p-Difluorobenzene in n-Hexane and Cyclohexane Richard A. Holroyd,* Chemistry Department, 5rookhaven National Laboratory, Upfon, New York 1 1973
Richard D. McCreary, and Georcge Bakale Division of Radiation Biology, Radiology Llepartment, Case Western Reserve University, Cleveland, Ohio 44 106 (Received September 27, 1978) Publication costs assisted by Brookhaven National Laboratory
The equilibrium reaction e; + p-CsH4FzG p-CsH4Fz-was studied as a function of temperature in n-hexane and cyclohexane solvents. The conductivity technique employed allowed measurement of both the attachment and detachment rates. The attachment rates are diffusion limited in both solvents. The detachment rates are characterized by very high preexponential factors and large activation energies. The reaction is exothermic in solution and proceeds with a large decrease in entropy. The latter can only partially be accounted for by polarization by the product anion, but in addition indicates a positive entropy of solution of the electron. A linear correlation of AG,,,, with electron affinity is presented. The results are consistent with an electron affinity of p-difluorobenzene equal to approximately -0.35 eV.
Introduction The equillibrium reaction (eq 1)of trapped electrons with e,-
+ p-C6H4F2r p-C6H4F2k2 k1
(1)
p-difluorobenzene has been observed in solution by a pulse 0022-3654/79/2083-O435$01.OO/O
conductivity technique. Previously, similar equilibria. have been reported for reaction of the electron with carbon dioxide,l and aromatic hydrocarbon^^,^ in nonpolar solvents. The reaction of solvated electrons with nnonofluorobenzene in liquid ammonia is suggested to involve such an eq~ilibrium.~ In this study with difluorobeinzene 0 1979 American Chemical Society
430
R, A, Holroyd, R,
The JOUrnRl o l Physlcal Cksml$try, Vol. 83, No, 4, 1979
two hydrocarbon solvents, n-hexane and cyclohexane, were employed. These solvents are not generally considered to be electrophilic, yet electron detachment from C6H4Ff to the solvent is facile and occurs thermally a t room temperature in spite of a large activation energy. Electron spin resonance studies have established that fluorinated benzene anions can be stabilized in an adamantane matrixaBThe ESR spectra show clearly that the excess electron in the p-difluorobenzene anion is in a T* orbital. This contrasts with the more heavily fluorinated benzenes in which the excess electron occupies a B* orbita1,6i6similar to the excess electron in perfluorocycloalkane anion^.^ The thermal stability of such anions appears to be a function of the extent of fluorination: whereas C6F0- was observed up to 240 K, the p-difluorobenzene anion was detected a t 110 K. The conductivity technique employed here permitted measurement of both the attachment and detachment rate constants for reaction 1 individually. From these data at various temperatures the thermodynamic parameters were evaluated. Previous results had shown that the free energies of these equilibrium reactions are related to the gas phase electron affinitiesO3The present results are compared with the previous data to derive a relationship between the gas phase electron affinities and the free energies of reaction observed in solution. Experimental Section The pulse conductivity technique described earlier was used for all m e a s ~ r e m e n t s . ~A, ~15-11s pulse of 0.9-MeV electrons was used for irradiation with the beam current adjusted to yield an initial electron current of -300 PA with a field of 30 kV/cm applied to the parallel plate electrodes of the ion chamber. Aldrich 1,4-difluorobenzene with a stated purity of 98% was added to n- or cyclohexane (Fisher 99 mol %) to prepare approximately millimolar base solutions, Aiiquots of freshly prepared base solutions were diluted to 1-6 pM solutions in n- or cyclohexane that was purified as described previou~ly.~ The sample solution was poured into the argon-purged ion chamber and argon was bubbled through the solution €or 10-15 min prior to each measurement. The chamber was refilled with fresh solution prior to measurements at each temperature. The solution temperature was controlled by a thermoelectric device encasing the ion chamber and the temperature was measured with a thermocouple attached to the chamber wall. Oscilloscope traces of the electron current decay were analyzed by a computer methoda3 The electron current per unit electrode area is a function of the concentration of electrons:
