J. Phys. Chem. 1994,98, 10891-10894
10891
Mixed-Metal Clusters in Aqueous Solution: Reactions of T12+ with Ag+, Cd2+,and Pb2+ and of Zn+ with TI+ B. G. Ershov,f E. Janata, and A. Henglein; Hahn-Meimer Institut, Abteilung Kleinteilchenforschung, 14109 Berlin, FRG Received: May 16, 1994@
Tlz+ is generated pulse radiolytically in 2 x M T1+ solutions containing small concentrations of Ag+, Cdz+, and Pb2+ ions. The reaction of Tlz' with Ag+ occurs with k = 3 x IO9 M-' s-', the first product being a mixed cluster to which the structure (TlAg)+ is attributed. It has strong absorption bands at 375 and 270 nm. The cluster rapidly reacts in a complex process with a second Ag+ ion (k = 5 x lo9 M-' s-'). T12+ reacts with Cdz+ with k = 9 x IO7 M-' s-l to yield a product that absorbs below 250 nm. The reaction is shown to be not a simple electron transfer; formation of a cluster, possibly (CdT1)2+, is proposed. Reaction of T l 2 + with Pbz+ consists of the transfer of an electron, k = 2 x IO9 M-' s-l. Zn+ transfers an electron to T1+ with k = 5 x IO9 M-' s-'. These reactions are discussed in terms of the electrochemical potentials of the abnormal valence states of the metal ions involved.
Introduction The method of pulse radiolysis has recently been applied to detect very small charged silver clusters such as (CdAg)2+and (CoAg)z+.l The clusters are formed in the reactions of the ions in abnormal valence states, Cd+ and Co+, with the Ag+ ion. The detection of the clusters has become possible through the development of very sensitive pulse radiolysis equipment, which allows one to work with concentrations in the order of M of the monovalent metal ions; at these small concentrations, second-order reactions between the short-lived intermediates are avoided to a great extent. Moreover, as the clusters have absorptions which often strongly overlap with those of the initiating species, Cd+ and Co+, sophisticated computer simulation had to be developed to determine rate constants and absorption spectra. In the present paper, the formation, decay, and optical absorption of a thallium-silver cluster is described. It is knownZ that the first products of the radiolytic reduction of T1+ are the free atom and the cluster T12+, formed in the processes eaq-
T1°
+ T1+ - T1°
+ T1+
T1,+
-
Ag+
+ e-
Cd+
(2)
(3)
is -1.8 V,3 Le., not significantly different from the thallium potentials. Electron transfer between the reduced thallium Home address: Institute of Physical Chemistry, Russian Academy of Sciences, Leninskii 31, 117915 Moscow, Russia. * To whom correspondence should be addressed. EI Abstract published in Advance ACS Abstracts, September 15, 1994.
0022-365419412098-10891$04.50/0
-
+ eaq-
Cdz+
(4)
is close to -1.8 V? Electron transfer from the fiist reduction products of thallium to CdZ+would have little driving force; however, as in the thallium-silver system, mixed-cluster formation could be expected. The standard potential of the system Pb+
-
Pbz+
+ e-
(5)
is -1.0 V.5 Electron transfer from reduced thallium to lead should be strongly exoergic. The standard potential of the redox system Zn+
Values between 2.3 x IO3 and 1.4 x IOz M-' have been reported for the constant of the equilibrium of eq 2,zb*cfrom which one calculates a small association free enthalpy of -0.17 to -0.13 eV. Both T1° and Tlz+ are strong reducing agents, the respective electrochemical potentials of the redox systems T1+/T1° (where T1° is afree atom in solution) and 2Tl+Dl2+ being - 1.9 and - 1.7 V.Zb,cThe standard potential of the system Ago
species and Ag+ should therefore have little driving force. As already pointed out,' the formation of mixed-metal clusters may be favored under these conditions. Experiments were also performed to detect electron transfer or cluster formation in the reactions of the abnormal valence state of thallium with cadmium, lead, and zinc. The standard potential of the redox system
-
Zn2+
+ eaq-
(6)
is as negative as -2.5 V.4 In this case, electron transfer from Zn+ to T1+ seemed probable.
