The Journal of Physical Chemistry, Vol. 83, No. 3, 1979 423
Communications to the Editor
COMMUNICATIONS TO THE EDITOR Kinetics of the Reaction of Solvated Electrons with Benzene in Water-Ammonia Mixtures
Sir: In a reclent communication Gordon and co-workers1 give proven evidence that the reaction of solvated electrons e,- with benzene in water can be described by the sequence
+ HzO
k2
C,H7*
OH-
(2)
where the protonating step (eq 2) occurs so rapidly that the benzene radical anion C6H6-cannot be observed. In aprotic solvents on the other hand the formation of the radical anion is an equilibrium r e a c t i ~ n ~ ! ~ e,-
kl
+ C6H6 k_l C&6-
1 0
I
(la)
the equilibrium constant of which K = kl/k-l is strongly solvent dependent. Assuming the same equilibrium in water and applying the concept of the quasi-stationary state for the C6H6-radical anion (d[C6H6-]/dt= 0)gives for the observed second-order rate constant of the disappearance of the electrons (cw is the water concentration) k1kzcw - h1kzcw/k-, (3) kobsd = k-1 + kzcw 1 + hzcw/h-l Therefore only in case of hzcw >> k-l the observed rate constant is identical with the rate constant hl of electron attachment to the substrate. Furthermore the reaction rate should slow down to zero in solvent mixtures with increasing content of solvents with which C6H6-cannot be protonated. This indeed we find in water-ammonia mixtures in which the solvated electrons were produced by laser-flash photolysis (10 * s, 100 mJ, 265 nm) of dissolved KI M) with CH30K (or KOH + methanol, M) as a very effective scavenger for reactive photolysis byproducts (electron lifetime in substrate free ammonia > 1 min). The observed rate constants (25 "C, normalized to the reported value for water) are displayed in Figure 1,showing their decrease by more than 7 powers of ten with decreasing water content. (No reaction was observed in pure ammonia. Also with the ESR technique we cannot observe the C6H6-radical anion in a metal-ammonia solution with added benzene, proving that the equilibrium constant of reaction l a in liquid ammonia is K < A few mole percent of water in the metal-ammonia solutions however leads to hydrogenated benzene (Birch r e a ~ t i o n via ) ~ reaction 2 which disturbs equilibrium la). The temperature dependence of the observed reaction rate corresponds to an activation energy of 4 f 1 kcal mol-' independent of solvent composiition. The experimental points of Figure 1 cannot be fitted by the rate constaiit (eq 3) with solvent independent h l and k2/k-1. However, considering that the entropy of the electrons Se; is known to be more positive (about 10 cal mol-' K-1)5and that of anions S A - is more negative (about -20 cal mol-' K-1)6in ammonia than in water the entropy change associated with reaction 1 (4) hs" = SCsHs- - SCsHG- se; 0022-3654/79/2083-0423$01 .OO/O
P
-2
.2
0
.6
.4
.8
Flgure 1. Observed second-order rate constant for the disappearance of solvated electrons by reaction with benzene in water-ammonia mixtures at 25 "C ( x IS the mole fraction of water). The full line was calculated according to eq 9.
