J. Phys. Chem. B 2000, 104, 6233-6239
6233
Spectral Properties and Redox Potentials of Silver Atoms Complexed by Chloride Ions in Aqueous Solution Isabelle Lampre, Pascal Pernot, and Mehran Mostafavi* Laboratoire de Chimie Physique, CNRS UMR 8610, UniVersite´ Paris-Sud, Centre d’Orsay, Baˆ t. 350, 91405 Orsay Cedex, France ReceiVed: February 10, 2000; In Final Form: April 10, 2000
Pulse radiolysis is used to determine the absorption spectra of the transient species formed during the decay of the hydrated electron in an aqueous solution containing AgCl43-. The mechanisms of AgCl43- reduction by solvated electrons and hydroxyalkyl radicals of 2-propanol are discussed. The redox potential of the couples AgCl43-/Ag0Cl44- and AgCl43-/Ag0 + 4Cl- is evaluated. It is shown that the redox potential of the solvated silver atom couple is lowered by the presence of Cl- ligands. The value of the E°(AgCl43-/Ag0Cl44-) is estimated to be around -2.3 VNHE.
Introduction As the early species of metal clusters produced after reduction such as atoms, dimers, and oligomers are short-lived, timeresolved observations of the reactions of these transients are carried out by the pulse radiolysis technique, coupled with optical absorption.1 The earliest2 and most complete data3 on the nucleation mechanism were obtained on silver clusters. Indeed, silver may be considered as a model system owing to the one-step reduction of the monovalent silver ions, hydrated or complexed by various ligands, and to the intense absorption bands of the transient atoms, oligomers, and final clusters. It is known that the optical absorption bands of Ag0 and Ag2+ strongly depend on the environment.4 They are red-shifted when the solvent polarity decreases. Indeed, the absorption bands in ethylenediamine and liquid ammoniac appear at longer wavelengths than in water.5 From electron spin-echo modulation analysis of Ag0 in ice or methanol glasses, it has been concluded that the absorption band is related to a charge transfer to solvent (CTTS).6 Ligands such as cyanide (CN-), ammonia (NH3), or ethylenediaminetetraacetate (Y4-), strongly interacting with the atom or the dimer, have also a great influence on the absorption spectra.7 It was shown that the transient product of the reduction of complexed silver ions is not the isolated hydrated atom but a complexed silver atom. In the presence of CN-, NH3, and Y,4- Ag0(CN)22-,8-10 Ag0(NH3)2,11 and Ag0(Y4-)9,12 are formed, respectively. Similar results were also obtained for gold13 and copper14 atoms. Oxidation reactions of Ag0 and Ag2+ by even mild oxidizing molecules were observed by pulse radiolysis.2,15 The high oxidation rate constants indicate a strong electron-donating character for the silver atom. That confirmed the evaluation of the redox potential of the silver atom couple (E°(Ag+/Ag0) ) -1.75 VNHE) which was derived from the difference between the metal electrode potential, E°(Ag+/Agm), and the metal sublimation energy.16 More generally, due to high values of the sublimation energy, the redox potential of any metal atom couple, E°(M+/M0), is expected to be quite negative. In fact, the redox potentials of other metal atom couples, for example gold13 and thallium,17 were found to be very low. * To whom correspondence should be addressed.
