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Comment on “Speciation of aqueous palladium(II) chloride solutions using optical spectroscopies” by C. D. Tait, D. R. Janecky, and P. S. Z. Rogers
0016.7037/93/$6.00
Geochimica el Cosmochimm Acta Vol. 57, pp. 1151-I 156 Copyright 0 1993 Pergamon Press Ltd.Printed in U.S.A.
+ .oO
COMMENT
Comment on “Speciation of aqueous palladium( II) chloride solutions using optical spectroscopies” by C. D. Tait, D. R. Janecky, and P. S. Z. Rogers ROBERT H. BYRNE’ and LEE R. KUMP*
‘Department of Marine Science, University of South Florida, St. Petersburg, FL 33701, USA ‘Department of Geosciences and the Earth System Science Center, The Pennsylvania State University, University Park, PA 16802, USA (Received April 23, 1992; accepted in revised form September 25, 1992) Abstract-Direct experimental evidence, statistical models, and linear free energy relationships indicate that mixed-ligand complex formation is an important aspect of Pd hydrolysis in natural solutions. Within the normal pH range of seawater the dominant hydrolyzed species of Pd( II) is PdC&OH *-. Comparisons of Pd( II) and Pt( II) chemistry indicate that the equilibrium characteristics of the two metals are quite similar, with Pt(I1) forming stronger complexes than Pd( II). Formulations of the solution chemical behavior of metals such as Pd(I1) should be viewed in the context of typical stepwise complexation behavior, statistically predicted complexation relationships, and the behavior of chemical analogs such as Pt(I1). PALLADIUM
INTRODUCIION RECENT EXPERIMENTAL STUDIES ( TAIT et al., 199 1; WOOD,
CHLORIDE
COMPLEXATION
Table 1 presents the results of palladium chloride formation constant determinations in which all four stepwise formation constants were determined. The stepwise formation constants, Kio, in Table 1 can be expressed as
199 1) have raised a number of questions concerning the chemical behavior of Pd in seawater-like solutions. A principal issue raised by TAIT et al. ( 199 1) and explored in our commentary is their contention that “. . . in natural systems, the stepwise replacement of chlorides by hydroxides would not be observed . .“. In contrast, we interpret the experimental data of KUMP and BYRNE ( 1989) as indicative of the importance of mixed-ligand species by just such a process. The formation of mixed-ligand complexes, MCli( 0H)j is a well-known aspect of the solution chemistry of certain metals (e.g., Au3+, T13+, In3+, Hg*+, Pt*+, Pd*+) having strong affinities for chloride ions ( BAES and MESMER, 1976 ). We feel that the basic experimental data of TAIT et al. ( 199 1) are generally meritorious. Our commentary is based upon disagreement in interpretation. The second important issue explored in this commentary is the quality of Pd*+ hydrolysis constant data. The stabilities of simple hydrolyzed palladium species, Pd( OH)j, are not well known. In recent works, WOOD ( 199 1) ultimately favors the hydrolysis characterizations of NABIVANETS and KALABINA ( 1970). However, this interpretation is not supported by the experimental data of WOOD ( 199 1) and is in conflict with the results of KUMP and BYRNE ( 1989). Our discussion begins with a critical assessment of what is perhaps the least controversial aspect of Pd2+ speciation, the formation of palladium chloride complexes. We then proceed to discuss the formation of mixed-ligand complexes in general, and the formation of mixed species PdCli(OH)j and PtCli( 0H)j as special cases of interest. Subsequent to a comparison of theoretically derived and experimentally observed formation constants for the species PdC13(OH)‘-, we examine the most controversial aspect of Pd*+ aqueous chemistry, the formation of simple hydrolyzed species, Pd( OH)j .
[ PdCli] log
KiO
=
1%
[ pdcli_,
and the cumulative expressed as
I
-log[Cl-1,
formation
constants
lOgBiO=log~-ilog[Cl-]
(1 Ii14),
(1)
in Table 1 can be
(2)
where brackets denote the concentration of each chemical species. It is seen that the Ki, results of the five investigations shown in Table 1 are in generally good agreement. The results of DROLL et al. ( 1957), which are in poor agreement with the Table 1 results, have been excluded from the table for a variety of reasons: (a) the results reported by DROLL et al. ( 1957) are appropriate to zero ionic strength while the results shown in Table 1 are appropriate to ionic strengths between 0.8 molar and 1.O molar; inclusion of these results in previous tabulations has led to some confusion; (b) the DROLL et al. ( 1957) results are in error (in some cases by an order of magnitude) because they were obtained at high ionic strength (these data were not reported ) and extrapolated to zero ionic strength using the Debye-Hiickel equation (the Debye-Hiickel equation is an appropriate description of activity coefficient behavior only at very low ionic strengths); (c) the quality of the DROLL et al. ( 1957) formation constant analysis was compromised by impurities in their experimental solutions (WEED, 1964; WOOD, 1993 ) . Due to these problems, the use of thermodynamic data derived from the DROLL et al. ( 1957) results ( SASSANI and SHOCK, 1990) is inadvisable. II51
R. H. Byrne and L. R. Kump
1152
Table 1: Stoicbiometric Formation Constants for PI-Cl Species in Cl--C104*Mixtures
Burgerand etal. nyrssen (I%0 (1963) 1=0.82ooc I=l.O25% [H+l=0.6M [H+]=l.OM
Shchukamv
IO8K,
Biiukov Eldiing andShIen(1972) Average Average (ET) F$gkyi I=I.O25oC I=l.O25oC 108Kio lo8Pi0 [H$O.7M [H+]=l.OM [H+]=l.OM
IO8KIO
4.34
3.88
3.98
4.4
4.47
4.21fo.27
4.21
108Kzo
3.54
3.03
3.24
3.34
3.29
3.29i0.18
1.50
log K30
2.68
2.18
2.3
2.34
2.41
2.38M.19
9.88
‘og &a
1.68
1.34
(2.0)’
1.38
1.37
1.44i0.29
11.32
log b~%o
-0.80
-0.85
-0.74
-1.06
-1.18
-0.93Zto.19
I
log K4&30
-0.85
-0.86
108K3dKzo
I
-1.00
-0.84
-0.94
-0.88
-1.00
I
I
1 (-0.3)’
1
I
-0.96
-0.91&0.06 I
-1.04
I
-0.96m.09 1
a. These. data have been excluded from the averaged results
Examination of the summary results in Table 1 demonstrates that to a good approximation, at 25°C and 0.8 M 5 I I 1.0 M, log Kio/Kti-l)o = -0.93. Consequently, we can approximate the stoichiometric formation constants with the following empirical relation: log Kio = 4.21 - 0.93(i - 1).
