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Complexes of silver(III) with oxoanions
Polyhedron Vol.4, No. 9, pp. 1573-1578, Printed in Great Britain
1985 0
COMPLEXES
0277-5387/85 $3.00 + .OO 1985 Pergamon Press Ltd
OF SILVER(II1) WITH OXOANIONS
JAMES D. RUSH* and LOUIS J. KIRSCHENBAUMt Department of Chemistry, University of Rhode Island, Kingston, RI 02881, U.S.A. (Received
17 January 1985 ; accepted 28 March 1985)
Ahatract-Silver(II1) has a half-life at pH 11 of several hundred seconds in aqueous solutions in the presence of 0.1-1.0 M concentrations of certain basic oxoanions (0x0) (phosphate, carbonate, borate, pyrophosphate, and arsenate). This compares with a lifetime of a few seconds at pH 11 in the absence of these oxoanions. UV-visible spectra and kinetic data for these solutions are interpreted as evidence for the following equilibria in the pH range 9-13. Ag(OH); r + Hz0 G=Ag(OH),H*O + OH -
(1)
Ag(OH), r + 0x0 = Ag(OH),Oxo + OH -
(2)
Ag(OH),Oxo + H,O + Ag(OH),(Oxo)H,O
+ OH -
(3)
Values of K, lie in the range 10e3 c K, < lo4 M for the systems studied. K, is estimated to be N 10’ for phosphate and slightly smaller for the other systems. Ag(OH); undergoes an unusual reaction with pyrophosphate at pH N 8 to form a novel silver(I1) complex, [Ag(P,0,),]6-, for which EPR and electronic absorption spectral parameters are reported.
The tetrahydroxo-argentate(II1) ion, Ag(OH);, is a square-planar complex of tervalent silver that is highly oxidizing and metastable in aqueous solution at pH > 13 (t,,, - 1.5 h in 1.2 M NaOH at room temperature). Its solutions decompose to AgO with resulting oxidation of water. Ag(OH); solutions are prepared electrochemically by oxidizing a silver metal anode in 1.2 M NaOH.‘*2 We have previously investigated the kinetics of Ag(OH); reactions with a number of reductants and complexing agents.3-8 When coordinated by strong ligands, such as periodate or tellurate,2 silver(II1) is quite stable even near pH neutrality. The redox stability of Ag(OH); solutions decreases as the pH is lowered2v7 suggesting that substitution of OH- by solvent water [eqn (l)] is the principal mechanism for decay of Ag(OH);. Ag(OH), + H,O 2 Ag(OH),(H,O)+OH-.
(1)
k-1
* Present address : Department of Chemistry, Building 555, Brookhaven National Laboratory, Upton, L.I., NY 11937, U.S.A. tAddress correspondence to this author at the University of Rhode Island.
The aquo species, Ag(OH),(H,O) rapidly oxidizes the solvent (k = 142 s-l).’ Its instability arises from the poor electron-donating properties of solvent water. AtpH - 11 decomposition of Ag(II1) is controlled by therateofaquation(k, = 2f 1 s-l)’ ofAg(OH);. The reverse reaction in eqn (l), abstraction by OH of a proton from the aquo complex, is relatively slow, presumably because of internal hydrogen bonding. This accounts for the importance of the solvent path for ligand substitution on Ag(II1) even at pH > 13.8 In this report we describe some labile complexes of silver(II1) with oxoanions (phosphate, borate, carbonate, arsenate, and pyrophosphate-general designation “0x0”), which show enhanced kinetic stability at pH < 13. The preparation of these species provides new approaches to the study of silver(II1) reactivity under conditions where the hydroxyl ion concentration is significantly reduced. We also report here the first trapping of an intermediate silver(I1) complex, resulting from the decomposition of the Ag(II1) pyrophosphate complex. This suggests that the oxidation of water to oxygen by Ag(II1) may proceed by one-electron processes.
1573
1574
J. D. RUSH and L. J. KIRSCHENBAUM EXPERIMENTAL
The electrochemical preparation of Ag(OH); has been described previous1y.l Stock solutions for the present study generally contained 3-5 x 10e4 M Ag{OH); as determined at its absorbance maximum at 267 nm (sZ6, = 1.17 x lO,M-‘cm-‘).Dilutionof [OH-] at a constant ionic strength was made using 1.2 M NaClO,. The chemicals used in all preparations were of reagent grade. The sodium hydroxide used was Mallinckrodt 50% solution (low carbonate), and the water was doubly-distilled. Oxoanion solutions were made from the sodium salts, and pH adjustments were made with perchloric acid or NaOH. The pH was measured with a Markson Electromark pH meter calibrated at pH 10 with standard buffers. Solutions of Ag(II1) with oxoanions at pH 9-13 were prepared in two ways: (1) slow addition of perchloric acid to an Ag(OH), solution containing the oxoanion ([0x0] > 0.1 M) while the solution was stirred rapidly, and (2) addition of a neutral or slightly acidic oxoanion solution to Ag(OH); in NaOH. The final pH was adjusted by varying maOH] in the Ag(II1) component prior to mixing. This method was used exclusively in the kinetic and spectral studies because the loss of Ag(II1) during mixing due to local regions of low pH is minimal. Concentrations of 0x0 in excess of 0.1 M were necessary. to neutralize the excess OH-. Because ionic strength could not be uniform for all experiments (the speciation and charges of the oxoanion vary over the pH range of the study), the AgQII) 0x0 preparations were made to a final p (1 M < ,u < 2 M) over which ionic strengths effects are expected to be small. Electronic spectra were obtained on a Cary 15 or a Varian DMS-90 UV-Vis spectrometer with a thermostated reaction cell. The decomposition kinetics of Ag(II1) solutions were measured at 20°C. Observed decay rate constants, k,, were calculated from plots of In (A -A,) vs time which were linear for most runs. In cases where catalytic effects from solid reaction products were evident, k, was obtained from the initial slope of the absorbance vs time data. Absorbance changes were generally monitored at 270 nm. Fast reactions were monitored in an AmincoMorrow stopped-flow apparatus.l A solution of the oxoanion was mixed with Ag(OH);, and absorbance changes were recorded on a storage oscilloscope. Analysis for the + 2 oxidation state of silver was done by EPR of frozen solution samples on a Bruker ER200 spectrometer. The interaction of the unpaired spin in the dg electronic configuration with the
PH
Fig. 1. pH dependence of the absorbance maximum of Ag(II1) phosphate solutions : [phosphate] = 0.35 M ; 20°C. The solid curve is calculated using eqn (4) (see Discussion).