where u, is the drift velocity of the electrons and d the electrode spacing. Since electron capturing impurities are almost always present, reaction 3 must also be considered. e-
+S
ks +
S-
(3)
Thus the continuity equation for electrons is an,/at = - ~ , ( d n , / a ~ )- hln,[DFB] - h,n,[S] + hzLDFB-1 (4) and the equation for DFB- is a[DFB-]/at = hln,[DFB] - hZlDFB-1
(5)
D,McCreary, and G,Bakale
n-HEXANE
1 OO"
0
.0,4 (usec)
7 0.6
' 0.8
Figure 1. Relatlve current vs, tlme In n-hexane solutlons of p-dlfluorobenzene. Experlmental polnts are from oscilloscape traces: ( 0 ) 5 '6,6.2 pM p-DFB; (0)20 O C , 4.1 pM p-DFB; (t) 25 OC, 2.9 p M p-DFB. Solid lines were calculated as descrlbed In text uslng the fallowing rate constants: (5 "C) k l 7,2 X I O " , k2 = 2.7 X I O 6 , k3 = 2.7 X lo6;(20 "C)k i = 1.2 X lo1*,k2 = 2.0X loe, k3 = 3.2 X lob;(25 O C ) k i = 1.2 X IO'*, k, = 3,5 X loe, k3 8 3 X I O 6 . 1
TABLE I : Rate Constants for e; t p-C,H,F, "p-C,H,F,' 2
5.0 10.0 20,O 23.0 25.0 28.6 7.0 9.0 20.0 25.0 28.0
n-Hexane 2.6 X l o 5 4.7 x lo5 1.7 X l o 6 1.9 X l o 6 3.6 X l o 6 3.25 X 10' Cyclohexane 2.5 X lo'* 5,2X lo5 2.0 x io*' 5.5 x 105 2.6 X 1Ola 2.7 x l o 6 2.5 x IOLa 4.7 x l o 6 2.5 X 10" 7.2 X lo6 7.0 X 10" 8.3 x 10" 1.1 x l o L 2 9.1 X 10" 1.25 X 10'' 8.4 X 10"
2.69 X 1.76 x 6.50 X 4.80 X 3.50 X 2.57 X
lo6 lo6 lo5
lo5 lo5 10'
4.82 X 10' 3 ,x i~o 6 9.4 x l o 5 5.3 x l o 5 9.0 X los
The solution of these coupled partial differential equations has been given,3 and n,(x,t) is a function of kl,kz,kB,u,, and d. The last three are fixed by experimental conditions; k&S] is measured from the current decay in the pure solvent, T o fit an oscilloscope trace requires calculating the current from eq 2 for a series of times after the pulse for a given kl and k2 and comparing to experiment. In practice the calculations were done on a PDP-11 computer with program EYEFIT; a Tektronix 613 storage display was utilized to compare the calculated curve to the experimental points. The rate constants were adjusted until a good fit was obtained. Examples of such calculated curves and experimental points are shown in Figure 1. Results Typical oscilloscope traces show a fast decay followed by an equilibrium current. The magnitude of the equilibrium current increases with temperature (see Figure l), an observation which is typical of these equilibria. The points in Figure 1 are experimental and the solid lines are computer calculated fits. The rate of the fast decay is very roughly given by hl[DFB] h2 + h3[S]and the height of the equilibrium current increases with kz. Values of kL and hz derived from the computer fits are given in Table I for various temperatures in both n-hexane and cyclohexane. The values are averages of a t least two runs at, each temperature. The electron attachment rate in n-hexane is approximately 0.9 X 1OI2 M-l s-l and in cyclohexane 2.5 X 10I2M-l s.-I. These rates do not change significantly over the limited temperature range investigated.
+
The Journal of Physical Chemistry, Vol. 83, No. 4, 1979 437
Equilibrium Reaction of Electrons with p-Difluorobenzene
TABLE 11: Thermodynamic Quantities for e; -kcal/mol
IO7
kcal/mol
-7.60 -1.75
n-hexane cyclohexane
- 16.8
r-11 t
i
c
1
I
lo5
--L...L LI 3.3
3.4
3.5
3.6
3.7
I / T x 1000
Figure 2. Plot!j of log Kvs. 1/T, where K = k l l k 2 ,for p-DFB in (0) n-hexane and ( 0 )cyclohexane.
The detachment rates on the other hand are very temperature sensitive. Least-square fits indicate the data are represented by h2 = 3.0 X lozo exp(-19200/RT) in n-hexane arid kz = 3.3 X loz2exp(-21600/RT) in cyclohexane. That is, the reactions are characterized by large activation energies and very large preexponentral factors. These equilibria can only be observed in a limited temperature range because of the following: if hz is too small, the eiquilibrium current is indistinguishable from baseline; and if kz is too large, the two components of the decay are not resolved. In practice it is also necessary to adjust hl[DFB] to be >> k3[S]; thus kl[DFB] is generally in the range 3-8 X lo6 s-l, anid in general 0.05 < h2/hl[DFB] < 2. Figure 2 shows plots of log K vs. 1/T for both solvents. The enthalpies of reaction 1 are derived from the slopes of the lines and the standard free energies from the values of K a t 25 'C. These values of AH",,,, AGor,,,, as well as AS',,,, are given in Table 11.