In previous studies, the redox potential of the silver atom complexed by Y4-, NH3, or CN- was evaluated with the help of thermodynamic cycles using optical properties of the silver atom. The value of the redox potential was found to be -2.2, -2.4, and -2.6 VNHE, respectively.8-12 To check these calculated values, radiolysis experiments were undertaken.8-12 The experimental work confirmed that the redox potential of the silver atom couple is lowered by ligands. In the present paper, we study the reduction of an unusual silver complex, AgCl43-, in solution using γ-radiolysis and pulse radiolysis techniques. The kinetics and the thermodynamics of the AgCl43- reduction are discussed, and the earliest reduced forms are characterized by their optical absorption spectrum. Experimental Section All the reagents were pure chemicals: AgClO4 from Aldrich; 2-propanol, NaCl, and acetone from Prolabo. The aqueous solutions of silver ions were prepared with an excess of Clfor complexing all silver ions. The solutions were degassed in vacuum. The γ-irradiation vessel was connected with a quartz cell having a 2 mm optical path length for spectrophotometric measurements under vacuum. The γ-irradiation source was a 60Co γ-facility of 7000 Ci with a dose rate of 1.75 kGy h-1. The steady-state absorption spectra were recorded with a single beam UV-vis spectrophotometer (Hewlett-Packard 8453). Electron pulses (3 ns duration) were delivered by a Febetron 706 accelerator (600 keV electron energy) to samples contained in a quartz suprasil cell with a thin entrance window (0.2 mm) for the beam and an optical path length of 1 cm.19 The solution was deaerated by a nitrogen flow and changed after each pulse. Absorption of transient species was analyzed by means of a classical xenon lamp, monochromator, and photomultiplier setup connected with a transient digitizer. Splitting of the beam made it possible to record simultaneously two signals at two different wavelengths. One signal for which the wavelength remained unchanged was used as a reference in order to take into account the electron pulse intensity variations. For the kinetic analysis, as bimolecular processes are generally involved, the variation of the initial electron concentration has to be carefully considered. Therefore, the raw kinetics signals at different wavelengths were first sorted out according to the
10.1021/jp000544n CCC: $19.00 © 2000 American Chemical Society Published on Web 06/10/2000
6234 J. Phys. Chem. B, Vol. 104, No. 26, 2000
Lampre et al. TABLE 2: Characteristics of the Four Types of Studied Aqueous Solutions and Measured Reduction Yield, G, of AgCl43- (Figure 3) solutiona
pH
a b c d
2.1 8.7 12.2 12.2
acetone (M)
G (10-7 mol J-1)
0.2
0.5 ( 0.1 2.6 ( 0.2 6.2 ( 0.3 6.5 ( 0.3
All the solutions contain 5 × 10-4 M AgClO4, 2.8 M NaCl, and 0.2 M 2-propanol. a
Figure 1. Percentage of the different complexes AgCln(n-1)- as a function of the Cl- concentration for a solution containing 10-3 M Ag+ according to the complexation constants Kn (Table 1). The calculations do not take into account the precipitation of AgCl (Ks ) 1.77 × 10-10 M2) which occurs for Cl- concentrations below 1.9 M.
TABLE 1: Complexation Constants, Kn, of Ag+ by Cl- To Form the Complexes AgCln(n-1)- in Aqueous Solution22 Kn
(M-n)
AgCl
AgCl2-
AgCl32-
AgCl43-
102.9
104.7
105
105.9
intensity, i.e. the time-integral of the reference signal. Then, sets of signals with close intensity values were selected and normalized against the intensity of the corresponding reference signal. The resulting signals were finally arranged into a photokinetic (wavelength × time) matrix A(λ,t) for further analysis. The data matrix can be written:
A(λ,t) ) ΣCi(t)i(λ)
Figure 2. Absorption spectrum of a solution containing 5 × 10-4 M AgClO4, 2.8 M NaCl, and 0.2 M 2-propanol at pH 8.7 (solution b) as a function of the absorbed irradiation dose (in gray): (1) 0, (2) 114, (3) 285, (4) 428, (5) 570, and (6) 855. The inset highlights the observed isosbestic point at 245 nm.
(1)
where Ci(t) is the concentration of the species i at time t and i(λ) is the molar extinction coefficient of the species i at a given wavelength λ. Recovery of the absorption spectrum of each species from the photokinetic matrix depends on the kinetic scheme. Global analysis method19 is designed to analyze the full photokinetic matrix in order to determine simultaneously the kinetic parameters (initial concentrations, reaction rate constants) and the absorption spectra. On the one hand, the global analysis method performs a nonlinear least-squares optimization of the kinetic parameters and the kinetic traces are computed with an adaptive stepsize fourth-order Runge-Kutta algorithm.20 On the other hand, the absorption spectra are determined by fitting the photokinetic matrix with a constrained linear least-squares method.21 Results In aqueous solutions containing silver ions, when the concentration of Cl- ions is very high, the precipitation of silver chloride is avoided and several complexes between Ag+ and Cl- are formed (Table 1).22 In the present study, we fix the concentration of Cl- ions at 2.8 M to ensure that all the solutions contain more than 96% of silver ions under the form AgCl43(Figure 1).23 This complex is very stable, even under illumination. It presents a specific absorption spectrum situated in the UV domain. The maximum is located at 218 nm with an extinction coefficient of 15 000 M-1 cm-1. We irradiated four types of aqueous solutions, all containing 5 × 10-4 M AgCl43and 0.2 M 2-propanol (Table 2). The pH before irradiation of the first solution (solution a) is 2.1 and that of the second one (solution b) is 8.7. The two other solutions (solutions c and d) are at pH 12.2, and only one of them contains in addition 0.2
Figure 3. Optical density at 218 nm at increasing irradiation dose for the four types of studied aqueous solutions.