(3)
This relation, which summarizes all of the formation constant data shown in Table 1, directly indicates that the Pd*+ complexes PdCl+, PdCl:, and PdCl; have similarly sized predominance fields on the order 0.93 log units wide. The log K40/K30 result obtained by TAUTet al. ( 199 1) at zero ionic strength (log K4o/K3o = - 1.3) is in fair agreement with the Table 1 summary of stability constant results. MIXED-LIGAND PALLADIUM COMPLEXES A number of works have examined Pd(I1) hydrolysis in chloride solutions. Equation 3 shows that for chloride concentrations greater than approximately 1O-‘.4 and ionic strengths near unity the dominant chloride complex of Pd( II) will be the species PdCl $- . Palladium hydrolysis at high chloride concentrations thus can be interpreted in terms of the stepwise replacement of chloride ions in this complex by OH _. The first step in a series of replacement reactions can be written as PdCl:- + OH- 2 PdC&OH*- + Cl-
(4)
where log
K3,
=
log
[PdC130H2-1 + log [PdCl:-]
[Cl-] -log
[OH-]. (5)
Assuming that the absorbance of hydrolyzed species at 290 nm is negligible (which is consistent with the modeling of KUMP and BYRNE, 1989), Eqn. 5 shows that when the solution absorbance of PdCl:- has fallen by one half due to the formation of PdC130H *- , log K3, can be directly calculated as log ([Cl-]/ [ OH -1). The results of KUMP and BYRNE ( 1989 ) indicate that in normal salinity seawater the concentrations of PdC&OH 2- and PdCl<- are approximately equal
when pH = 8.7. In this case, assuming log [H+][OH-] = -13.75 (BAESand MESMER,1976), it follows that, at 25°C and ionic strength = 0.7 M, log K,, = 4.8.
(6)
The PdCl$- absorbance vs. pH depictions of TAIT et al. ( 199 1) at log [Cl-] = -0.25 can be used to derive a log K3, result essentially equal to the result given above. However, TAIT et al. ( 1991) concluded that their results (supporting the importance of mixed-ligand species) should not be applied to studies of the natural behavior of Pd in seawater. They presumed that in their experiments, the association of Pd*+ and OH- depleted the free OH- concentration in solution to a point where no additional OH- was available to further replace Cl- associated with Pd*‘. This misinterpretation constitutes our major point of disagreement with TAIT et al. ( 1991). Since the formation of metal hydroxide complexes can be conceptualized in terms of hydrolysis equilibria, M*+ + nH20 = M(OH)f-’
+ nH+
(7)
and MCI:-i + jH20 = MQ_j(OH)f-’
+ jH+ + jCl-,
(8)
hydroxide can never be a limiting reagent. Thus, the Fig. 5 stability field diagram of TAIT et al. ( 199 1) which depicts a broad predominance field for mixed PdCli(OH)j species is not conditioned on high palladium/ hydroxide concentration ratios. These results of TAIT et al. ( 199 1) are appropriate to natural solutions. Our experimental result given in Eqn. 6 is in reasonable agreement with the results of KAZAKOVAand PTITSYN ( 1967) and MILIC and BUGARCIC ( 1984) as well as the results of TAIT et al. ( 199 1). In one molar NaCl at 25°C the result given by KAZAKOVAand PTITSYN (1967) is log K3, = 5.7. Recognizing that the symbol Pd*+ in the experimental work of MILIC and BUGARCIC( 1984) represents all nonhydrolyzed Pd species (including chlorides) under their experimental conditions (one molar NaCl at 25’C), their data can be used to obtain the result log K3, = 4.45. Thus, available experimental evidence indicates that 4.4 I log K3, 5 5.7. We prefer
1153
Discussion: Comment our Eqn. 6 estimate for log KS,, since it was obtained at relatively low total Pd concentrations (5 to 10 PM), minimizing potential problems attributable to the formation of polymeric Pd( II) complexes.