Ag107*10g (I = l/2)
gives a characteristic spectrum.g(*)*(b)The diamagnetic trivalent oxidation state in solution was detected by addition of sodium periodate. The bis periodatoargentate(II1) complex is bright yellow and indefinitely stable.2
RESULTS Spectra of oxoanion solutions. The addition of an oxoanion causes a pH-dependent shift in the absorption maximum of Ag(II1). For 0x0 = borate, phosphate, arsenate, and pyrophosphate, A,,,, changes from 267 to - 276 nm as the pH is lowered from 13 to 10. The spectral data for phosphate shown in Fig. 1 are typical of the other systems as well. At pH > 12.5, the, spectrum is indistinguishable from that of Ag(OH),. In the pH range 11-12 the spectrum broadens and exhibits a small shoulder near 280 nm which merges into a single peak at - 276 nm as the pH is further reduced (Fig. 2). These pH-dependent shifts are reversible upon addition of NaOH. The presence oflabile complexes oftervalent silver was confirmed by the addition of sodium periodate. Over the same pH range, solutions, containing 0.2 M carbonate become orange coloured owing to the formation of a Ag(II1) carbonate complex with the spectrum shown in Fig. 2. The colour change seems to be associated with the pK, of the complex. Ag(II1) borate solutions have spectra similar to that of phosphate when pH > 10. At lower pH, further reactions result in a new spectrum which is also shown in Fig. 2. As with carbonate, the change in spectrum appears to be associated with the pK, of the oxoanion complex. The extinction coefficients of the spectra in Fig. 2 are approximate (+ 15%). The addition of0.2 M pyrophosphate to Ag(OH); with HClO, sufficient to reduce the pH to - 8 (a
Complexes of silver(II1)with oxoanions
IO-
9-
8I
7
i
5 T
6-
I
:
5-
0 x
v
4-
31
0
200
1
1
I
I
I
250
300
350
400
450
500
A, nm
Fig. 2. Solution spectra of Ag(II1)oxoanion complexes at selected pHs. (A): 0.35 M phosphate at pH 10.5.Borate, arsenate, and pyrophosphate solutions are similar at this pH ; (B): 0.4 M carbonate at pH 10.5; (C): 0.6 M borate at pH N 9.5.
buffer region for P,O+-/HP,O:-) causes the immediate formation of a pale yellow-brown solution containing a silver(I1) species with the absorption spectra shown in Fig. 3. Spectrum A is indefinitely stable in 0.1 M pyrophosphate at pH 6-8. Its stability decreases at lower pH. At pH 2-2.5 a species with the spectral peak shown in B predominates which has a half-life of -3 s. A spectrum identical to A is obtained when an Ag(I) solution is oxidized with sodium persulphate under the same conditions. The silver in solutions prepared
I
I
I
I
I
I
I
I I
1 6 T E "5 T I q-J42 * w 3-
1575
from Ag(OH); (frozen at - lOOC, pH = 8.2) had the following EPR parameters : g1 = 2.051, gll = 2.269, and A, = 32.3 G, A,, = 39.3 G. These parameters are typical of many tetragonal Ag(I1) complexes.g At pH 8, the complex is predominately [Ag(II)(P,O,)Jby analogy with the Cu(I1) pyrophosphate system.i’ Products of A&II) oxoanion decomposition. With the exception of pyrophosphate, Ag(II1) oxoanion solutions of pH < 12 decay almost entirely to silver(I) products. Black precipitates which result from slow decomposition of Ag(OH); contain mostly AgO, a mixed Ag(I)/Ag(III) oxide which, in base, may form from the fast disproportionation of Ag(I1)” or the very slow reaction between Ag(OH); and Ag(OH);.” However, the rapid decomposition of Ag(OH), at reduced pH gives Ag(I),’ suggesting that AgO is formed via the latter pathway in strongly alkaline solution. By contrast, the addition of pyrophosphate causes Ag(II1) to be reduced entirely to Ag(I1) over the pH range 8-14, and eventually results in precipitates of AgO at pHs where the decompositions of other Ag(II1) 0x0 systems yield Ag(1) as a product. AgO is detected. by its reaction with bipyridine in acid to give the red bis-bipyridyl Ag(I1) complex.g(b) Formation of A&II) 0x0 complexes. Qualitative observations of the reactions between Ag(OH)i and some oxoanions were made in the stopped-flow apparatus in the pH range 10-l 3. Concentrated 0x0 solutions (0.4 M phosphate, 0.6 M borate and carbonate) which produced a final pH > 12.5 caused no discernable spectral changes when mixed with Ag(OH);. At lower pH a fast decrease in the absorbance (t,,, N 8-15 ms) was monitored at 270 nm in all three systems. The amplitude of the absorbance changes, though not the rates, depended upon the pH. In the cases of borate (pH < 10) and carbonate (pH lO-11.5), the spectra of Fig. 2 were attained by slower processes (tl12 - 3 s and 75 ms, respectively) following the initial, much faster, reaction. Only the faster reaction was observed for the phosphate system. These kinetic observations and the solution spectra (Figs. 1 and 2) show that Ag(OH); is in a pHdependent equilibrium with silver(II1) oxoanion complexes where one or more hydroxyl ligands have been substituted by the oxoanion as in eqn (2).
2-
Ag(OH); + nOxo 5 Ag(OH),_,(Oxo),
+ nOH-.
I-
(2) O
I 250
I 350
I
I 450
I
I 550
Fig. 3.Spectra of Ag(I1)complexes in 0.1 M pyrophosphate at (A): pH 7.8, and (B): pH 2.2.