M acetone (solution d). For solutions b-d, an isosbestic point is present at 245 nm on the spectra recorded as a function of the irradiation dose (Figure 2, inset), indicating that the decrease of the complex concentration observed at 218 nm is accompanied by the formation of silver particles which display a very broad absorption band in the visible domain (Figure 2). Indeed, the formed clusters are very large, and the well-known plasmon absorption band around 400 nm of silver clusters is not observed. That can be understood by the very high value of the ionic strength which accelerates the coalescence process and by the absence of any stabilizer. We observe that the spectral evolution with the irradiation dose drastically differs for the four solutions. Figure 3 presents the optical density at 218 nm as a function of the irradiation dose. At pH 2.1 (solution a) the absorbance declines slightly, whereas at pH 8.7 (solution b) the optical density decreases linearly versus the dose. At pH 12.2, in the absence (solution c) or presence (solution d) of acetone, the decrease of the optical density at 218 nm is more pronounced (Figure 3). We deduce the reduction yield, G, of AgCl43- from
Spectral Properties and Redox Potentials of AgCl43-
J. Phys. Chem. B, Vol. 104, No. 26, 2000 6235
the slope of the curves in Figure 3 (Table 2). The yield is less than 0.05 (0.5) for solution a, 0.26 (2.5) for solution b, and around 0.63 µmol J-1 (6.1 species/100 eV) for solutions c and d at pH 12.2. It is worth reminding that without complexation of silver ions, for instance in sulfate or in perchlorate medium, the reduction yield of silver ions is close to 0.62 µmol J-1 (6.0 species/100 eV). The main reactive radicals generated by irradiation are e-hyd, OH•, and H•. H• atoms are scavenged by 2-propanol as follows: 24
H• + (CH3)2CHOH f (CH3)2C•OH + H2; k2 ) 7.4 × 107 M-1 s-1 25 (2) The OH• radicals could react in acidic medium with Cl- ions which are in excess in solution: H+
OH• + Cl- 98 Cl• + OH-; k3 ) 3.0 × 109 M-1 s-1 26 (3)
Figure 4. Reciprocal optical density at 720 nm as a function of time for an aqueous solution containing 2.8 M NaCl and 0.2 M 2-propanol at pH 8.7. Inset: corresponding transient absorption signal.
Chloride atoms are complexed by Cl- very quickly:
Cl• + Cl- f Cl2•-; k4 ) 6.5 × 109 M-1 s-1 27
(4)
Cl2•- reacts with (CH3)2CHOH to form the alcohol radical:
Cl2•- + (CH3)2CHOH f (CH3)2C•OH + 2Cl- + H+; k5 ) 1.5 × 105 M-1 s-1 28 (5) In basic medium, the OH• radicals are scavenged by 2-propanol:
OH• + (CH3)2CHOH f (CH3)2C•OH + H2O; k6 ) 1.9 × 109 M-1 s-1 25 (6) Under our experimental conditions, the reactions 2-6 are fast. Moreover, the alcohol radical exists under two forms depending on the pH (pKa ) 12):29
(CH3)2C•OH a (CH3)2C•O- + H+
(7)
At pH 2.1, the solvated electrons are scavenged by H+ to form H• atoms. Therefore, at this pH and in the presence of 2-propanol, (CH3)2C•OH is the only reducing species formed under irradiation. But at pH 2.1 (solution a), while the redox potential of the alcohol radical couple (CH3)2CO + H+/ (CH3)2C•OH is -1.4 VNHE, the reduction of the silver complex is almost suppressed. Only at very high doses are a few silver clusters formed. At pH 8.7, the reducing species are e-hyd and (CH3)2C•OH formed with a radiolytic yield of 0.28 and 0.34 µmol J-1, respectively. In this case, the reduction yield is found to be 0.26 µmol J-1 and not 0.62 µmol J-1 (solution b). At pH 12.2, the dominated form of the alcohol radical is the deprotonated radical, (CH3)2C•O-. Moreover, at pH 12.2 and in the presence of acetone (solution d), the hydrated electrons are scavenged according to the following reaction:
e-hyd + (CH3)2CO f (CH3)2C•O-; k8 ) 6.5 × 109 M-1 s-1 25 (8) This reaction is very fast; therefore, at pH 12.2 and in the presence of acetone, the reducing species formed under irradiation with G ) 0.62 µmol J-1 is the alcohol radical. At pH 12.2, in the presence or absence of acetone (solutions c and d) we obtain a reduction yield close to 0.63 µmol J-1. It is worth noting
Figure 5. Time profile of transient absorption signals monitored at different wavelengths upon pulse radiolysis of solution b.