metal coordination sphere. This term is generally considered negligible for substitutions of equally charged ligands. We will consider the case i + j = 4, for both Pt(ll) and Pd( II) complexes. The statistically predicted Pt( II) results shown in Table 2 were obtained using the result log fl& Pt ) = 13.99 (ELDING, 1978) and a cumulative Pt(OH):- formation constant (log PO4= 30.4 1) derived from the stepwise PtCl:- hydrolysis constants of PESHCHEVITSKIet al. ( 1962). The statistically predicted Pd(l1) results shown in Table 2 were obtained using the Table 1 result log &(Pd) = 11.32 and the result log Po4(Pd) = 26.4 of NABIVANETSand KALABINA( 1970). The directly observed results in Table 2 were obtained using the above log pea results plus the following equations,
STATISTICALLY PREDICTED COMPLEXATION CONSTANTS The formation constants of mixed-l&and species for the metals Au, Tl, In, and Hg can be predicted from statistical arguments ( BAES and MESMER, 1976), and such predictions are generally in good agreement (within an order of magnitude) with direct observations. Statistical predictions for mixed-ligand species, MClI( OH)j are based upon the use of cumulative complexation constants, /3Ioand &j, ho
[MChl = [M2+l[cl_li
and
Poj =
[M(OH)jl
[M’+][OH-]I
(9)
log P3, = log 840 + 1% K3,;
log
P22 =
log
P3 I +
log
log
log
Pod =
1%
P13 +
lois KM,
013
=
lOis 022
+
log
K13;
with log BIo= C log Kio and log poj = c log Z&j.
(12)
i
where the constants Kij (Pt) are given as ( PESHCHEV~TSKI et al., 1962)
Formation constants for mixed-ligand Complexes can be calculated using the equation of BELEVANTSEV et al. ( 1982): log
Pij
=
&
1%
P(i+j)O
+
&
log
K3,tRj
PO(i+j)
K
(10)
-i(aij+lOg6ij)
22
tRj
= [~C130H2-I[C1-l [nCl$-][OH-] =
= 105’08;
PG(OW:-l[Cl-1 [PtC130H2-][OH-]
where log p-u = log
B-XOW:-l[Cl-1
[MC1i(oH)jl - i log [Cl-]
=
103.79.
’
Ko4(R)
= ~RCI~OH~~_lIOH_l
[WOW:-l[Cl-1
Table 2: Observed and Statistically Modeled Mixed-Ligand Formation Constants (O.SMB 5 l.OM, 2S“C) Observed
log Pij P40 @)
13.99
IO881 (PO
19.08
log p22 m
I
= 103.24
(13)
and log K3,( Pd) = 4.8 (this work, Eqn. 6). It is seen in Table 2 that the directly observed and statistically predicted stability constants are in generally good agreement. The statistically predicted constants are smaller than directly observed values by 0.2 to 0.4 log units, possibly due to ligand effects.
The term aij in Eqn. 10 accounts for the well-known promotion of mixed-l&id complex formation, relative to single ligand-type species, through statistical effects. In the case that i + j = 4, the term “II = -0.402 ( BELEVANTSEV et al., 1982 ). The term log 6Ij accounts for “ligand effects” which include steric hindrance, ligand-ligand bonds, and changes in the
23.38
log PI3PO
27.17
IO8PO4(n)
30.41
IO8P.IO (Pd)
11.32
1%PSI@‘dd)
16.12
PICdjCtd
18.72 I
23.04 26.96
15.69
log &a (Pd)
19.66
h? h 8d)
23.23
log PM W)
=
- j log [OH-].
(11)
I
= 104’30’
[PtC12(OH):-][OH-]
K’3(R)
Df*+l
1%
K22 ;
26.4
I
R. H. Byrne and L. R. Kump
1154 LINEAR FREE ENERGY RELATIONS
Linear free energy calculations provide an additional means of estimating log &(Pd). The comparison of PtX:- and PdX:- formation constants of HANCOCKet al. ( 1977) indicated that log &( Pd) = 0.8 log &( Pt). The &, results shown in Table 2 indicate that log &,( Pd) = 0.809 log &,(Pt), and the PO4results shown in Table 2 indicate that log &(Pd) = 0.868 log &(Pt), Using the average of our Table 2 &, and PO4results for Pt and Pd, log &( Pd) = 0.84 + 0.03 log &( Pt ), in addition to the directly observed P3, (Pt) value shown in Table 2, log &,(Pt) = 19.08, the cumulative PdCl,OH 2- fo~ation constant is calculated as log &, (Pd) = 16.0. This value, combined with the observed value for log ,&,( Pd) in Table 1, yields log X;, (Pd) = 4.7, a value that compares favorably to our experimental determination (KUMP and BYRNE, 1989; Eqn. 6). Thus, our log &(Pd) and log KS1(Pd) results are in good agreement with (a) the direct experimental observations of others ( TAIT et al., 199 1; MILK and BUGARCIC, 1984; KAZAKOVA and PTITSYN, 1967), (b) statistical predictions involving the &( Pd) and &f Pd) results of others, and (c) linear free energy relationships employing the &, (Pt) result obtained from the works of PESHCHEVITSKIet al. ( 1962 ) and ELDING( 1978). In contrast to the conclusion of TAIT et al. ( 199 1 ), mixed-ligand complexes PdCIi(OH)j should be significant species under natural conditions. KINETICS OF HYDROLYSIS WOOD ( 1991 ) indicated that if the hydrolysis results of NABIVANETSand KALABINA( 1970) are correct, Pd( II) hydrolysis should be much more extensive than the degree of hydrolysis observed by KUMP and BYRNE ( 1989). WOOD ( 199 1) suggested that equilibrium may not have been attained in the work of KUMP and BYRNE( 1989), and cited the works of KAZAKOVAand PTITSYN ( 1967) and TAIT et al. ( 199 1) as evidence that equilibrium with respect to hydrolysis is attained only after several days. The work of KAZAKOVAand PTITSYN( 1967) does not support this suggestion. KAZAKOVA and PTITSYN ( 1967) indicated that PdCl:- reacts very rapidly