The relatively small spectral shifts and the single step formation reaction suggest that n = 1 when 0x0 = phosphate. At pH > 10 this condition also seems to hold for all except the carbonate system where
J. D. RUSH and L. J. KIRSCHENBAUM
1576
further substitution may occur. These substitution reactions are reversible upon changing the pH. In the case of oxoanions which have protolytic equilibria in this pH range, a proton ambiguity exists as to the source of the spectral changes upon which eqn (2) is, so far, based. In eqn (2) we have treated “0x0” without regard to its degree of deprotonation. Some justification for this exists. In the range 12 5 pH 5 10 the spectral shifts which occur on the phosphate system (Fig. 1) are qualitatively similar to spectral changes in the borate and pyrophosphate systems, in which the ligands have p&s below this range. Thus, we have attributed the speciation to a single equilibrium, eqn (2), with the realization that ligand protonations might also contribute to the complexity ofthe systems. It should be noted that the rather pronounced spectral changes due to protonations in the borate and carbonate systems may arise from further substitutions or chelation, rather than protonation at a remote site which would not result in a structural change in the complex. Kinetics of the decomposition of Ag(ZZZ) 0x0 solutions. The first-order rate constants for decomposition of Ag(II1) in oxoanion solutions, k,, were measured over the range 9 < pH < 13 at 20°C. At pH > 12.5 half-lives generally exceeded 9 h (except for pyrophosphate preparations) and are very susceptible to the catalytic effects of solids (e.g. AgO) which often complicate the slower reactions of Ag(OH);.4 At lower pH the decay remains firstorder in [Ag(III)], although values of k, varied by as much as 50% over repeat runs due to impurities or the effect of solid products. Rate constants were selected to minimize these effects. The Ag(II1) phosphate system shows typical characteristics and was the most extensively studied. Values of k, obtained at different phosphate concentrations are shown in Table 1. At pH 11 k, is virtually independent of [phosphate] in the range 0.0654.5 M. Lower concentrations could not be achieved because of limitations imposed by the preparative procedure. However, in the absence of
Table 1. Phosphate dependence of k,, (pH 11,2O”C, 1 M (~~22) [Phosphate], 0.5 0.46 0.35 0.21 0.175 0.1 0.065
M
103 k&s-l 2.6 2.5 1.3 2.3 1.6 2.0 1.9
Fig. 4. pH dependence of the rate constant for decay of Ag(II1) Oxoanion systems at 20°C: A, 0, 0.35 M phosphate; B, A, 0.40 M carbonate; C, x, 0.10 M pyrophosphate; D, n, 0.33 M borate; QO.2 M arsenate. Ionic strengths vary between 1 and 2 M.
stabilizing ligands, Ag(II1) decomposes within a few seconds’ at this pH. The pH dependence of k, at constant [0x0] is shown in Fig. 4 for several systems. For Ag(II1) phosphate, log k, varies linearly with pH with a slope of -0.4 in the range 10.3 < pH < 12.5. The slopes increase sharply at a “break point” which occurs between pH 10 and 11 for all of the systems studied. Values of k, for the different systems (except for arsenate) are correlated with solid lines in Fig. 4. In the pH range of greater stability the values of kd vary by less than an order of magnitude between the least stable (borate) and the most stable (phosphate) systems. This presumably reflects differences in the stability constants (K,) of the Ag(II1) 0x0 complexes. The pH dependence of k, in the Ag(II1) carbonate system reflects the-combined effect of two complex p&s on Ag(II1) speciation. This produces a more pronounced dependence above pH 10.5 than is seen in the other systems. The Ag(II1) arsenate system appears to have intermediate stability. However, the scatter in the data obtained for this system causes us to present several observations of k, in Fig. 4 without attempting to correlate them. The acid/base properties of the ligand have no specific effect upon the pH dependence of k, except in those systems where pronounced spectral changes appear to be associated with a pK, of the ligand. The rapid increase in k, between pH 10 and 11 in these
Complexes of silver(III) with oxoanions
1577
Table 2. Equilibrium constants for aquation of Ag(II1) oxoanions differ in structure as well as other oxoanion complexes (K3) (20°C 1 M < p < 2M) properties, their common ability to form hydrogen bonds at several sites suggests a possible mechanism 0x0 Log K, for complex formation. If, as in other reactions of Ag(OH), 3,5 the axial attack on Ag(II1) by the Phosphate - 3.75&-0.2 entering nucleophile occurs, the oxoanion might Borate - 3.2f 0.2 assist in transferring a proton to a hydroxyl ligand -3.25kO.2 Pyrophosphate via a sequence such as those shown in Scheme 1. - 3.7f 0.2 Carbonate Path (a) would result in deprotonation of protonated ligands (e.g. B(OH);, H,PO:-, HAsO;-) while path (b) would be possible for most systems indicates a protonic equilibrium in this species throughout our pH range. Path(b) would be required for pyrophosphate (which is P,O: ~ down range which we suggest is due to reaction (3). to about pH 8) and would be inaccessible to borate Ag(OH),Oxo + H,O both as B(OH); and B(OH),. The five-coordinate Ag(II1) transient may be 2 Ag(OH),(Oxo)(H,O) + OH -. (3) stabilized by the formation of a chelate ring which effectively protonates a hydroxyl ligand. Tetrahedral The qualitatively similar variations in k, (at pH > oxoanions might indeed form two such rings within 10) observed in the phosphate, borate, and pyrophosphate systems support the contention that the the five-coordinate intermediate. The formation of acid/base properties of the Ag(OH), moiety in the four-coordinate entity Ag(OH),Oxo is then Ag(OH),Oxo, rather than those of the ligands, are assisted by elimination of an aquo rather than a the primary cause of the decreasing stability of the hydroxyl ligand. Similar hydrogen bond assisted systems as pH is lowered. Values of log K, for the pathways are important, e.g. in the complexation of various systems were evaluated at the inflection molybdate and tungstate with catecholate points of the data in Fig. 4 and are listed in Table 2. dianion.i3 Transient silver(II1) species in which one OHIn the borate system, k, reaches a constant value of N 10. A complex (spectrum C in Fig. ligand is replaced by a monodentate ligand have -3s-‘topH 2) which decays at a rate almost independent of pH is spectra very similar to that of the parent Ag(OH); 4.5 formed in competition with the pH-dependent decay whereas multidentate ligands typically produce charge-transfer bands at longer wavelengths.6s’4 of its precursor. The formation of a monosubstituted complex, as in Scheme 1, appears to be common to all the oxoanions DISCUSSION studied However, as the pH is reduced, slower Formation reactions. The fast reactions observed processes are observed in the borate and carbonate in the stopped-flow experiments between Ag(OH); systems which appear to be further substitutions. and the 0x0 ligands are too rapid to be mediated by Decomposition reactions. Our analysis of the pH the aquation of Ag(OH); [eqn (l)]; k, - 2 s-l,5 stability data will be qualitative and is based, on indicating that a direct reaction occurs. Although the assumptions which seem to be in accord with the
Ag(OH&
t 0x0
Scheme 1
J. D. RUSH and L. J. KIRSCHENBAUM
1578
experimental observations. These are (1) decay of Ag(II1) proceeds mainly through steady-state amounts of aquated Ag(II1) species, (2) Ag(II1) speciation is basically described by eqns (l)-(3), and (3) coordination by either OH- or 0x0 should affect the hydrolysis properties of the silver(II1) moiety similarly. At pH > 11, where decomposition is slow, the speciation of an Ag(II1) 0x0 solution is described approximately by eqn (4).