Discussion If electron attachment is diffusion controlled then the rate constant should be given by the Smo1uchow;ki-Debye equation: k l = 4ir(D, + DDFB)R,where the encounter radius ( R ) iEi given Iby the electron-induced dipole interaction energy. Previously it was shown that the rates of attachment to a series of nitro-substituted benzenes in cyclohexane are given by the above equation and are therefore difffusion limited.g For solutes with no dipole moment the rate constant was found to be 2.7 X lo1*M-l s-'. This value is very close to that observed in this study in cyclohexane. Similarly the observed rate of electron
A S " reac
AS,
*
cal/deg mol
19.2 21.6
-21.3
r-7
-
AE~*
A H 0reac
A G'reac,
solvent
I o7
t p C 6 H 4 F 22 p-C6H4F2-
-30.9 -45.5
34.2 43.5
attachment to p-difluorobenzene in n-hexane is very close to the rate of attachment to biphenyl and perylene (equal to 1.0 X 10l2M-l s-l) in this solvent and these reactions have been shown to be diffusion-controlled reactions.1° It is concluded that electron attachment to p-difluorobenzne is diffusion limited in these solvents. The fluorinated benzene anions are stable a t low temperatures5 but the possibility should be considered that they may dissociate a t higher temperatures. ESR studies indicate carbon--fluorinebond rupture does occur in certain compounds following electron capture.ll The occurrence of dissociation or any other reaction which removes DFB(eq 6) would kinetically be equivalent to increasing h3 upon DFB-
-
products
(6)
the addition of difluorobenzene because the effect of such a consecutive reaction is to increase the rate of decay of the equilibrium ~ u r r e n t .That ~ is, the electron current decay is the sum of two exponentials as shown in Figure 1. The rate of the slower (equilibrium) decay is very approximately given by (hzh3[S]+ hl[DFB]k6)/(kl[DFB] hz h3[S] k6). Since the computer fits indicated no significant change in h3 on addition of difluorobenzene, it is concluded thait dissociation of the anion, or any reaction of DFB- other than detachment, is of minor importance under these experimental conditions. The detachment rates are characterized by high activation energies (G*) of 19.2 and 21.6 kcal/mol in nhexane and cyclohexane, respectively. The actiwtion energies are very comparable in magnitude to the observed values of the enthalpies of reaction (AH'), but opposite in sign (Table 11). With such high activation energies detachment would be expected to be very slow a t room temperature. What makes the reaction fast is the high preexponential factors. If these factors are interpreted in terms of transition state theory, where the frequency factor is taken to be 1013,the activation entropies (AS*) are 34 and 44 cal/deg mol in n-hexane and cyclohexane, respectively. Note that the magnitudes of the entropies of activation are very close to the observed entropy changes for reaction 1 (Table 11). Large positive entropies of activation were ;also observed for detachment from styrene large and a-methylstyrene anions in n - h e ~ a n e .These ~ entropy changes are associated in part with the polarization of solvent molecules around the anions but also with a positive entropy of solution of the electron in hexane and cyclohexane (see below). The equilibrium constant for reaction 1 is 3.5 X lo!! M-l in n-hexane a t 25 O C and the value of AG0,,,,(1) is --0.33 eV. Thus the reaction is exothermic and the equilibrium is strongly shifted to the right. These equilibria are more exothermic in solution because of the extra stabilization gained in polarization of the solvent by the product anion. That is, AGoreac(l) is given by (6), where the free energy
+ +
+
AGore,,(l) := -EA
+ AGoPo1(DFB-)- AGosol(e-) (6)
of solution of the electron, AGosol(e-)must also be taken into account. The polarization energy of CGHstin rare gas matrices has been experimentally measured and is -0.93 eV in an argon matrix.12 This is within 5% of what one would estimate from the Born equation: A G , = -(e2/2r)(l
438
The Journal of Physical Chemistry, Vol. 83, No. 4, 7979
R. A. Holroyd, R. D. McCreary, and G. Bakale
- l/c), assuming an effective radius calculated from the molar volume. For difluorobenzene this radius would be 3.38 A and we can estimate AGoPo1= -1.0 eV in n-hexane and -1.08 eV in cyclohexane. The enthalpy of reaction, AHoreac(l), is even more negative; in n-hexane AH",,,, = -0.73 eV compared to AG",,, = -0.33 eV. This indicates there is a large negative entropy change associated with reaction 1. By analogy to eq 6 we can write
AHoreac(l) = -EA
+ AH",,l(pDFB-)
-
--IT-
i
?