should be solvent dependent. According to the transition state theory the rate constant of a reaction depends on the activation entropy AS*
k,
-
exp(AS*,/R)
(5)
(R is the gas constant) and the reaction entropy is given by the difference of activation entropies of the forward and back reaction
AS' = AS'& - AS'?,
(6)
Therefore the rate constants hl and kl should be solvent d e ~ e n d e n t . ~In the following we make the simplified assumption that the reaction entropy of (1)in the mixture AS," and also the activation entropies of both the forward and back reaction AS*,,, and vary linearily with mole fraction of water x
AS," = ASA" AS*,,, = AS*,,*
+ ( S + S)X
+ SX
AS*-l,, = AS*-1,A
(7)
+ sx
(8)
(AS*", are the reaction entropy and activation entropy of the forward and back reaction, respectively, in pure ammonia ( x = o).) So the reaction entropy contributes to both activation entropies depending on the choice of parameter s. The rate constant (eq 3) then can be written ~ I , exp((s A kobsd
+ S)x/R)hzcw/(k-l,A
exp(sx/R)) (9) hZcW/(k-l,A exP(sx/R))
= -~
-
1 4-
-
with kl,* exp(ASI1,A/R)and k-lA exp(AS*_l,A/R), rate constants in pure ammonia. The solid line of Figure 1 is the best fit obtained with S = 30 and s = -24 cal mol-' K-l (Le., the solvent dependence of the reaction entropy is mainly given by that of the activation entropy of the back reaction) and kl,* = 5 x l o 5 M-l s-' a nd k2/k_,,A= 3 x M-'. For water ( x
0 1979 American Chemical Society
424
The Journal of Physical Chemistry, Vol. 83, No. 3, 1979
= 1)it follows kl,w = 1 X lo7M-l s-I and k2/k-l,w = 50 M-l. The protonating reaction (eq 2) as a radical reaction might be diffusion controlled with the rate constant k2 of around 1 X 1O'O M-' s-l, independent of solvent. Then k l win water and k I , A in ammonia would be around 2 X lo6 and 3X s-l and the equilibrium (K = kl/k_l) constant of reaction 1 around Kw = 5 X M-' and KA = 1.7 X in the two respective pure solvents, a11 at 25 "C. Although the treatment might seem speculative, the results do justify it. 1. In water the overall rate is given by the rate of the electron attachment (k,,w = hobvawith k 2 >> k-l,w) in accordance with the assumption of Gordon and co-workers.' 2. The difference of the reaction entropy according to (7) in water and in ammonia (S = 30 cal mol-' K-l) agrees with the value expected on the basis of the differences of entropies of the solvated electrons and the anions in the two solvents, as given above. 3. In ammonia the equilibrium constant KAof reaction 1 is in agreement with the ESR experiment, The absence of any reaction in ammonia thus is a consequence of the unfavorable equilibrium (eq 1)in conjunction with the goor proton-donating property of ammonia, which does not allow reaction 2 to procede. 4.With AGO = -RT In K it follows that the free energy change of the electron attachment reaction 1with benzene is positive both in water and in ammonia (AGw"= 1.8 and AGA" = 10.6 kcal mol-l, 25 "C) in accordance with the negative electron affinity of benzene8 and in contrast to all higher aromatic moleculesgJOwhich react very fast with solvated electrons and for which K >> 1, or AGO < 0 or electron affinity > 0 (e.g., biphenyl in ammonia:'l K = 8 X lo4 M-l, AGO == -7 kcal mol-l, AHo = -20 kcal mol-l, and AS" N -45 cal mo1-l K-l, kl N 8 X 1O'O M-l s?, k-l 1 x 106 s-1). 5. From AG," - AGW" = 8.8 kcal mol-I and ASAO - ASw" = -30 cal mol-l K-I (eq 7) follows AHAo - AHwo = 4.1 kcal mol-', i.e., the reaction enthalpy of (1) is almost independent of the solvent composition. This is to be expected since AHAo- AHw" is mainly given by the difference of the solvation enthalpies of the electron in water and in ammonia, both of which are about12 -40 kcal mol-], the difference of solvation enthalpies of the benzene and its anion cancelling, because of little interaction with the solvent due to zero or small charge density of both species. Also the observed solvent independent activation energy of the overall reaction supports a small difference of AHA' - AHw",if reaction 2 has the solvent independent activation energy of a diffusion-controlled reaction. Thus it seems that the given rate constants k l , kl,and k2 for the e--CsHs reaction sequence in ammonia and in water can explain the observed reaction kinetics and also the different behavior of benzene compared to that of the higher aromatic molecules with respect to the formation of their radical anions.