C4GW,Oxol C&(111)1,,,.
=
K:[oxo] [OH-]
+K;[Oxo]’
(4)
species. In base (pH > 9) this complex hydrolyzes to
produce AgO precipitates, but at pH 6-8 (following probable addition of a second pyrophosphate ligand) the complex is quite stable. In contrast to this system, the addition of bipyridyl (which also forms strong Ag(I1) complexes) to Ag(OH); solutions does not result in spontaneous one-electron reduction.15 This is probably because the weakly basic bipyridyl nitrogens are unable to mediate a proton transfer to OH- and thus cannot complex Ag(II1) in strong base. Acknowledgements-We
wish to thank Professor P. H. Rieger for his help with the EPR experiments and Professor E. Mentasti for helpful discussions under the
If it is allowed that. the spectral shifts shown in Fig. 1 arise from two species, Ag(OH); and the silver(II1) NATO Collaborative Grants Program (RG 803/83). phosphate complex, which contribute to the shift in peak maximum proportionately to their concentrations, eqn (4) can be applied to obtain a rough REFERENCES estimate for Kz. Using values of A,, = 267 nm for 1. L. J. Kirschenbaum, J. H. Ambrus and G. Atkinson, Ag(OH)i and Iz,, N 276 nm for Ag(OH),(PO,) the Inorg. Chem. 1973,12,2832. data in Fig. 1 yields Ki = 10m2 and the calculated 2. G. L. Cohen and G. Atkinson, J. Electrochem. Sot. curve. 1968,115,1236. The reaction of excess 0x0 with aquo Ag(II1) 3. L. J. Kirschenbaum and J. D. Rush, Znorg. Chem. 1983, with Ag(OH), and species in equilibrium 22,3304. Ag(OH),Oxo [eqns (1) and (311 prevents the rapid 4. E. T. Borish and L. J. Kirschenbaum, Znorg. Chem. decomposition of the solutions even at pH c pK, or 1984,23,2355. pKs. Still, when the pK,s of the complexes are 5. J. D. Rush and L. J. Kirschenbaum, Znorg. Chem. (in press). approached the concentration of reactive aquo 6. L. J. Kirschenbaum and J. D. Rush, J. Am. Chem. Sot. complex, and hence kd, increases. The values of K, 1984,106, 1003. obtained for these systems (see Table 2) suggest that 7. L. J. Kirschenbaum and L. Mrozowski, Znorg. Chem. 10e3 M > K, > 10m5 M, in agreement with earlier 1978,17,3718. estimates.’ 8. L. J. Kirschenbaum, J. Znorg. Nucl. Chem. 1976, 38, The above interpretations of the formation, 881. speciation, and decay of the Ag(II1) oxoanion 9. (a) H. N. PO, Coord. Chem. Rev. 1976,20, 1971 (and complexes is qualitative and accounts only for references therein);(b) W. G. Thorpe and J. K. Kochi, general features. The differences among the J. Znorg. Nucl. Chem. 1971,33,3962. oxoanion systems have not been emphasized. 10. 0. E. Schupp, P. E. Sturrock and J. I. Watters, Znorg. However, it is clear from Fig. 4 that K2 and K3, and Chem. 1963,2, 106. hence kd, depend on the nature of the oxoanion 11. A. Kamar and P. Neta, J. Phys. Chem. 1979,83,3091. 12. Addition of dilute Ag(1) (slightly soluble in base as ligand. Ag(OH);) to an Ag(OH); solution results in the slow The reaction of Ag(OH), withpyrophosphate. This (t 1/Z- 15 min) formation of AgO precipitates. AgO reaction, which at pH < 8 results in the oneitself is insoluble in base. E. T. Borish (unpublished electron Ag(I1) reduction product rather than Ag(1) results). warrants further discussion. This system differs from 13. K. Gilbert and K. Kustin, J. Am. Chem. Sot. 1976,98, the others in the sequence of reactions which follow 5502. eqn (3), aquation of the monosubstituted complex, 14. A. Balikungieri and M. Pelletier, Znorg. Chim. Acta since the pH dependence of k, is quite similar to that 1978,29, 141. of the other oxoanion systems. Evidently, closure of 15. Bipyridyl reduces solutions of Ag(OH); very slowly. the pyrophosphate ring must follow aquation at a The rate of many such redox reactions correlates strongly with the nucleophilicity of the reductant cis-position resulting in a chelated Ag(II1) transient (Ref. 5). which is reduced (by solvent) to form an Ag(I1)
1985 0
COMPLEXES
0277-5387/85 $3.00 + .OO 1985 Pergamon Press Ltd
OF SILVER(II1) WITH OXOANIONS
JAMES D. RUSH* and LOUIS J. KIRSCHENBAUMt Department of Chemistry, University of Rhode Island, Kingston, RI 02881, U.S.A. (Received
17 January 1985 ; accepted 28 March 1985)
Ahatract-Silver(II1) has a half-life at pH 11 of several hundred seconds in aqueous solutions in the presence of 0.1-1.0 M concentrations of certain basic oxoanions (0x0) (phosphate, carbonate, borate, pyrophosphate, and arsenate). This compares with a lifetime of a few seconds at pH 11 in the absence of these oxoanions. UV-visible spectra and kinetic data for these solutions are interpreted as evidence for the following equilibria in the pH range 9-13. Ag(OH); r + Hz0 G=Ag(OH),H*O + OH -
(1)
Ag(OH), r + 0x0 = Ag(OH),Oxo + OH -
(2)
Ag(OH),Oxo + H,O + Ag(OH),(Oxo)H,O
+ OH -
(3)
Values of K, lie in the range 10e3 c K, < lo4 M for the systems studied. K, is estimated to be N 10’ for phosphate and slightly smaller for the other systems. Ag(OH); undergoes an unusual reaction with pyrophosphate at pH N 8 to form a novel silver(I1) complex, [Ag(P,0,),]6-, for which EPR and electronic absorption spectral parameters are reported.