-4 Nti
1
AHosoln(e-) (7)
The enthalpy of polarization by the difluorobenzene anion differs from AGwl by the TAS,,, term. An estimate of the polarization entropy is obtained from the Born equation:
AS,,, = -aGPoI/3T = ( e 2 / 2 r e 2 )at.jaT
(8,
The variation of the dielectric constant with temperature is used to calculate AS,,l, thus AHpol(DFB) i s - 1.26 eV in n-hexane and -1.33 el7 in cyclohexane. The reaction is more exothermic in the liquid than in the gas by approximately this amount, depending on the magnitude of AHsoj(e-),which is not B large term (see below). As stated above, these reactions involve a large negative entropy change which is associated with the polarization of the medium by the difluorobenzene anion but also includes a positive entropy of solution of the electron. If we subtract eq 6 from 7, we obtain
Values of AS,,, are estimated from ey 8 to be -20 eu in n-hexane and -19 eu in cyclohexane. Thus a large share of the observed entropy of reaction can be accounted for by the entropy of polarization. However, putting these values into eq 9 we find ASsoln(e-)is 4-11 eu in n-hexane and $26 eu in cyclohexane. The value for n-hexane is in good agreement with the value of $12 f 4 eu reported earlier.3 Comparable positive entropies of solution have been reported of 4-13 and +16 eu for amrnonia,l3J4 and 4-15.6 eu for water.I5 The positive entropy of solvated electrons can be associated with the large number of states available to the electron in a disordered liquid. If the estimated polarization terms along with the observed free energies of reaction are substituted into eq 6, one obtains EADFBis equal to -0.67 the following results: AG,,,(e-) eV in n-hexane and -0.74 eV in cyclohexane. In a similar equal to -0.53 eV in way one obtains: AHsOh(e-)4- EADFB n-hexane and -0.40 eV for cyclohexane. The energies of solution reported earlier3 for n-hexane were AHSol,(e-) = -0.18 eV and AGsoln(e-)= -0.33 eV. These values combined with the above results yield -0.34 eV for the electron affinity of p-difluorobenzene. Recent reported values are higher than this; an electron transmission experiment indicated the electron affinity of p-difluorobenzene to be -0.54 eV.lG In a similar experiment the lowest unoccupied T state was found a t -0.53 eV.17 This should be the lowest state since ESR evidence indicates the excess electron is in a x orbital. On the other hand, the thermal electron attachment technique yielded a positive value for the electron affinity of p-difluorobenzene of 4-0.18 Our results are inconsistent with a positive electron affinity. In an earlier study of the electron equilibrium reaction with styrene the value of AGO,,,, was reported to be -0.56 eV in n-hexane a t 25 'C3 Since A G O , , , , for p-CGHdF2 is -0.33 eV it follows that the electron affinity of p-difluorobenzene is more negative than that of styrene, and styrene has a negative electron af-
+
I
1-0.3 + 0 2
to1
0 -0.1 E A eV
-0.2 -0.3 - 0 4 -05
Figure 3. Free energy of reaction vs. electron affnty: ( phenanthrene, naphthalene, biphenyl, styrene, and a-methylstyrene in tetramethylsilane (ref 3 and 20); (0) naphthalene, styrene, and amethylstyrene in 2,2,44rimethylpentane (ref 3); (0) styrene, amethylstyrene, and p-difiuorobenzene in n-hexane (this work and ref
3).