with OH -, and modeled their titration data in terms of Eqn. 4 by assuming rapid replacement of Cl- by OH- followed by slow processes corresponding to polymerization. The data of KUMP and BYRNE ( 1989) were obtained through rapid titration in which the replacement of Cl- by OH- occurred in minutes to seconds followed by slower changes which we interpreted as polymerization. The work of TAIT et al. ( 199 I) did not produce evidence that the reaction shown in Eqn. 4 is a slow process. TAIT et al. ( 199 1) observed slow hydrolysis, as have all experimentalists investigating Pd*+ hydrolysis, but did not distinguish between the possibility of rapid monomeric reactions and slow polymeric reactions. TAIT et al. ( 1991) appear to assume that all Pd2+ hydrolysis reactions are slow based on the slow kinetics of Pt2+ hydrolysis (WV et al., 1990). It must be noted, however, that the kinetics of Pt*+ complexation is typically many orders of magnitude slower than the kinetics of Pd*+ complexation ( ELDING, 1973; BAESand MESMER, 1976). PURE HYDROLYSIS OF PALLADIUM Although investigation of Pd2+ hydrolysis in chloride solution is demanding, investigation of Pd2+ hydrolysis in noncomplexing media (e.g., perchlorates) appears to be even more challenging. Pd*+ has strong affinities for a wide variety of ligands, making concerns over solution impurities of extreme importance. Pd*+ has strong tendencies toward polymerization in perchlorate solutions over the range of pH in which the tmnsitions between Pd*‘, PdOH+, and Pd(OH): occur ( NABIVANETSand KALABINA,1970). Table 3 compares previous quantitative assessments of Pd’+ hydroxide complexation constants, &j( Pd) and &j( Pd) where
Poj(Pd)=
[Pd(OH)jl [pd2+l[oH_]j
lPd(OH)jl ~“‘(Pdf = [Pd(OH)j_~][OH-]
.
(14)
The stepwise stability constant behavior in Table 3 is very unusual. Whereas it is generally expected that the ratio of stepwise stability constants is much smaller than one (e.g.,
Table 3: Formation Constants for Palfadium Hydroxide Species ,
and
1155
Discussion: Comment log K20/K,O = -0.9, Table 1) the ratio of K&Ko, obtained by IZATT et al. ( 1967) is as large as five. According to the results shown in Table 3, addition of OH- to Pd2+ actually increases the affinity of palladium for OH- : PdOH+ has a larger affinity for OH - than does Pd’+ . The stepwise hydrolysis constant results of NABIVANETSand KALABINA ( 1970) are especially unusual in that Ko2/KoI > 1 and Ko, /I& < 10 -I’. Given the results which have been obtained for PdClI formation ( KzoJKlo _NK301Kzo N Km/ Kjo), and the absence of any explanation for the unusual stepwise stability constant behavior shown in Table 3, we feel that some of the data in Table 3 may be greatly in error. We suspect that the data of IZATT et al. ( 1967) are in error because the authors assumed negligible polynuclear species formation. IZATT et al. ( 1967) used total Pd concentrations on the order of one millimolar and greater, and equilibrated their solutions for five hours prior to each measurement. In contrast, NABIVANETSand KALABINA ( 1970) carefully accounted for the substantial influence of polymerization in their solutions. We suspect that the principal failing of the NABIVANETS and KALABINA ( 1970) experimental procedures was an inability to measure the shape of the Pd( II) solubility curve at total Pd( II) concentrations much lower than 1 X 10e5 molar. If the apparent solubility minimum observed by NABIVANETS and KALABINA( 1970) were somewhat lower than 1 X 10e6 molar, the equilibrium constants PO,and PO2would be smaller, and likely consistent with our observation of the dominance of PdCl:- in experimental seawater solutions even at pH = 8.3. It is interesting, and important, to note that a lower minimum in the Pd( II) solubility curve of NABIVANETSand KALABINA( 1970) would not affect their estimate for &: using the Fig. 1 solubility curve of NABIVANETSand KALABINA ( 1970), it is seen that total dissolved monomeric Pd concentrations are equal to 1 X 10m3molar at both pH = 0.5 and pH = 13.6. Because the free Pd*+ concentration in a solubility equilibrium with Pd( OH)201 decreases by a factor of 1026.2(i.e., 10ZApH)between pH = 0.5 and pH = 13.6, it follows that [Pd”] = 10e3- 10-26.2 at pH = 13.6 and [ Pd(OH):-] = 10e3. Noting that the [OH-] concentration
6
7
8
9
10
PH
FIG. 1. Predominance diagram for palladium chloride and hydroxide species in seawater (S = 35, 25”C, 1 bar) based on the formation constants recommended in the text. Numbers on graph indicate the i, j values for the species PdCli(OH )pi’.