The tetrahydroxo-argentate(II1) ion, Ag(OH);, is a square-planar complex of tervalent silver that is highly oxidizing and metastable in aqueous solution at pH > 13 (t,,, - 1.5 h in 1.2 M NaOH at room temperature). Its solutions decompose to AgO with resulting oxidation of water. Ag(OH); solutions are prepared electrochemically by oxidizing a silver metal anode in 1.2 M NaOH.‘*2 We have previously investigated the kinetics of Ag(OH); reactions with a number of reductants and complexing agents.3-8 When coordinated by strong ligands, such as periodate or tellurate,2 silver(II1) is quite stable even near pH neutrality. The redox stability of Ag(OH); solutions decreases as the pH is lowered2v7 suggesting that substitution of OH- by solvent water [eqn (l)] is the principal mechanism for decay of Ag(OH);. Ag(OH), + H,O 2 Ag(OH),(H,O)+OH-.
(1)
k-1
* Present address : Department of Chemistry, Building 555, Brookhaven National Laboratory, Upton, L.I., NY 11937, U.S.A. tAddress correspondence to this author at the University of Rhode Island.
The aquo species, Ag(OH),(H,O) rapidly oxidizes the solvent (k = 142 s-l).’ Its instability arises from the poor electron-donating properties of solvent water. AtpH - 11 decomposition of Ag(II1) is controlled by therateofaquation(k, = 2f 1 s-l)’ ofAg(OH);. The reverse reaction in eqn (l), abstraction by OH of a proton from the aquo complex, is relatively slow, presumably because of internal hydrogen bonding. This accounts for the importance of the solvent path for ligand substitution on Ag(II1) even at pH > 13.8 In this report we describe some labile complexes of silver(II1) with oxoanions (phosphate, borate, carbonate, arsenate, and pyrophosphate-general designation “0x0”), which show enhanced kinetic stability at pH < 13. The preparation of these species provides new approaches to the study of silver(II1) reactivity under conditions where the hydroxyl ion concentration is significantly reduced. We also report here the first trapping of an intermediate silver(I1) complex, resulting from the decomposition of the Ag(II1) pyrophosphate complex. This suggests that the oxidation of water to oxygen by Ag(II1) may proceed by one-electron processes.
1573
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J. D. RUSH and L. J. KIRSCHENBAUM EXPERIMENTAL
The electrochemical preparation of Ag(OH); has been described previous1y.l Stock solutions for the present study generally contained 3-5 x 10e4 M Ag{OH); as determined at its absorbance maximum at 267 nm (sZ6, = 1.17 x lO,M-‘cm-‘).Dilutionof [OH-] at a constant ionic strength was made using 1.2 M NaClO,. The chemicals used in all preparations were of reagent grade. The sodium hydroxide used was Mallinckrodt 50% solution (low carbonate), and the water was doubly-distilled. Oxoanion solutions were made from the sodium salts, and pH adjustments were made with perchloric acid or NaOH. The pH was measured with a Markson Electromark pH meter calibrated at pH 10 with standard buffers. Solutions of Ag(II1) with oxoanions at pH 9-13 were prepared in two ways: (1) slow addition of perchloric acid to an Ag(OH), solution containing the oxoanion ([0x0] > 0.1 M) while the solution was stirred rapidly, and (2) addition of a neutral or slightly acidic oxoanion solution to Ag(OH); in NaOH. The final pH was adjusted by varying maOH] in the Ag(II1) component prior to mixing. This method was used exclusively in the kinetic and spectral studies because the loss of Ag(II1) during mixing due to local regions of low pH is minimal. Concentrations of 0x0 in excess of 0.1 M were necessary. to neutralize the excess OH-. Because ionic strength could not be uniform for all experiments (the speciation and charges of the oxoanion vary over the pH range of the study), the AgQII) 0x0 preparations were made to a final p (1 M < ,u < 2 M) over which ionic strengths effects are expected to be small. Electronic spectra were obtained on a Cary 15 or a Varian DMS-90 UV-Vis spectrometer with a thermostated reaction cell. The decomposition kinetics of Ag(II1) solutions were measured at 20°C. Observed decay rate constants, k,, were calculated from plots of In (A -A,) vs time which were linear for most runs. In cases where catalytic effects from solid reaction products were evident, k, was obtained from the initial slope of the absorbance vs time data. Absorbance changes were generally monitored at 270 nm. Fast reactions were monitored in an AmincoMorrow stopped-flow apparatus.l A solution of the oxoanion was mixed with Ag(OH);, and absorbance changes were recorded on a storage oscilloscope. Analysis for the + 2 oxidation state of silver was done by EPR of frozen solution samples on a Bruker ER200 spectrometer. The interaction of the unpaired spin in the dg electronic configuration with the
PH
Fig. 1. pH dependence of the absorbance maximum of Ag(II1) phosphate solutions : [phosphate] = 0.35 M ; 20°C. The solid curve is calculated using eqn (4) (see Discussion).