References and Notes (1) S. Gordon, K. H. Schmidt, and E. J. Hart, J . Phys. Chem., 81, 104 (1977). (2) T. R. Tuttle and S. I. Weissman, J. Am. Chem. Soc., 80, 5342 (1958). (3) K. W. Boddeker, G. Lang, and U. Schindewolf, Angew. Chem., Int. Edit. Engl., 8 , 138 (1969). (4) H. Smith, "Organic Reactions in Liquid Ammonia", Vieweg and Sohn, Braunschweig; Interscience, New York, 1963. (5) J. Jortner and R. M. Noyes, J . Phys. Chem., 70, 770 (1966); G. Lepoutre and J. Jortner, bid., 76, 683 (1972). (6) C. M. Criss, R. P. Held, and E. Liksha, J . Phys. Chem., 72, 2970 (1968). ( 7 ) P. Krebs and E. San Roman, Chem. Phys. Lett., in press. (8) G. Briegleb, Angew. Chem., 76, 326 (1964). (9) B. Bockrath and L. M. Dorfman, J. Phys. Chem., 77, 1002 (1973). (10) R. A. Holroyd, Ber. Bunsenges. Phys. Chem., 81, 298 (1977).
0022-3654/79/2083-0424$01.00/0
Communications to the Editor (11) Farhataziz and L. M. Perkey, J. Phys. Chern., 80, 122 (1976). (12) N. R. Kestner in "Electron-Solvent and Anion-Solvent Interactions", L. Kevan and B. C. Webster, Ed., Elsevier, Amsterdam, 1976. Institut fur Physikalische Chemie und Elektrochemie Universitit Karlsruhe D-7500 Karlsruhe, West Germany
U. Schlndewolf * B. Neumann
Received January 10, 1978; Revised Manuscript Received December 5, 1978
Proton Phototautomerism in a Metal Complex. Determination of pK," for a Monodentate 2,2'-Bipyridine Complex of Iridium(II1) Publication costs assisted by the Department of Energy
Sir: We report here some tentative evidence for the first observation of proton phototautomerism in a transition metal complex. Several well-documented examples of this phenomenon have been found for organic molecules such as salicylic a ~ i d l and - ~ 7 - a ~ a i n d o l e ;single ~ - ~ proton phototautomerism is displayed by the former, while twoproton phototautomerism is evident in the dimer of the latter, which has been presented as a model for the study of electronic interactions in DNA base pairs.4 Photoinduced charge-transfer phenomena, which are common among transition metal complexes, should lead to a wide variety of phototautomerizations in this class of compounds if ionizable protons and two or more basic sites are available. The complex ion, Ir(bpy)zHaO(bpy)3+, whose structure and ground state acid dissociation are depicted in I and 11, represents an example of one such c ~ m p o u n d .Recent ~ 2+
I
I1
success in the use of I8to sensitize the photoisomerization of norbornadieneg with high efficiency has led us to undertake a study of the quenching of its luminescence by a variety of acceptors. During the course of this study it was necessary to perform a luminescence titration of I to establish its pKa*. The results of this pKa* determination along with previous luminescence data suggest t,hat single proton phototautomerization occurs in I. Previous measurement7 of the ground state pK, of I yielded a value of 3.0 f 0.1. Due to the long luminescence lifetimes of I and I1 in aqueous room temperature solutions (12.2 and 10.0 ks, respectively), and the large bimolecular rate constants normally associated with intermolecular proton transfer processes (ionization),lo acid-base equilibrium is likely to be established within the lifetime of the luminescence states. We were unable to demonstrate intermolecular proton transfer within the excited state lifetime directly, as has been done" in the case of rutheniumbis(2,2'-bipyridine) (2,2'-bipyridine-4,4'-dicarboxylic acid)2+,due to the much smaller difference between pK, and pK,* in our case. However, the sharp inflection region in the luminescence titration curve (Figure 1) provides indirect evidence for establishment of the ionization equilibrium within the excited state lifetime.3 0 1979 American Chemical Society