that, in the basic medium, the redox potential of the couple (CH3)2CO/(CH3)2C•O- is -2.1 VNHE. To obtain the spectrum of the first reduced form of AgCl43and to check the role of the deprotonated form of the alcohol radical in the reduction mechanism, we carried out a pulse radiolysis study. First, we observed at 720 nm the decay of the solvated electron in the absence of silver complex for aqueous solution at pH 8.7 containing 2.8 M NaCl and 0.2 M 2-propanol (Figure 4, inset). The decay of the hydrated electron follows a second-order kinetics law from which we deduce a rate constant k ) 2.45 × 1010 M-1 s-1 (with 720 ) 18 200 M-1 cm-1). The decay is mostly due to the following reaction:
e-hyd + e-hyd f 2OH- + H2
(9)
But, we cannot exclude possible reactions with H2O2 or impurities, for instance from NaCl salt, in solution. The rate constant that we found under our experimental conditions is 5 times as high as that known at low ionic strength for reaction 9.30 In the presence of the silver complex in solution b, at pH 8.7, the decay of the solvated electron observed at 600 nm is correlated to a rise in the optical density observed in the UV domain (Figure 5). We followed the decay of the solvated electron versus the AgCl43- concentration which is much higher than the hydrated electron concentration produced by the pulse. In this case, the decay follows a pseudo-first-order law with an observed rate constant, kobs. From the linear dependence of kobs versus the AgCl43- concentration, the rate constant of the
6236 J. Phys. Chem. B, Vol. 104, No. 26, 2000
Lampre et al.
Figure 6. First-order rate constant of the optical change observed at 720 nm, kobs, as a function of the Ag+ concentration for solutions containing 2.8 M NaCl and 0.2 M 2-propanol at pH 8.7.
AgCl43- reduction reaction by the solvated electron is found to be 7.8 × 109 M-1 s-1 (Figure 6). The obtained value is close to that of a reaction controlled by diffusion. It is worth noting that the rate constant is determined at high ionic strength (2.8 M) where the charge effect is weakened. While the rise in transient optical density recorded around 350 nm is correlated to the decay of the solvated electron and reaches its maximum around 350 ns, the optical density at 310 nm reaches its maximum later, around 750 ns (Figure 5). Therefore, we present the absorption spectra recorded at 350 and 750 ns after the pulse for solutions b and d (Figure 7). For solution b, soon after the pulse, the spectrum displays a wide absorption band with a maximum at 360 nm and a shoulder around 430 nm and a less intense absorption band above 500 nm due to the hydrated electron (Figure 7, top). At a longer time (750 ns) the shape of the spectrum is different from that obtained at a shorter time. The spectrum is shifted to the blue and exhibits a band around 330 nm. For solution d at pH 12.2 and in the presence of acetone, both absorption spectra recorded after the pulse are similar and display a band peaking at 330 nm (Figure 7, bottom). The shapes of these spectra are close to that obtained at 750 ns for solution b without acetone (Figure 7). Note that at pH 12.2 the alcohol radical weakly absorbs in the range of 200-400 nm, too.31 Moreover, like γ-radiolysis experiments, the pulse radiolysis study shows that for solution d the reduction yield of silver is roughly twice as high as that obtained for solution b. Indeed, we observed the evolution of the optical density at 400 nm in the time range of seconds when the growing silver clusters absorb. The signal for solution d (when acetone is present in the solution) is twice as intense as that observed for solution b (Figure 8). This result indicates that all the radicals (CH3)2C•Oparticipate in the silver reduction. Discussion The time evolution of the spectra observed for solution b until 800 ns (Figures 5 and 7) is fitted with a global analysis method taking into account the presence of three absorbent species in the range of 300-600 nm: the solvated electron, the AgCl43reduction product, and a charged dimer (Ag2+,Cl-) issued from the association of the silver atom with AgCl43- in excess. As was reported in the Results section, in the case of solution b, the reduction yield is 0.26 µmol J-1. Thus, we consider that only the solvated electron participates in the reduction of
Figure 7. Transient absorption spectra recorded at 350 and 750 ns after the pulse for solutions b and d.