Ag107*10g (I = l/2)
gives a characteristic spectrum.g(*)*(b)The diamagnetic trivalent oxidation state in solution was detected by addition of sodium periodate. The bis periodatoargentate(II1) complex is bright yellow and indefinitely stable.2
RESULTS Spectra of oxoanion solutions. The addition of an oxoanion causes a pH-dependent shift in the absorption maximum of Ag(II1). For 0x0 = borate, phosphate, arsenate, and pyrophosphate, A,,,, changes from 267 to - 276 nm as the pH is lowered from 13 to 10. The spectral data for phosphate shown in Fig. 1 are typical of the other systems as well. At pH > 12.5, the, spectrum is indistinguishable from that of Ag(OH),. In the pH range 11-12 the spectrum broadens and exhibits a small shoulder near 280 nm which merges into a single peak at - 276 nm as the pH is further reduced (Fig. 2). These pH-dependent shifts are reversible upon addition of NaOH. The presence oflabile complexes oftervalent silver was confirmed by the addition of sodium periodate. Over the same pH range, solutions, containing 0.2 M carbonate become orange coloured owing to the formation of a Ag(II1) carbonate complex with the spectrum shown in Fig. 2. The colour change seems to be associated with the pK, of the complex. Ag(II1) borate solutions have spectra similar to that of phosphate when pH > 10. At lower pH, further reactions result in a new spectrum which is also shown in Fig. 2. As with carbonate, the change in spectrum appears to be associated with the pK, of the oxoanion complex. The extinction coefficients of the spectra in Fig. 2 are approximate (+ 15%). The addition of0.2 M pyrophosphate to Ag(OH); with HClO, sufficient to reduce the pH to - 8 (a
Complexes of silver(II1)with oxoanions
IO-
9-
8I
7
i
5 T
6-
I
:
5-
0 x
v
4-
31
0
200
1
1
I
I
I
250
300
350
400
450
500
A, nm
Fig. 2. Solution spectra of Ag(II1)oxoanion complexes at selected pHs. (A): 0.35 M phosphate at pH 10.5.Borate, arsenate, and pyrophosphate solutions are similar at this pH ; (B): 0.4 M carbonate at pH 10.5; (C): 0.6 M borate at pH N 9.5.
buffer region for P,O+-/HP,O:-) causes the immediate formation of a pale yellow-brown solution containing a silver(I1) species with the absorption spectra shown in Fig. 3. Spectrum A is indefinitely stable in 0.1 M pyrophosphate at pH 6-8. Its stability decreases at lower pH. At pH 2-2.5 a species with the spectral peak shown in B predominates which has a half-life of -3 s. A spectrum identical to A is obtained when an Ag(I) solution is oxidized with sodium persulphate under the same conditions. The silver in solutions prepared
I
I
I
I
I
I
I
I I
1 6 T E "5 T I q-J42 * w 3-
1575
from Ag(OH); (frozen at - lOOC, pH = 8.2) had the following EPR parameters : g1 = 2.051, gll = 2.269, and A, = 32.3 G, A,, = 39.3 G. These parameters are typical of many tetragonal Ag(I1) complexes.g At pH 8, the complex is predominately [Ag(II)(P,O,)Jby analogy with the Cu(I1) pyrophosphate system.i’ Products of A&II) oxoanion decomposition. With the exception of pyrophosphate, Ag(II1) oxoanion solutions of pH < 12 decay almost entirely to silver(I) products. Black precipitates which result from slow decomposition of Ag(OH); contain mostly AgO, a mixed Ag(I)/Ag(III) oxide which, in base, may form from the fast disproportionation of Ag(I1)” or the very slow reaction between Ag(OH); and Ag(OH);.” However, the rapid decomposition of Ag(OH), at reduced pH gives Ag(I),’ suggesting that AgO is formed via the latter pathway in strongly alkaline solution. By contrast, the addition of pyrophosphate causes Ag(II1) to be reduced entirely to Ag(I1) over the pH range 8-14, and eventually results in precipitates of AgO at pHs where the decompositions of other Ag(II1) 0x0 systems yield Ag(1) as a product. AgO is detected. by its reaction with bipyridine in acid to give the red bis-bipyridyl Ag(I1) complex.g(b) Formation of A&II) 0x0 complexes. Qualitative observations of the reactions between Ag(OH)i and some oxoanions were made in the stopped-flow apparatus in the pH range 10-l 3. Concentrated 0x0 solutions (0.4 M phosphate, 0.6 M borate and carbonate) which produced a final pH > 12.5 caused no discernable spectral changes when mixed with Ag(OH);. At lower pH a fast decrease in the absorbance (t,,, N 8-15 ms) was monitored at 270 nm in all three systems. The amplitude of the absorbance changes, though not the rates, depended upon the pH. In the cases of borate (pH < 10) and carbonate (pH lO-11.5), the spectra of Fig. 2 were attained by slower processes (tl12 - 3 s and 75 ms, respectively) following the initial, much faster, reaction. Only the faster reaction was observed for the phosphate system. These kinetic observations and the solution spectra (Figs. 1 and 2) show that Ag(OH); is in a pHdependent equilibrium with silver(II1) oxoanion complexes where one or more hydroxyl ligands have been substituted by the oxoanion as in eqn (2).
2-
Ag(OH); + nOxo 5 Ag(OH),_,(Oxo),
+ nOH-.
I-
(2) O
I 250
I 350
I
I 450
I
I 550
Fig. 3.Spectra of Ag(I1)complexes in 0.1 M pyrophosphate at (A): pH 7.8, and (B): pH 2.2.