Table 4: Proposed Hydrolysis Constants for Palladium (OJM sI 5 l.OM, ZS’C!)
log PO4
26.5
at pH 13.6 is approximately IO-‘.*, the formation constant log PO4is then calculated as log PO4= log ‘p;E;;‘-l
- 2 log [OH-]
= 26.2 + 0.4 = 26.6.
(15)
This result is in good agreement with the reported result of NABIVANETSand KALABINA( 1970). We suspect this result may be the only truly quantitative formation constant which can be directly derived from their work. A final extension of this analysis of the hydrolysis of Pd can be made by assuming that the ratios of Kocj+l,/Koj for stepwise Pd2+ complexation are similar for OH- and Cl-. If Kocj+l,fKoj = 0.1, then /304= K~,(O.1K,,)(O.O1Ko,)(O.OO1Ko,)
= 1026.5
where 1026.5is an average of our estimate and the estimate OfNABIVANETSand KALABINA( 1970). The cumulative OHcomplexation constants, thus obtained, are shown in Table 4. It is interesting to note that the recent results of WOOD ( 199 1) shown in Table 3 are in much better agreement with these results than the Bo3and PO2results of NABIVANETSand KALABINA( 1970). A MODEL
FOR THE SPECIATION OF PALLADIUM IN SEAWATER
Through the use of a combination of experimental results, linear free energy relationships and statistical modeling, we have developed an internally consistent set of formation constants for PdCli(OH)j species in solutions at intermediate ionic strength. Recommended log fiio results are found in Table 1. Our recommended log p3I (Pd) result is the directly observed value in Table 2, log &, (Pd) = 16.1. Noting that the statistically predicted log 83, (Pt ) and log ,822(Pt ) results underestimate observed results by about 0.35 log units (Table 2), our recommended p22and /?I3values for Pd are log f12*(Pd) = 20.0 and log &(Pd) = 23.6. Suggested formation constants for Pd( 0H)j are shown in Table 4. Figure 1 shows the predicted speciation of Pd(l1) in seawater using our recommended constants. The species PdCl$- dominates throughout the normal range of seawater pH, 7.4 I pH 5 8.4. The next most abundant species are the three-chloride complexes PdC&OH 2- and PdCl 5 . Mixedligand complexes, PdC&OH 2- and PdC120H :-, become dominant above pH 8.7. Interpretation of palladium hydrolysis experiments is presently controversial. Indeed, in our earlier work ( KUMP and BYRNE, 1989), we were careful to report our experimental data in terms of the observed extent of hydrolysis
R. H. Byrne and L. R. Kump
1156
without specifying the relative importance of single-ligandtype hydrolyzed species (Pd( 0H)j) vs. mixed-ligand hydrolyzed species (PdCli( OH)j). Herein we have extended our analyses, by using statistical and linear free energy calculations, to elucidate the potential importance of mixed-ligand species. Clearly, these arguments are subject to considerable uncertainty and cannot bc substituted for direct observations. However, even order of magnitude uncertainties are dwarfed by the range of hydrolysis constants which have recently been reported for the species Pd(OH)i and Pd(OH); (Table 3). Our hydrolysis observations in seawater ( KUMP and BYRNE, 1989), plus the arguments advanced in the present work, indicate that PdC&OH ‘- is the dominant, hydrolyzed form of Pd within the normal pH range of seawater. Nevertheless, we know of no direct experimental observations which preclude the possibility that we have underestimated the significance of Pd(OH); due to very slow formation kinetics for this species. Although we feel there is little reason to suspect that the kinetics of Pd( OH)! formation is substantially slower than the formation kinetics of PdC130H2- or PdC12(OH):-, we are currently initiating long-term experiments to investigate this possibility.
CONCLUSION In this paper we have quantitatively summarized previous Pd( II) chloride complexation results and quantitatively compared these results with the results of TAIT et al. ( 199 1). We have produced recommended formation constants for the species PdC130H2- and other mixed species, and have shown that our recommended PdCljOH *- data are in agreement with (a) direct experimental evidence including the work of TAIT et al. ( 199 1 ), (b) statistical predictions based on the well-known characteristics of mixed-ligand complexation, and (c) linear free energy relationships using experimentally observed Pt( II) complexation characteristics. Finally, we have produced recommended Pd( OH)j hydrolysis constants which are consistent with Pt( OH):- hydrolysis data and which are in better agreement with the recent data of WOOD ( 199 I ) than the data of NABIVANETSand KALABINA (1970). We feel that our narrative produces a coherent picture of Pd( II) complexation and provides a framework within which future experimental investigations of both Pd( II) and Pt( II) complexation can be assessed.
Acknowledgments-This work has been partially supported hy grants from the Marine Chemistry and Geology and Paleontology Programs of the National Science Foundation. We express our gratitude to Drew Tait, Scott Wood, and an anonymous reviewer for their thoughtful critiques of the manuscript.