Experimental Section The pulse radiolysis was carried out with 3.8 MeV electrons. Pulses of 0.1 ps or 5 ns duration could be generated. The equipment has been described in detaiL6 The signals were averaged over 10 pulses. The absorbance changes are presented as E values (absorbance change divided by the concentration of hydrated electrons generated in one pulse and the thickness of the optical cell, 1.5 or 0.5 cm). The absorbed dose per pulse is given in terms of the concentration of hydrated electrons generated. In the computer simulations, the differential equations for the elementary reactions were numerically solved. The iteration step size was continuously adjusted during the calculation to ensure fast speed and high accuracy. The software will be described elsewhere.' The solutions contained the two metal ions in quite different concentrations. The hydrated electrons generated in the pulse 0 1994 American Chemical Society
10892 J. Phys. Chem., Vol. 98, No. 42, 1994
Ershov et al.
2.0
-I
5
c
t
4
g 1.0 L
W
"200
400
600
4
800
X fnml Figure 1. Absorption spectra of T12+ and TlO. Inset: kinetic traces at two different doses (expressed as concentrations of hydrated electrons). M T12S04 and 1 M propanol-2. Spectra were Solution: 1 x
wo
corrected for the small absorption of the organic radical. reduce the ion that is present in much higher concentration to produce the abnormal valence state, which then reacts with the second metal ion within about 20 ps. As the dose was very low (eaq- concentrations of the order of lop7M), second-order reactions of the intermediates could practically be neglected. The solutions contained 1.O M propanol-2 to scavenge the OH radicals and H atoms, which are also formed in the radiolysis of aqueous solutions. In these scavenging reactions, which occur during the pulse, 1-hydroxymethylethyl radicals, (CH3)2COH, are generated, which have only a weak absorption in the
uv.
The reagents were commercial samples used without additional purification. T1C104 was prepared from T12S04 (99.9995%), as previously described.2b
Results and Discussion Spectra of Tlz+ and TlO. The spectra of these two reduced forms of thallium were previously observed.2b$cThey were remeasured as more sophisticated pulse radiolysis equipment was available which allows to extend the measurements to shorter wavelengths. The new spectra of T1*+ and Tlo are shown in Figure 1. A high concentration of T1+ (1 x loF2 M T12SO4), at which the equilibrium of eq 2 lies on the right side, was used to determine the T12+ spectrum. The solution also contained 1.0 M propanol-2. The spectrum was fully developed at 1 ps after the pulse. It contains the well-known absorption band at 420 nm, E = 1.6 x lo4 M-' cm-'. In addition, a new band was found at 700 nm, which has a small absorption coefficient, E = 1.5 x lo3 M-' cm-', as well as a strong band in the UV, A = 245 nm, E = 2.3 x lo4 M-' cm-I. The spectrum of T1° was determined at 0.7 ps after the pulse by using a solution of low T1+ concentration (2 x M T1+), for which the equilibrium of eq 2 lies on the left side. It contains two broad bands at 450 nm, E = 5 x lo3 M-' cm-', and 260 nm, E = 9 x lo3 M-' cm-'. The 450 nm maximum corresponds to a photon energy of 2.75 eV, which is substantially smaller than the first excitation energy of 3.29 eV of gaseous Tlo. A similar red shift in water was previously reported for the silver atom.* The inset of Figure 1 shows two kinetic traces at 420 nm for different doses. At the lower dose, the absorption slightly increases after the pulse. This is attributed to the additional formation of T12+, as the 1-hydroxymethylethyl radicals generated during the pulse reduce T1+ ions: T1'
+ (CH,),COH
-
T1°
+ (CH3),C0 + €3'
(7)
Figure 2. Kinetic traces at various wavelengths in the presence of 1 x M AgC104. Other concentrations as in Figure 1.