Figure 8. Time profile of transient absorption signals monitored at 400 nm upon pulse radiolysis of solutions b and d.
AgCl43-. The reactions 2-7 are fast and occur within the pulse duration. Therefore, in the kinetics simulations, three major reactions are considered: reaction 9, AgCl43- reduction, and Ag2+,Cl- formation. For the first two reactions, we use the experimental rate constants obtained above (2.45 × 1010 and 7.8 × 109 M-1 s-1, respectively). Moreover, the well-known absorption spectrum of the solvated electron is imposed. So, using the kinetic analysis of the experimental transient spectra, we determine the rate constant of Ag2+,Cl- formation (1.1 × 1010 M-1 s-1), the absorption spectrum of the AgCl43- reduction product, and that of Ag2+ complexed by Cl- (Figure 9). The absorption spectrum of the AgCl43- reduction product presents a broad band peaking at 360 nm with a slight shoulder
Spectral Properties and Redox Potentials of AgCl43-
J. Phys. Chem. B, Vol. 104, No. 26, 2000 6237 TABLE 3: Calculated Values of the Redox Potential of the Different Couples (AgICln(n-1)-/Ag0 + nCl-) According to the Thermodynamic Cycle I
E° (VNHE)
AgICl/Ag0 + Cl-
AgICl2-/Ag0 + 2Cl-
AgICl32-/Ag0 + 3Cl-
AgICl43-/Ag0 + 4Cl-
-1.97
-2.08
-2.10
-2.15
the basic form of the radical alcohol (E°((CH3)2CO/(CH3)2C•O-) ) -2.1 VNHE). Now, if we assume that the couple AgCl43-/Ag0Cl44- is also involved in the reduction reaction, we need to know the value of its redox potential. An estimation of the redox potential can be done within the framework of the continuum model using the absorption properties of Ag0Cl44-. Such an estimation method was already used in the case of CN-, NH3, and EDTA ligands.8-12 First, according to the continuum model,34 as for the halides,35 we may consider the thermodynamic cycle II:
Figure 9. Absorption spectrum of the reduction product of AgCl43in solution b (b, top), of the free hydrated silver atom (O, top), of Ag2+,Cl- in solution b (b, bottom), and of the free hydrated Ag2+ (O, bottom). The spectra of the free hydrated species are taken from ref 32.
around 430 nm (Figure 9, top). That spectrum is different from the known absorption spectrum of the free hydrated silver atom which displays a less broad band peaking at 360 nm32 (Figure 9). We consider the absorption spectrum of the AgCl43reduction product as an overlap of two bands: one due to the free silver atoms peaking around 360 nm and another one corresponding to the complexed silver atoms peaking around 430 nm. Indeed, as was previously reported, in the presence of NH3, CN-, and Y4-, the formation of complexed silver atom induces a red shift of the absorption spectrum of the silver atom. Therefore, let us consider two different possible couples for the reduction of the silver complex: AgCl43-/Ag0 + 4Cl- and AgCl43-/Ag0Cl44- involved in the following reactions:33
e-hyd + AgCl43- f Ag0 + 4Cl-
(10)
e-hyd + AgCl43- f Ag0Cl44-
(11)
As a rule, the redox potential of the AgCln(n-1)-/Ag0 + nClcouples in which the silver atom would not be complexed by Cl- can be calculated from the following cycle:
including the CTTS transitions of complexed (430 nm) and uncomplexed (360 nm) silver atoms and taking into account the two complexation constants of Ag+ and Ag0 by Cl-, K4, and K4′, respectively. This cycle gives us the relation:
hνAg0Cl44- - hνAg0 ) 0.059 (log K4′ - log K4)
(12)
Hence the value of K4′ is found to be 10-3.6. This very low value means that the stability of the complexed atom by Cl- is weak. Second, if we consider the thermodynamic cycle III:
including the redox potentials E° of the couples Ag+/Ag0 and AgCl43-/Ag0Cl44- and the two complexation constants K4 and K4′, we have:
|e|E°(AgCl43-/Ag0Cl44-) - |e|E°(Ag+/Ag0) ) 0.059 (log K4′ - log K4) (13) As the silver ions are more strongly complexed by Cl- ligands than the silver atoms are (K4 . K4′), the redox potential of the monomeric couple is lowered by complexation (eq 13). Taking E°(Ag+/Ag0) ) -1.75 VNHE, we find:
E°(AgCl43-/Ag0Cl44-) ) -2.35 VNHE Note that the redox potential depends on the complexation constant Kn (Table 3). The value of E°(AgCl43-/Ag0 + 4Cl-) is -2.15 VNHE. This value is higher than that of the solvated electron (E°(H2O/ehyd-) ) -2.87 VNHE), much lower than that of the acidic form of the alcohol radical (E°((CH3)2CO + H+/ (CH3)2C•OH) ) -1.4 VNHE), and slightly lower than that of
(14)
This value is much higher than the redox potential of the hydrated electron, which is -2.87 VNHE, but it is lower than the redox potential of the couple (CH3)2CO/(CH3)2C•O-. Given the values of the reduction yield obtained for the solutions at pH 2.1 and 8.7 (close to 0 and 0.26 µmol J-1, respectively), we can assume that the acidic form of the alcohol
6238 J. Phys. Chem. B, Vol. 104, No. 26, 2000 radical cannot reduce the complexed ions AgCl43-. That means that the redox potential of the silver complex is lower than that of the couple (CH3)2CO + H+/(CH3)2C•OH. This result is in agreement with the above estimation of the redox potentials. For the solutions at pH 12.2, the reduction yields measured under γ-irradiation (Table 2) and pulse radiolysis (Figure 8) indicate that the deprotonated form of the alcohol radical participates in the reduction of the silver complexes. Nevertheless, we do not observe immediately after the pulse the same absorption band in the presence of acetone at pH 12.2 as that formed at pH 8.7 (Figure 7). Thus we can exclude the mechanism which implies a direct electron transfer from (CH3)2C•O- to AgCl43- leading to the complexed silver atom, Ag0Cl44-. To explain the formation of silver clusters by the radical (CH3)2C•O- that we observed by pulse radiolysis, we propose another mechanism. As a first step a complexation between the silver complex and the basic form of the hydroxyalkyl radical is established:
(CH3)2C•O- + AgCl43- a (Cl-xAgI(CH3)2C•O)x- + (4-x)Cl- (15) Then an electron-transfer reaction involving a second silver complex occurs:
(Cl-xAgI(CH3)2C•O)x- + AgCl43- f (Ag2+,Cl-) + (CH3)2CO + yCl- (16) The overall reaction can be written as follows:
(CH3)2C•O- + 2AgCl43- f (Ag2+,Cl-) + (CH3)2CO + zCl- (17) Indeed, the complexation of certain metal cations by hydroxyalkyl radicals was already established: for example the reactions of hydroxyalkyl radicals with the free hydrated copper cations, Cu2+ and Cu+, were widely studied,37 and the reaction of alcohol radicals with CuI complex, CuIY4-, was also reported.37 Moreover, for the free hydrated cation Ag+ 16 and for silver ions complexed by NH3 and CN-,10,11 a reduction mechanism through the complexation by the alcohol radical, similar to reactions 15-17, was already proposed. In the presence of acetone and at pH 12.2, the spectrum found after the pulse (Figure 7, bottom) is different from that of Ag0 and that of Ag0Cl44-, but it is very close to that found for Ag2+ complexed by Cl- (Figure 9, bottom). Therefore we can attribute the spectrum found at 750 ns for solution d to that of Ag2+ complexed by Cl- which is formed through the reactions 1517. That implies that the product of reaction 15, (Cl-xAgI(CH3)2C•O)x-, either absorbs out of the observed range or disappears very quickly by reaction 16. The redox couples involved in the silver complex reduction through the reactions 15-17 are 2AgCl43-/(Ag2+,Cl-) and (CH3)2CO/(CH3)2C•O-. According to another thermodynamic cycle, we can write:
E°((2AgCl43-)/(Ag2+,Cl-)) ) E°(AgCl43-/Ag0Cl44-) + ∆G°/e (18) where ∆G° is the free enthalpy for the dissociation of Ag2+,Clinto Ag0Cl44- and AgCl43-. Since this last term is positive, we can deduce:
Lampre et al.