Editorial handling: G. Faure
REFERENCES E. and MESMER R. E. ( 1976) The Hydrolysis of Cations. J. Wiley & Sons. BELEVANTSEV V. I., KOLONING. R., and PESHCHEVITSKIY B. I. ( 1982) Use of stepwise ligand replacement in estimating the stability constants of mixed complexes in geochemically important systems. Geochem. Intl. 19, 169-179. BIRYUKOVA. A. and SHLENSKAYAV. I. (1964) Composition and stability constants of chloro-complexes of palladium( II). Russ. .I. Inorg. Chem. 9,450-452. BURGERK. and DYRSSEND. ( 1963) On the complex formation of palladium with dimethylglyoxime. Acta Chem. &and. 17, 14891501. DROLL H. A., BLOCKB. P., and FERNELIUSW. C. (1957) Studies on coordination compounds. XV. Formation constants for chloride and acetylacetonate complexes of palladium( II). J. Phys. Chem. 61, lOOf-1004. ELDINGL. I. ( 1972) Palladium( II) halide complexes. I. Stabilities and spectra of palladium( II) chloro and bromo aqua complexes. Inorg. Chim. Acta 6,647-65 1. ELDINGL. I. ( 1973) Palladium( II) halide complexes. III. Acid hydrolyses and halide anations of cis- and trans-Dichlorodiaquapalladium( 11)and -Dibromodiaquapalladium( II). Znorg. Chim. Acta 7,58 l-588. ELDINGL. I. ( 1978) Stabilities of platinum(B) chloro and bromo complexes and kinetics for anation of the tetraaquaplatinum( II) ion by halides and thiocyanate. Inorg. Chim. Acta 28, 255-262. HANCOCKR. D., FINKELSTEIN N. P., and EVERSA. ( 1977) A linear free-energy relation involving the formation constants of palladium( II) and olatinum( II 1.J. Inure. NW/. Chem. 39. 1032- 1034. IZATTR. M., EATOUCHD., and CH&TENSEN J. J. ( 1967) A study of Pd2+( aq) hydrolysis. Hydrolysis constants and the standard potential for the Pd, Pd’+ couple. J. Chem. Sot. (A), 1301-1304. KAZAKOVAV. I. and PTITSYNB. V. f 1967 ) Hvdrolvsis of haloeenocomplexes of palladium. Russ. I. Inorg. ‘Chem. i2, 323-326. KUMP L. R. and BYRNER. H. (1989) Palladium chemistry in seawater. Environ. Sci. Tech. 23, 663-665. MILIC N. B. and BUGARCICZ. D. ( 1984) Hydrolysis of the palladium( 11)ion in a sodium chloride medium. Transition Met. Chem. 9, 173-176. NABIVANETS B. 1. and KALABINAL. V. ( 1970) State of palladium(B) in perchlorate solutions. Russ. I. Inorg. Chem. 15, 8 18-82 1. PESHCHEVITSKI B. I., PTITSYN B. V., and LESKOVA,N. M. ( 1962) Izv. Sibirsk. Otd. Akad. Nauk SSSR 11,143- 145 (Cited in: Stability Constants ( 1984), The Chemical Society, London, p. 284). SASSANID. C. and SHOCKE. L. ( 1990) Speciation and solubility of palladium in aqueous magmatic-hydrothermal solutions. Geology 18,925-928. SHCHUKAREVS. A., LOBANEVA0. A., IVANOVAM. A., and KoNONOVA M. A. ( 196 I ) Spectrophotometric investigations of the chloride complexes of bivalent palladium in aqueous solutions. Vestnik. Leningrad Univ. 16, (Ser. Fiz. i Khim. No. 2), 152-155. TAIT C. D., JANECKYD. R., and ROGERSP. S. Z. ( 199 1) Speciation of aqueous palladium( II) chloride solutions using optical spectroscopies. Geochim. Cosmochim. Acta 55, 1253- 1264. WEEDE. D. ( 1964) A study of the palladium( II)-chloride complexes in aqueous solution. Ph.D. dissertation, Ohio State University. WOOD S. A. ( I99 I ) Experimental determination of the hydrolysis constants of Pt2+ and Pd’+ at 25°C from the solubility of Pt and Pd in aqueous hydroxide solutions. Geochim. Cosmochim. Acta 55, 1759-1761. WOOD S. A., MOUNTAIN B. W., and PAN P. (1993) The aqueous geochemistry of platinum, palladium, and gold: Recent experimental constraints and a reevaluation of theoretical predictions. Canadian Mineral. (in press). WV L., SCHWEDERSKI B. E., and MORGAND. W. (1990) Stepwise hydrolysis kinetics of tetrachloroplatinate( II) in base. Inorg. Chem. 29, 3578-3584. BAES C.
Geochimica el Cosmochimm Acta Vol. 57, pp. 1151-I 156 Copyright 0 1993 Pergamon Press Ltd.Printed in U.S.A.