In the previous work, only the reaction of the base form of the radical, which is present at pH = 13, with T1+ had been observed.2b.c The specific rate of reaction 7 is 6 x lo5 M-' s-l. At the higher dose, the absorption after the pulse is practically constant. This is due to the fact that second-order reactions, in which T12+ is consumed, have already occurred, At still higher doses, one observes a decrease in T12+ absorption, as the second-order reactions become dominant; all three absorption bands of T12+ decay at the same rate. Reaction of Tlz+ with Ag+. Figure 2 shows kinetic traces at various wavelengths for a solution containing 1 x lo-* M TlzS04, 1 x M AgC104, and 1 M propanol-2. At 420 nm, the Tlz+ absorption present right after the pulse disappears rapidly. A comparison with the traces in the inset of Figure 1 shows that the Ag+ ion obviously accelerates the Tl2+ decay. At 370 nm, the trace has a maximum at 3 ps after the pulse; the trace at 340 nm goes through a minimum at about 1 ps and through a slight maximum at 7 ps. This complex behavior would be understandable if the reaction of T l 2 + with Ag+ yielded a first product which is not stable and is converted into a second one. The traces in Figure 2 did not change when the dose per pulse was varied between 2 x lo-' and 10 x M. When the concentration of Ag+ was decreased by a factor of 2, the traces essentially kept their shape but were temporally extended by the factor of 2. This indicates that a second Ag+ ion is involved in the conversion of the first reaction product to the second one. The absorption spectrum of the solution at various times after the pulse is shown in Figure 3. At 0.3 ps, the spectrum of T12+ can be seen. As it decays, a new absorption at about 370 nm appears. As the 370 nm absorption disappears, two maxima at 310 and 265 nm are built up. These maxima are typical for small silver clusters such as Ag2+ and Ag32+.9 A mechanism in which the first product rapidly reacts with a second Ag+ ion has recently been proposed for the reactions of Cd+ and Co+ with Ag+.' The first product is a mixed metal cluster; we propose a similar mechanism for the T12+ reaction:
T1;
+ Ag'
-
(TlAg)'
+ T1+
(8)
One could also discuss the formation of the cluster (TlAg)? T1,+ -tAg+
-
(T12Ag)2+
(9)
However, since the free energy of the Tlo-T1+ bond is very
Mixed Metal Clusters in Aqueous Solution
.-, -2.0 LC/*../.T
J. Phys. Chem., Vol. 98, No. 42, 1994 10893
1I
!
c
~
243nm
5 k
c 4
2 1.0 I
W
m
'250
400 A Inml
300
500
s
Figure 3. Absorption spectrum at various times after the pulse (1 x M T12S04,1 x M AgC104, 1.0 M propanol-2). Dose per pulse: 5.5 x lo-' M hydrated electrons.
,
2.5
10
20
30ps
Figure 5. Kinetic traces at various wavelengths. Solution: 1 x M TlzS04, 1 x M CdS04, 1.0 M propanol-2. Dose: 6.6 x lo-' M hydrated electrons.
I
( T I Ag)'
= 2.0 c
-
0
1.5
k
4
g 1.0 L
0.5
0
A Inml Figure 6. Absorption spectrum of a solution containing 1 x lo-* M "@04,5 x M CdS04, and 1 M propanol-2 at various times after M hydrated electrons. the pulse. Dose: 4.7 x
A fnml Figure 4. Absorption spectrum of (TlAg)'
small (-0.17 eV, see above), it appears plausible to expect that the addition of Ag+ to T12+ would lead to the dissociation of a T1+ ion from (T12Ag)2+ because of Coulomb repulsion. We therefore prefer to assign the formula (TlAg)+ to the first reaction product. In analogy to the reaction mechanism described in the previous work, the reaction of this cluster with a second silver ion is proposed: (TlAg)'
+ Ag'
-
Ag2+
+ T1+
(10)
Ag2+, which absorbs at 310 nm, finally reacts with Ag+ to yield Ag32+, Ag,f
+ Ag+
-
Ag32+
(1 1)
which also absorbs at 310 nm and has a second maximum at 265 nme9 The rate constant of reaction 8 could readily be obtained from the pseudo-first-order decay of the T12+ absorption at 700 nm where none of the following products absorb. A value of kg = 3 x 109 M-' s-1 was obtained. We first describe the fitting of the experimental absorption vs time curves in a short time range (< 10 ps) after the pulse. Using a value of klo of 5 x lo9 M-' s-l, it was possible to get very good fits at all wavelengths. The E values of (TlAg)+, which were used in the computer simulation, are shown in Figure 4. The absorption spectrum of the (TlAg)+ cluster has two strong bands at 375 and 270 nm. That these bands do not appear strongly in the course of the reaction (figure 3) is due to the fact that the rate of disappearance of the cluster is greater than that of its formation, the result being that the cluster is not enriched in a substantial concentration. When one tried to fit the kinetic traces also at longer times (> 10 p s ) , where reactions 10 and 11 mainly occur (knowing the spectra of the silver clusters and the specific rate of reaction 119),good fits could not be obtained at all wavelengths. One must therefore conclude that the above scheme does not fully