E°((2AgCl43-)/(Ag2+,Cl-)) > E°(AgCl43-/Ag0Cl44-) (19) Then considering the values of the redox potentials, reaction 17 is thermodynamically more favorable than the direct reduction of AgCl43- by (CH3)2C•O-. It is worth noting that there is some uncertainty about the estimated values of the redox potentials due to the uncertainty about the complexation constants reported in the literature and the use of the continuum model. Consequently, the direct reduction of AgCl43- by (CH3)2C•O- cannot be ruled out only on the basis of the calculated redox potentials but mainly on the shape of the absorption spectrum of the reduction product as we mentioned above. Conclusion The reduction of AgCl43- by the solvated electron yields Ag0 and Ag0Cl44- species. The Ag0Cl44- species is autodissociative and absorbs around 430 nm. The redox potential of the couples AgCl43-/Ag0Cl44- and AgCl43-/Ag0 + 4Cl- are estimated to be -2.30 and -2.15 VNHE, respectively. While the acidic form of the alcohol radical cannot reduce AgCl43-, the basic form, (CH3)2C•O-, first complexes AgCl43- and then transfers an electron through a reaction with another AgCl43- to form Ag2+ complexed by Cl-. The present study provides further evidence about the influence of ligands on the redox potential and the absorption spectrum of silver atoms. References and Notes (1) Henglein, A. Chem. ReV. 1989, 89, 1861. Belloni, J.; Amblard, J.; Marignier, J. L.; Mostafavi, M. In Clusters of Atoms and Molecules; Haberland, H., Ed.; Springer-Verlag: New York, 1994; p 290. Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1995, 99, 903. Belloni, J.; Mostafavi, M.; Remita, H.; Marignier J. L.; Delcourt, M. O. New J. Chem. 1998, 22, 1239. Belloni, J.; Mostafavi, M. In Metal Clusters in Chemistry; Braunstein, P., Oro, R., Raithby, J., Eds.; Wiley: New York, 1999; p 1213. (2) Baxendale, J. H.; Fielden, E. M.; Keene, J. P.; Ebert, M. In Pulse Radiolysis; Keene, J. P., Swallow, A., Baxendale, J. H., Eds.; Academic Press: London, 1965; p 207. Von Pukies, J.; Roebke, W.; Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1968, 72, 842. Henglein, A. Ber. BunsenGes. Phys. Chem. 1977, 81, 556. (3) Mostafavi, M.; Marignier, J. L.; Amblard, J.; Belloni, J. Radiat. Phys. Chem. 1989, 34, 605. Kapoor, S.; Lawless, D.; Kennepohl, P.; Meisel, D.; Serpone, N. Langmuir 1994, 10, 3018. Janata, E.; Lilie, J.; Martin, M. Radiat. Phys. Chem. 1994, 43, 353. Janata, E. Radiat. Phys. Chem. 1994, 44, 449. (4) Belloni, J.; Khatouri, J.; Mostafavi, M.; Amblard, J. In Ultrafast reaction dynamics and solVent effects; Rossky, P. J., Gauduel, Y., Eds.; American Institute of Physics: Woodbury, NY, 1993; p 541. (5) Belloni, J.; Delcourt, M. O.; Marignier, J. L.; Amblard, J. In Radiation Chemistry; Hedwig, P. L., Schiller, R., Eds.; Akad. Kiado: Budapest, 1987; p 89. (6) Kevan, L. J. Phys. Chem. 1981, 85, 1828. (7) Mostafavi, M.; Belloni, J. Recent Res. DeVel. Phys. Chem. 1997, 1, 459. (8) Remita, S.; Archirel, P.; Mostafavi, M. J. Phys. Chem. 1995, 99, 13198. (9) Remita, S.; Mostafavi, M.; Delcourt, M. O. J. Phys. Chem. 1996, 100, 10187. (10) Texier, I.; Mostafavi, M. Radiat. Phys. Chem. 1997, 49, 459. (11) Texier, I.; Remita, S.; Archirel, P.; Mostafavi, M. J. Phys. Chem. 1996, 100, 12472. (12) Mostafavi, M.; Remita, S.; Delcourt, M. O.; Belloni, J. J. Chim. Phys. 1996, 93, 1828. (13) Mosseri, S.; Henglein, A.; Janata, E. J. Phys. Chem. 1989, 93, 6791. (14) Ershov, B. G.; Sukhov, N. L. Radiat. Phys. Chem. 1990, 36, 93. (15) Tausch-Treml, R.; Henglein, A.; Lilie, J. Ber. Bunsen-Ges. Phys. Chem. 1978, 82, 1335. (16) Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1977, 81, 556. (17) Rao, P. S.; Hayon, E. J. Phys. Chem. 1975, 79, 865. Butler, J.; Henglein, A. Radiat. Phys. Chem. 1980, 15, 603. (18) Belloni, J.; Billiau, F.; Cordier, P. Delaire, J.; Delcourt, M. O. J. Phys. Chem. 1978, 82, 532. (19) Beechem, J. M.; Ameloot, M.; Brand, L. Chem. Phys. Lett. 1985, 120, 466.