The relatively small spectral shifts and the single step formation reaction suggest that n = 1 when 0x0 = phosphate. At pH > 10 this condition also seems to hold for all except the carbonate system where
J. D. RUSH and L. J. KIRSCHENBAUM
1576
further substitution may occur. These substitution reactions are reversible upon changing the pH. In the case of oxoanions which have protolytic equilibria in this pH range, a proton ambiguity exists as to the source of the spectral changes upon which eqn (2) is, so far, based. In eqn (2) we have treated “0x0” without regard to its degree of deprotonation. Some justification for this exists. In the range 12 5 pH 5 10 the spectral shifts which occur on the phosphate system (Fig. 1) are qualitatively similar to spectral changes in the borate and pyrophosphate systems, in which the ligands have p&s below this range. Thus, we have attributed the speciation to a single equilibrium, eqn (2), with the realization that ligand protonations might also contribute to the complexity ofthe systems. It should be noted that the rather pronounced spectral changes due to protonations in the borate and carbonate systems may arise from further substitutions or chelation, rather than protonation at a remote site which would not result in a structural change in the complex. Kinetics of the decomposition of Ag(ZZZ) 0x0 solutions. The first-order rate constants for decomposition of Ag(II1) in oxoanion solutions, k,, were measured over the range 9 < pH < 13 at 20°C. At pH > 12.5 half-lives generally exceeded 9 h (except for pyrophosphate preparations) and are very susceptible to the catalytic effects of solids (e.g. AgO) which often complicate the slower reactions of Ag(OH);.4 At lower pH the decay remains firstorder in [Ag(III)], although values of k, varied by as much as 50% over repeat runs due to impurities or the effect of solid products. Rate constants were selected to minimize these effects. The Ag(II1) phosphate system shows typical characteristics and was the most extensively studied. Values of k, obtained at different phosphate concentrations are shown in Table 1. At pH 11 k, is virtually independent of [phosphate] in the range 0.0654.5 M. Lower concentrations could not be achieved because of limitations imposed by the preparative procedure. However, in the absence of
Table 1. Phosphate dependence of k,, (pH 11,2O”C, 1 M (~~22) [Phosphate], 0.5 0.46 0.35 0.21 0.175 0.1 0.065
M
103 k&s-l 2.6 2.5 1.3 2.3 1.6 2.0 1.9
Fig. 4. pH dependence of the rate constant for decay of Ag(II1) Oxoanion systems at 20°C: A, 0, 0.35 M phosphate; B, A, 0.40 M carbonate; C, x, 0.10 M pyrophosphate; D, n, 0.33 M borate; QO.2 M arsenate. Ionic strengths vary between 1 and 2 M.
stabilizing ligands, Ag(II1) decomposes within a few seconds’ at this pH. The pH dependence of k, at constant [0x0] is shown in Fig. 4 for several systems. For Ag(II1) phosphate, log k, varies linearly with pH with a slope of -0.4 in the range 10.3 < pH < 12.5. The slopes increase sharply at a “break point” which occurs between pH 10 and 11 for all of the systems studied. Values of k, for the different systems (except for arsenate) are correlated with solid lines in Fig. 4. In the pH range of greater stability the values of kd vary by less than an order of magnitude between the least stable (borate) and the most stable (phosphate) systems. This presumably reflects differences in the stability constants (K,) of the Ag(II1) 0x0 complexes. The pH dependence of k, in the Ag(II1) carbonate system reflects the-combined effect of two complex p&s on Ag(II1) speciation. This produces a more pronounced dependence above pH 10.5 than is seen in the other systems. The Ag(II1) arsenate system appears to have intermediate stability. However, the scatter in the data obtained for this system causes us to present several observations of k, in Fig. 4 without attempting to correlate them. The acid/base properties of the ligand have no specific effect upon the pH dependence of k, except in those systems where pronounced spectral changes appear to be associated with a pK, of the ligand. The rapid increase in k, between pH 10 and 11 in these
Complexes of silver(III) with oxoanions
1577
Table 2. Equilibrium constants for aquation of Ag(II1) oxoanions differ in structure as well as other oxoanion complexes (K3) (20°C 1 M < p < 2M) properties, their common ability to form hydrogen bonds at several sites suggests a possible mechanism 0x0 Log K, for complex formation. If, as in other reactions of Ag(OH), 3,5 the axial attack on Ag(II1) by the Phosphate - 3.75&-0.2 entering nucleophile occurs, the oxoanion might Borate - 3.2f 0.2 assist in transferring a proton to a hydroxyl ligand -3.25kO.2 Pyrophosphate via a sequence such as those shown in Scheme 1. - 3.7f 0.2 Carbonate Path (a) would result in deprotonation of protonated ligands (e.g. B(OH);, H,PO:-, HAsO;-) while path (b) would be possible for most systems indicates a protonic equilibrium in this species throughout our pH range. Path(b) would be required for pyrophosphate (which is P,O: ~ down range which we suggest is due to reaction (3). to about pH 8) and would be inaccessible to borate Ag(OH),Oxo + H,O both as B(OH); and B(OH),. The five-coordinate Ag(II1) transient may be 2 Ag(OH),(Oxo)(H,O) + OH -. (3) stabilized by the formation of a chelate ring which effectively protonates a hydroxyl ligand. Tetrahedral The qualitatively similar variations in k, (at pH > oxoanions might indeed form two such rings within 10) observed in the phosphate, borate, and pyrophosphate systems support the contention that the the five-coordinate intermediate. The formation of acid/base properties of the Ag(OH), moiety in the four-coordinate entity Ag(OH),Oxo is then Ag(OH),Oxo, rather than those of the ligands, are assisted by elimination of an aquo rather than a the primary cause of the decreasing stability of the hydroxyl ligand. Similar hydrogen bond assisted systems as pH is lowered. Values of log K, for the pathways are important, e.g. in the complexation of various systems were evaluated at the inflection molybdate and tungstate with catecholate points of the data in Fig. 4 and are listed in Table 2. dianion.i3 Transient silver(II1) species in which one OHIn the borate system, k, reaches a constant value of N 10. A complex (spectrum C in Fig. ligand is replaced by a monodentate ligand have -3s-‘topH 2) which decays at a rate almost independent of pH is spectra very similar to that of the parent Ag(OH); 4.5 formed in competition with the pH-dependent decay whereas multidentate ligands typically produce charge-transfer bands at longer wavelengths.6s’4 of its precursor. The formation of a monosubstituted complex, as in Scheme 1, appears to be common to all the oxoanions DISCUSSION studied However, as the pH is reduced, slower Formation reactions. The fast reactions observed processes are observed in the borate and carbonate in the stopped-flow experiments between Ag(OH); systems which appear to be further substitutions. and the 0x0 ligands are too rapid to be mediated by Decomposition reactions. Our analysis of the pH the aquation of Ag(OH); [eqn (l)]; k, - 2 s-l,5 stability data will be qualitative and is based, on indicating that a direct reaction occurs. Although the assumptions which seem to be in accord with the
Ag(OH&
t 0x0
Scheme 1
J. D. RUSH and L. J. KIRSCHENBAUM
1578
experimental observations. These are (1) decay of Ag(II1) proceeds mainly through steady-state amounts of aquated Ag(II1) species, (2) Ag(II1) speciation is basically described by eqns (l)-(3), and (3) coordination by either OH- or 0x0 should affect the hydrolysis properties of the silver(II1) moiety similarly. At pH > 11, where decomposition is slow, the speciation of an Ag(II1) 0x0 solution is described approximately by eqn (4).