explain the phenomena which occur after the (TlAg)+ cluster is formed. It is conceivable that an additional intermediate cluster is formed, such as (T1Ag2)2+, which absorbs in the wavelength range of the other intermediates. Reaction of Tlz+ with Cd2+. The absorptions of T12+ at 270, 420, and 700 nm decay rapidly in the presence of 5 x M CdS04. The rate of this decay, which is of pseudofirst order, does not change with the dose. It is concluded that Tl2+ reacts with Cd2+. As can be seen from Figure 5, the kinetic traces are more complicated at other wavelengths. At 300 nm, the absorption is almost constant during the first 10 ps after the pulse and then decays slowly. At 243 nm, there is a fast decay during about 4 ps, followed by an almost constant absorption at longer times. It is concluded that a first product is formed which absorbs almost as strongly as T12+ at 300 nm, slightly less at 243 nm, and much less at 420 nm. This product is transformed into a second one which absorbs at 243 nm roughly as strong as the first one. None of the products have an absorption at 700 nm. The specific rate of the reaction of T12+ with Cd+ was therefore derived from the pseudo-first-order decay of the TlzC absorption at this wavelength. A value of 9 x lo7 M-' s-l was obtained. The absorption spectrum of the pulsed solution is shown in Figure 6 at various times after the pulse. Whereas the T12+ absorption at 420 nm decays substantially, only small changes are seen at wavelengths below 250 nm. No new maxima appear in the visible or near-UV region of the absorption spectrum during the decay of T12+, which would allow one to recognize the reaction of T12+ with Cd2+and the subsequent processes in more detail. However, it can be said that the reaction is not a simple electron transfer,
since the strong absorption band of Cd+ at 300 nm ( E = 1.8 x lo4 M-' cm-l)' did not appear. We therefore conclude that a
10894 J. Phys. Chem., Vol. 98, No. 42, 1994
;:t
Ershov et al. of T1° appears (because of the small Tl+ concentration used, the equilibrium of eq 2 lies on the left side). The inset of the figure shows the buildup of the 460 nm absorption of T1° and its decay at longer times. Knowing the absorption spectra of the species involved, one could readily simulate the kinetic traces. The best fit was obtained by using a specific rate of 5 x lo9 M-' s-l for the reaction
1
I
f 1.0 s I
W
"200
Zn+ 400
300
A lnml
cluster, possibly (CdTl)?+, is the first product,
+ Cd2+- (CdTl)*+ + T1+
(13)
which is not surprising because of the small driving force for electron transfer as mentioned above. Experiments were also carried out with solutions containing a high concentration (2.5 x M) of CdS04 and low concentrations of T12S04, (0.5-2.5) x M. In these solutions, the absorption of Cd+ is present after the pulse, which then decays according to second-order kinetics. The rate of this decay was not influenced by the presence of the T1' ions. It is concluded that Cd+ cannot react with T1+ under our experimental conditions. Reaction of Tlz+ with Pb2+ and of Zn+ with T1+. T12+ also reacts with Pb2+ ions. Solutions containing 2 x lop2M TlC104 and (1 -2) x M Pb(C104)2 were pulsed. The 420 nm absorption of T12+ decayed according to pseudo-first-order kinetics. As the T12+ absorption disappeared, an absorption at 320 nm appeared, which is attributed to Pb+.l0 It is concluded that electron transfer occurred: T1:
+ Pb2+ - 2T1+ + Pb+
(15)
and 2.0 x 1Olo M-' s-l for the specific rate of disappearance of the T1° atoms. (They react with each other as well as with the organic radicals generated during the pulse.2b)
500
Figure 7. Absorption spectrum immediately after the pulse and at longer times. Solution: 4 x lo-* M ZnS04, 1 x M TlzS04, and 1 M propanol-2. Dose: 5.3 x lo-' M hydrated electrons.