Spectral Properties and Redox Potentials of AgCl43(20) Press, W. H.; Teukolsky, S. A.; Vetterling, W. T.; Flannery, B. P. Numerical Recipes; Cambridge University Press: Cambridge, 1992. (21) Hanson, R. J. Linear least squares with bounds and linear constraints. Report SAND82-1517, Sandia Laboratories, August 1982. (22) Ringbom A. Les complexes en chimie analytique; Dunod: Paris, 1967; p 293. (23) Several complexes between Ag+ and Cl- are reported in the literature as well as different complexation constants. We used the values of the complexation constants given in ref 22. In a few papers the formation of AgCl43- is not taken into account (see, for example, Kim, J. I.; Duschner, H. J. Inorg. Nucl. Chem. 1977, 39, 471), whereas in others the presence of AgCl54- at high Cl- concentration is also admitted (see, for example, Demakhin, A. G.; Zimakov, I. E.; Sinitsin, V. I. Zh. Neorg. Khim. 1974, 29, 245). (24) Baxendale, J. H.; Busi, F. The Study of fast Processes and transient Species by Electron Pulse Radiolysis. NATO ASI Series 86, D. Reidel, 1982. (25) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513. (26) Grigo’rev, A. E.; Makarov, I. E.; Pikaev, A. K. High Energy Chem. 1987, 21, 99. (27) Klaning, U. K.; Wolff. T. Ber. Bunsen-Ges. Phys. Chem. 1985, 89, 243.
J. Phys. Chem. B, Vol. 104, No. 26, 2000 6239 (28) Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1982, 86, 241. (29) Asmus, K. D.; Henglein, A.; Wigger, A.; Beck, G. Ber. BunsenGes. Phys. Chem. 1966, 70, 756. Schwarz, H. A.; Dodson, R. W. J. Phys. Chem. 1989, 93, 409. (30) Schmidt, K. H.; Bartels, D. M. Chem. Phys. 1995, 190, 145. (31) Simic, M.; Neta, P.; Hayon, E. J. Phys. Chem. 1969, 73, 3794. (32) Ershov, B. G.; Janata, E.; Henglein, A.; Fojtik, A. J. Phys. Chem. 1993, 97, 4589. (33) To simplify the discussion, in the present work, we do not take into account possible silver complex reduction reactions combined with Cl- ligand detachment such as e-hyd + AgCl43- f Ag0Cl33- + Cl-. (34) Treinin, A. J. Phys. Chem. 1964, 68, 893. (35) Blandamer, M. J.; Fox, M. F. Chem. ReV. 1970, 59, 9. Grossweimer, L. I.; Matheson, M. S. J. Phys. Chem. 1957, 61, 1089. (36) Buxton, G. V.; Green, J. C. J. Chem. Soc., Faraday Trans. 1987, 74, 697. Freiberg, M.; Mulac, W. A.; Schmidt, K. H.; Meyerstein, D. J. Chem. Soc., Faraday Trans. 1980, 76, 1838. Cohen, H.; Meyerstein, D. J. Chem. Soc., Faraday Trans. 1988, 84, 4157. (37) Buitenhuis, R.; Bakker, C. N. M.; Stock, F. R.; Louwrier, P. W. F. Radiat. Phys. Chem. 1980, 16, 5.