C4GW,Oxol C&(111)1,,,.
=
K:[oxo] [OH-]
+K;[Oxo]’
(4)
species. In base (pH > 9) this complex hydrolyzes to
produce AgO precipitates, but at pH 6-8 (following probable addition of a second pyrophosphate ligand) the complex is quite stable. In contrast to this system, the addition of bipyridyl (which also forms strong Ag(I1) complexes) to Ag(OH); solutions does not result in spontaneous one-electron reduction.15 This is probably because the weakly basic bipyridyl nitrogens are unable to mediate a proton transfer to OH- and thus cannot complex Ag(II1) in strong base. Acknowledgements-We
wish to thank Professor P. H. Rieger for his help with the EPR experiments and Professor E. Mentasti for helpful discussions under the
If it is allowed that. the spectral shifts shown in Fig. 1 arise from two species, Ag(OH); and the silver(II1) NATO Collaborative Grants Program (RG 803/83). phosphate complex, which contribute to the shift in peak maximum proportionately to their concentrations, eqn (4) can be applied to obtain a rough REFERENCES estimate for Kz. Using values of A,, = 267 nm for 1. L. J. Kirschenbaum, J. H. Ambrus and G. Atkinson, Ag(OH)i and Iz,, N 276 nm for Ag(OH),(PO,) the Inorg. Chem. 1973,12,2832. data in Fig. 1 yields Ki = 10m2 and the calculated 2. G. L. Cohen and G. Atkinson, J. Electrochem. Sot. curve. 1968,115,1236. The reaction of excess 0x0 with aquo Ag(II1) 3. L. J. Kirschenbaum and J. D. Rush, Znorg. Chem. 1983, with Ag(OH), and species in equilibrium 22,3304. Ag(OH),Oxo [eqns (1) and (311 prevents the rapid 4. E. T. Borish and L. J. Kirschenbaum, Znorg. Chem. decomposition of the solutions even at pH c pK, or 1984,23,2355. pKs. Still, when the pK,s of the complexes are 5. J. D. Rush and L. J. Kirschenbaum, Znorg. Chem. (in press). approached the concentration of reactive aquo 6. L. J. Kirschenbaum and J. D. Rush, J. Am. Chem. Sot. complex, and hence kd, increases. The values of K, 1984,106, 1003. obtained for these systems (see Table 2) suggest that 7. L. J. Kirschenbaum and L. Mrozowski, Znorg. Chem. 10e3 M > K, > 10m5 M, in agreement with earlier 1978,17,3718. estimates.’ 8. L. J. Kirschenbaum, J. Znorg. Nucl. Chem. 1976, 38, The above interpretations of the formation, 881. speciation, and decay of the Ag(II1) oxoanion 9. (a) H. N. PO, Coord. Chem. Rev. 1976,20, 1971 (and complexes is qualitative and accounts only for references therein);(b) W. G. Thorpe and J. K. Kochi, general features. The differences among the J. Znorg. Nucl. Chem. 1971,33,3962. oxoanion systems have not been emphasized. 10. 0. E. Schupp, P. E. Sturrock and J. I. Watters, Znorg. However, it is clear from Fig. 4 that K2 and K3, and Chem. 1963,2, 106. hence kd, depend on the nature of the oxoanion 11. A. Kamar and P. Neta, J. Phys. Chem. 1979,83,3091. 12. Addition of dilute Ag(1) (slightly soluble in base as ligand. Ag(OH);) to an Ag(OH); solution results in the slow The reaction of Ag(OH), withpyrophosphate. This (t 1/Z- 15 min) formation of AgO precipitates. AgO reaction, which at pH < 8 results in the oneitself is insoluble in base. E. T. Borish (unpublished electron Ag(I1) reduction product rather than Ag(1) results). warrants further discussion. This system differs from 13. K. Gilbert and K. Kustin, J. Am. Chem. Sot. 1976,98, the others in the sequence of reactions which follow 5502. eqn (3), aquation of the monosubstituted complex, 14. A. Balikungieri and M. Pelletier, Znorg. Chim. Acta since the pH dependence of k, is quite similar to that 1978,29, 141. of the other oxoanion systems. Evidently, closure of 15. Bipyridyl reduces solutions of Ag(OH); very slowly. the pyrophosphate ring must follow aquation at a The rate of many such redox reactions correlates strongly with the nucleophilicity of the reductant cis-position resulting in a chelated Ag(II1) transient (Ref. 5). which is reduced (by solvent) to form an Ag(I1)