+ .oO
COMMENT
Comment on “Speciation of aqueous palladium( II) chloride solutions using optical spectroscopies” by C. D. Tait, D. R. Janecky, and P. S. Z. Rogers ROBERT H. BYRNE’ and LEE R. KUMP*
‘Department of Marine Science, University of South Florida, St. Petersburg, FL 33701, USA ‘Department of Geosciences and the Earth System Science Center, The Pennsylvania State University, University Park, PA 16802, USA (Received April 23, 1992; accepted in revised form September 25, 1992) Abstract-Direct experimental evidence, statistical models, and linear free energy relationships indicate that mixed-ligand complex formation is an important aspect of Pd hydrolysis in natural solutions. Within the normal pH range of seawater the dominant hydrolyzed species of Pd( II) is PdC&OH *-. Comparisons of Pd( II) and Pt( II) chemistry indicate that the equilibrium characteristics of the two metals are quite similar, with Pt(I1) forming stronger complexes than Pd( II). Formulations of the solution chemical behavior of metals such as Pd(I1) should be viewed in the context of typical stepwise complexation behavior, statistically predicted complexation relationships, and the behavior of chemical analogs such as Pt(I1). PALLADIUM
INTRODUCIION RECENT EXPERIMENTAL STUDIES ( TAIT et al., 199 1; WOOD,
CHLORIDE
COMPLEXATION
Table 1 presents the results of palladium chloride formation constant determinations in which all four stepwise formation constants were determined. The stepwise formation constants, Kio, in Table 1 can be expressed as
199 1) have raised a number of questions concerning the chemical behavior of Pd in seawater-like solutions. A principal issue raised by TAIT et al. ( 199 1) and explored in our commentary is their contention that “. . . in natural systems, the stepwise replacement of chlorides by hydroxides would not be observed . .“. In contrast, we interpret the experimental data of KUMP and BYRNE ( 1989) as indicative of the importance of mixed-ligand species by just such a process. The formation of mixed-ligand complexes, MCli( 0H)j is a well-known aspect of the solution chemistry of certain metals (e.g., Au3+, T13+, In3+, Hg*+, Pt*+, Pd*+) having strong affinities for chloride ions ( BAES and MESMER, 1976 ). We feel that the basic experimental data of TAIT et al. ( 199 1) are generally meritorious. Our commentary is based upon disagreement in interpretation. The second important issue explored in this commentary is the quality of Pd*+ hydrolysis constant data. The stabilities of simple hydrolyzed palladium species, Pd( OH)j, are not well known. In recent works, WOOD ( 199 1) ultimately favors the hydrolysis characterizations of NABIVANETS and KALABINA ( 1970). However, this interpretation is not supported by the experimental data of WOOD ( 199 1) and is in conflict with the results of KUMP and BYRNE ( 1989). Our discussion begins with a critical assessment of what is perhaps the least controversial aspect of Pd2+ speciation, the formation of palladium chloride complexes. We then proceed to discuss the formation of mixed-ligand complexes in general, and the formation of mixed species PdCli(OH)j and PtCli( 0H)j as special cases of interest. Subsequent to a comparison of theoretically derived and experimentally observed formation constants for the species PdC13(OH)‘-, we examine the most controversial aspect of Pd*+ aqueous chemistry, the formation of simple hydrolyzed species, Pd( OH)j .
[ PdCli] log
KiO
=
1%
[ pdcli_,
and the cumulative expressed as
I
-log[Cl-1,
formation
constants
lOgBiO=log~-ilog[Cl-]
(1 Ii14),
(1)
in Table 1 can be
(2)
where brackets denote the concentration of each chemical species. It is seen that the Ki, results of the five investigations shown in Table 1 are in generally good agreement. The results of DROLL et al. ( 1957), which are in poor agreement with the Table 1 results, have been excluded from the table for a variety of reasons: (a) the results reported by DROLL et al. ( 1957) are appropriate to zero ionic strength while the results shown in Table 1 are appropriate to ionic strengths between 0.8 molar and 1.O molar; inclusion of these results in previous tabulations has led to some confusion; (b) the DROLL et al. ( 1957) results are in error (in some cases by an order of magnitude) because they were obtained at high ionic strength (these data were not reported ) and extrapolated to zero ionic strength using the Debye-Hiickel equation (the Debye-Hiickel equation is an appropriate description of activity coefficient behavior only at very low ionic strengths); (c) the quality of the DROLL et al. ( 1957) formation constant analysis was compromised by impurities in their experimental solutions (WEED, 1964; WOOD, 1993 ) . Due to these problems, the use of thermodynamic data derived from the DROLL et al. ( 1957) results ( SASSANI and SHOCK, 1990) is inadvisable. II51
R. H. Byrne and L. R. Kump
1152
Table 1: Stoicbiometric Formation Constants for PI-Cl Species in Cl--C104*Mixtures
Burgerand etal. nyrssen (I%0 (1963) 1=0.82ooc I=l.O25% [H+l=0.6M [H+]=l.OM
Shchukamv
IO8K,
Biiukov Eldiing andShIen(1972) Average Average (ET) F$gkyi I=I.O25oC I=l.O25oC 108Kio lo8Pi0 [H$O.7M [H+]=l.OM [H+]=l.OM
IO8KIO
4.34
3.88
3.98
4.4
4.47
4.21fo.27
4.21
108Kzo
3.54
3.03
3.24
3.34
3.29
3.29i0.18
1.50
log K30
2.68
2.18
2.3
2.34
2.41
2.38M.19
9.88
‘og &a
1.68
1.34
(2.0)’
1.38
1.37
1.44i0.29
11.32
log b~%o
-0.80
-0.85
-0.74
-1.06
-1.18
-0.93Zto.19
I
log K4&30
-0.85
-0.86
108K3dKzo
I
-1.00
-0.84
-0.94
-0.88
-1.00
I
I
1 (-0.3)’
1
I
-0.96
-0.91&0.06 I
-1.04
I
-0.96m.09 1
a. These. data have been excluded from the averaged results
Examination of the summary results in Table 1 demonstrates that to a good approximation, at 25°C and 0.8 M 5 I I 1.0 M, log Kio/Kti-l)o = -0.93. Consequently, we can approximate the stoichiometric formation constants with the following empirical relation: log Kio = 4.21 - 0.93(i - 1).