T1;
+ TI+ - Zn2+ + ~ 1 '
(14)
This is understood in terms of the great driving force mentioned above. The specific rate of reaction 14 was found to be 2 x 109 M-' s-1. No reaction between T12+ and Zn2+ was found. However, reaction in the reverse direction is possible. A solution containing a high ZnS04 concentration, 4 x M, and low TlzS04 concentrations, ( 2 - 5 ) x M, was pulsed. The absorption of the Zn+ ion was present immediately after the pulse. In the presence of T1+, its decay after the pulse was significantly accelerated. In fact, it followed pseudo-first-order kinetics. Figure 7 shows the absorption spectrum at different times after the pulse. As the absorption of Zn+ disappears, that
Final Remarks In most of the mixed-metal clusters in aqueous solution observed to date, the silver ion is one of the constituents. It seems that Ag+ has a particularly great ability to bind reactive species carrying a lone electron. The formation of pure silver clusters, such as Ag2+ and Ag32+,has been known for a long time, especially by ESR experiments with irradiated frozen silver salt solutions" and with zeolites containing Ag+ ions.12 Our present and previous investigations show that mixed clusters of Ag+ with abnormal valence states of other metals are formed in aqueous solution at ambient temperature. A thalliumcadmium cluster was also found, and one may expect that many more mixed clusters of various metals will be detected in the future. The conditions under which clusters form, i.e., a small difference in the electrochemical potentials of the reacting metal ions, have been outlined above and in the previous paper,' and they have been found to be a suitable guide in the search for new clusters.
References and Notes (1) Ershov, B. G.; Janata, E.; Henglein, A. J . Phys. Chem. 1994, 98, 7619. (2) (a) Cercek, B.; Ebert, M.; Swallow, A. J. J . Chem. SOC.A 1966, 612. (b) Butler, J.; Henglein, A. Radiat. Phys. Chem. 1980, 15, 603. (c) Schwarz, H. A.; Dodson, R. W. J . Phys. Chem. 1989, 93,409. (3) Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1977, 81, 557. (4) Baxendale, J. H.; Dixon, R. S. Z. Phys. Chem. (Munich) 1964,43, 161. (5) Breitenkamp, M.; Henglein, A,; Lilie, J. Ber. Bunsen-Ges. Phys. Chem. 1976, 80, 973. (6) Janata, E. Radiat. Phys. Chem. 1993, 40, 437. (7) Janata, E. Radiat. Phys. Chem., submitted. (8) Ershov, B. G.; Janata, E.; Henglein, A,; Fojtik, A. J. Phys. Chem. 1993, 97,4589. (9) Janata, E.; Henglein, A.; Ershov, B. G. J . Phys. Chem., preceding paper in this issue. (10) Henglein, A.; Janata, E.; Fojtik, A. J . Phys. Chem. 1992, 96,4734. (11) Forbes, C. E.; Symons, M. C. R. Mol. Phys. 1974, 27, 467. (12) (a) Ozin, G. A.; Baker, M. D.; Godber, J. J . Phys. Chem. 1984, 88, 4902. (b) Morton, J. R.; Preston, K. F. In Electronic Magnetic Resonance of the Solid State; Weil, J. A., Ed.; Canadian Society for Chemistry: Ottawa, 1987, p 295. (c) Van der Pol, A.; Reijerse, E. J.; De Boer, E.; Wasowicz, T.; Michalik, J. Mol. Phys. 1992, 75, 37.