finity.lg Finally if the value of the electron affinity of p-difluorobenzene a t -0.34 eV is used to calculate the energies of solution of the electron in cyclohexane one obtains AGsoln(e-)= -0.2 eV and AHRoln(e-)= 4-0.14 eV. Equation 6 suggests a linear relationship between free energy and electron affinity. This would only be linear for molecules of approximately the same size so that the polarization energies would be comparable. Figure 3 shows such a plot where data from other studies are combined with the present data. For a series of six aromatic molecules studied in tetramethylsilane (solid circles) there is a rough linear relationship. The data for 2,2,4-trimethylpentane and n-hexane fall on different lines because the value of the free energy of solution of the elect,ron is different in each solvent. Clearly more data are needed to establish these lines; however, we may use this figure to estimate equilibrium constants for other compounds. For example for the reaction of the electron with benzene in n-hexane, the graph indicates that the equilibrium constant would be a t 25 O C ; that is, no reaction is expected. Acknowledgment. The authors thank Dr. IJ. Sowada for stimulating discussions. This research was supported by the U.S. Department of Energy under Contracts EP7843-02-4746and EY-76-C-02-0016. References and Notes R. A. Holroyd, T. E. Gangwer, and A. 0.Allen., Chem. fhys. l.ett., 31. 520 (1975). J. M. Waiman, M. P. de Haas, E.Zador, and A. Hummel, Chem. fhys. Lett., 35, 383 (1975). R. A. Holroyd, Ber. Bunsenges. Phys. Chem., 81, 298 (1977). E. San Roman, P. Krebs, and U. Schindewolf, Chem. f h y s . Lett., 49, 98 (1977). M. B. Yim and D. E. Wood, J . Am. Chem. Sac., 98, 2053 (1976). M. C. R. Symons, J . Chem. SOC , Chem. Commun., 408 (1977). A. Hasegawa, M. Shiotani, and F. Williams, Faraday Discuss., 63, 157 (19771. G. Bakale, E. C. Gregg, and R. D.McCreary, J . Chem. Phys., 57, 4246 (1972). G. Bakale, E. C. Gregg, and R. D. McCreary, J . Chem. Phys., 67, 5788 (1977). J. H. Baxendale, B. P. H. M. Geelen, and P. H. G. Scharpe, Int. J . Radiat. fhys. Chem., 8, 371 (1976). L. D. Kispert, ACS Symp. Ser., No. 66, 349 (1978); K. Toriyama and M. Iwasaki, J . fhys. Chem., 76, 1824 (1972). J. Jortner and A. Gaathon, Can. J . Chem., 55, 1801 (1977). Farhataziz and L. M. Perkey, J . fhys. Chem., 80, 122 (1976). G. Lepoutre and A. Demorter, Ber. Bunsenges. fhys. Chem., 75, 647 (1971). A. G. Ryabukhin, Russ. J. fhys. Chem., 51, 573 (1977).
Reduction of Co(bpy);+
and Co(phen);+
The Journal of Physical Chemistry, Vol. 83, No. 4, 7979 439
in Solution
(16) P. D. Burrow, private communication. (17) J. R. Frazier, L. G. Ghristophorou, J. G. Carter, and H. C. Schweinler, J . Chem. Pbys., 69, 3807 (1978). (18) S.Lin, Ph.D. Thesis, IJniversityof Houston, 1973. University Microfilms, Internatioiial, Ann Arbor, Mich.
(19) P. D. Burrow, J. A. Michejda, and K. D. Jordan, J. Am. Cbem. Soc., 98, 6392 (1976). (20) J. M. Warman M. P. de Haas, and A. Hummel in "Conduction and Breakdown in Dielectric Liquids", J. M. Goldschwartz, Ed., Delft University Press, 1975, p 70.
One-Electron Reduction of Tris(2,2'-bipyridine) and Tris( 1;IO-phenanthroline) Complexes of Cobalt(II1) iin aqueous Solution' M. G. Simic," M. 2. Hoffman,*2b R. P. Cheney,2band Q. G. Mulazzani*2c3d Food Engineering Laboratory, US. Army Natick Research and Development Command, Natick, Massachusetts 0 1760, Department of Chemistry, Boston University, Boston, Massachusetts 022 15, and Laboratorio di Fotochirnica e Radiazioni d' Alia Energia, Cons@/ioNazionale Delle Ricerche, 40 726 Bologna, Italy (Received August 16, 19;'8) Publication costs assisted by the National Science Foundation
The reaction of Co(bpy)QB+with radiation-generated reducing radicals (ea;, .C02-, (CH3)&0H,and CH,OH) in aqueous solution quantitatively and rapidly ( h = 108-1010M-' s-') yields high-spin (ti, e:) C ~ ( b p y ) (A,~, ~ + 300 nm, e,,, 4.2 X IO4 M-' cm-l) which slowly ( k = 3.4 s-' at pH 0.5-10.5; 8.0 s-l at pH 0.3) equilibrates with the loss of bpy; C ~ ( b p y ) ~does ' + not transfer an electron to O2 or p-benzoquinone. Equilibrium mixturles of Co(bpy):+ react with ea; (h > 1O1O M-l s- ) to produce Co(bpy),+ (Amc 620 nm, emax 5.1 X IO3M-' cm-'). Reduction of Co(phen)QB+yields C ~ ( p h e n ) (A~, ~ + 270 nm, e,, 5.6 X lo4 M-l cm-I); conversion to the equilibrium mixture (Co(phen)2+)occurs in a time frame (>I s) too long to be detected by pulse radiolysis. The relationship of these results to those of Waltz and Pearson and Baxendale and Fiti is discussed.
Introductioin In a paper that has become a classic in the literature of the radiation chemistry of transition metal coordination complexes, Waltz and Pearson3 reported the result of the pulse radiolyrsis of methanolic aqueous solutions containing C ~ ( b p . y ) (bpy ~ ~ + = :!,2'-bipyridine). The reaction of ea; produced a weak tramient absorption a t h >400 nm that was attributed to law-spin (t,",eg) Co(bpy):+; high-spin (ti, e:) C ~ ( b p y )was ~ ~ +the ultimate product. The rate of spin relaxation was estimated as 1 5 x lo3 s?. The transient intermediate reacted with C ~ ( b p y ) , ~(k+ = 8 X lo8 M-l s-l), with Fe(CN):-, Co(CNP:-, Cr(CN)63-,and O2 (k > lo9M-l s-l), and with Zn2+,Mn2+,Mg2+,and Ba2+(k = 106-107 ~ - s-1). 1 Baxendale and Fiti4 reexamined the system and also observed a broad, weak transient absorption in the 4001100-nm region with a pronounced maximum at -620 nm. They confirmed the rate constants of the reaction of that species with Co(bpy)l+,Oz, and Cr(CN),3- but were unable to observe any reaction with Zn2+. The effect of Mg2+and Ba2+was attributed to kinetic salt effects or to the presence of reducing impurities in the solutes. They rejected the low-spin Co(I1) assignment and suggested that the transient precurslor to the Co(I1) product in the reaction of ea; with Co(bpy)ipbe dlescribed as an electron adduct to the ligand (Co(~,py)z(bl~y-.)2+) Coordinated to the Co(1II) center. They admitted that the mechanism whereby this coordinated radical species is converted to the Co(I1) product is not clear, requiring reaction with the substrate and not ligand-to-metal intramolecular electron transfer. In a preliminary communication5 we reported that the reaction of ea; with C ~ ( b p y ) , ~produces + an intense long-lived transient absorption a t 300 nrn which was attributed to ai reduced ligand radical coordinated to the Co(1IIP center and assumed to decay via intrarnolecular electron transfer. A:, part of our continuing investigation of the interaction of radiation-generated free radicals with 0022-3654/79/2083-0439$0 1.OO/O
polypyridines and their coordination we have reexamined the reaction of Co(bpy)gB+and its analogue, Co(phen)?+ (phen = 1,lO-phenanthroline),with reducing radicals and eaq- using pulse and continuous radiolysis techniques. Experimental Section Materials. Co(bpy),(C104),.2H20 was prepared and purified according to described procedures;1° Co(phen)3(C104)3.1/2H20was provided by Professor L. Moggi from his photochemical study.ll Aqueous solutions of these complexes are stable in acidic, neutral, and alkaline media; no changes in their absorption spectra are detectable as a function of pH. Methanol and 2-propanol were Baker Analyzed reagents and were used witlhout further purification or were further purified by the method of Baxendale and Wardman." tert-Butyl alcohol (Mallinckrodt) was used as received. Sodium formate was recrystallized tvvice from water. NzO was either used directly or was passed through a column of NaOH pellets. Free bpy (Aldrich) was recrystallized three times from methanol. All other chemicals were of the highest grade available and were used without further purification. Solutions werie made up in high purity triply distilled water with the plH adjusted with NaOH, phosphate buffer, HzS04, or HC1014. Radiation Techniques. The pulse radiolysis equiprnent a t the U.S. Army Natick Research and Development Command and the C.N.R. Laboratory (Bologna) have been described in detai1.13J4 Transient absorption spectra were obtained by optical spectrophotometry with a time resolution of -0.5 ys. The radiation dose per pulse was established by the use of SCN- d 0 ~ i m e t r y . lIn ~ order to avoid or minimize photodecomposition induced by the analyzing light, solutions were exposed to as little UV Light as possible. Continuous radiolyses were conducted in @Co y sources with dose rates of 4-18 X 10l6 eV mL-l min-l 0 1979 American Chemical Society