(3)
This relation, which summarizes all of the formation constant data shown in Table 1, directly indicates that the Pd*+ complexes PdCl+, PdCl:, and PdCl; have similarly sized predominance fields on the order 0.93 log units wide. The log K40/K30 result obtained by TAUTet al. ( 199 1) at zero ionic strength (log K4o/K3o = - 1.3) is in fair agreement with the Table 1 summary of stability constant results. MIXED-LIGAND PALLADIUM COMPLEXES A number of works have examined Pd(I1) hydrolysis in chloride solutions. Equation 3 shows that for chloride concentrations greater than approximately 1O-‘.4 and ionic strengths near unity the dominant chloride complex of Pd( II) will be the species PdCl $- . Palladium hydrolysis at high chloride concentrations thus can be interpreted in terms of the stepwise replacement of chloride ions in this complex by OH _. The first step in a series of replacement reactions can be written as PdCl:- + OH- 2 PdC&OH*- + Cl-
(4)
where log
K3,
=
log
[PdC130H2-1 + log [PdCl:-]
[Cl-] -log
[OH-]. (5)
Assuming that the absorbance of hydrolyzed species at 290 nm is negligible (which is consistent with the modeling of KUMP and BYRNE, 1989), Eqn. 5 shows that when the solution absorbance of PdCl:- has fallen by one half due to the formation of PdC130H *- , log K3, can be directly calculated as log ([Cl-]/ [ OH -1). The results of KUMP and BYRNE ( 1989 ) indicate that in normal salinity seawater the concentrations of PdC&OH 2- and PdCl<- are approximately equal
when pH = 8.7. In this case, assuming log [H+][OH-] = -13.75 (BAESand MESMER,1976), it follows that, at 25°C and ionic strength = 0.7 M, log K,, = 4.8.
(6)
The PdCl$- absorbance vs. pH depictions of TAIT et al. ( 199 1) at log [Cl-] = -0.25 can be used to derive a log K3, result essentially equal to the result given above. However, TAIT et al. ( 1991) concluded that their results (supporting the importance of mixed-ligand species) should not be applied to studies of the natural behavior of Pd in seawater. They presumed that in their experiments, the association of Pd*+ and OH- depleted the free OH- concentration in solution to a point where no additional OH- was available to further replace Cl- associated with Pd*‘. This misinterpretation constitutes our major point of disagreement with TAIT et al. ( 1991). Since the formation of metal hydroxide complexes can be conceptualized in terms of hydrolysis equilibria, M*+ + nH20 = M(OH)f-’
+ nH+
(7)
and MCI:-i + jH20 = MQ_j(OH)f-’
+ jH+ + jCl-,
(8)
hydroxide can never be a limiting reagent. Thus, the Fig. 5 stability field diagram of TAIT et al. ( 199 1) which depicts a broad predominance field for mixed PdCli(OH)j species is not conditioned on high palladium/ hydroxide concentration ratios. These results of TAIT et al. ( 199 1) are appropriate to natural solutions. Our experimental result given in Eqn. 6 is in reasonable agreement with the results of KAZAKOVAand PTITSYN ( 1967) and MILIC and BUGARCIC ( 1984) as well as the results of TAIT et al. ( 199 1). In one molar NaCl at 25°C the result given by KAZAKOVAand PTITSYN (1967) is log K3, = 5.7. Recognizing that the symbol Pd*+ in the experimental work of MILIC and BUGARCIC( 1984) represents all nonhydrolyzed Pd species (including chlorides) under their experimental conditions (one molar NaCl at 25’C), their data can be used to obtain the result log K3, = 4.45. Thus, available experimental evidence indicates that 4.4 I log K3, 5 5.7. We prefer
1153
Discussion: Comment our Eqn. 6 estimate for log KS,, since it was obtained at relatively low total Pd concentrations (5 to 10 PM), minimizing potential problems attributable to the formation of polymeric Pd( II) complexes.
metal coordination sphere. This term is generally considered negligible for substitutions of equally charged ligands. We will consider the case i + j = 4, for both Pt(ll) and Pd( II) complexes. The statistically predicted Pt( II) results shown in Table 2 were obtained using the result log fl& Pt ) = 13.99 (ELDING, 1978) and a cumulative Pt(OH):- formation constant (log PO4= 30.4 1) derived from the stepwise PtCl:- hydrolysis constants of PESHCHEVITSKIet al. ( 1962). The statistically predicted Pd(l1) results shown in Table 2 were obtained using the Table 1 result log &(Pd) = 11.32 and the result log Po4(Pd) = 26.4 of NABIVANETSand KALABINA( 1970). The directly observed results in Table 2 were obtained using the above log pea results plus the following equations,
STATISTICALLY PREDICTED COMPLEXATION CONSTANTS The formation constants of mixed-l&and species for the metals Au, Tl, In, and Hg can be predicted from statistical arguments ( BAES and MESMER, 1976), and such predictions are generally in good agreement (within an order of magnitude) with direct observations. Statistical predictions for mixed-ligand species, MClI( OH)j are based upon the use of cumulative complexation constants, /3Ioand &j, ho
[MChl = [M2+l[cl_li
and
Poj =
[M(OH)jl
[M’+][OH-]I
(9)
log P3, = log 840 + 1% K3,;
log
P22 =
log
P3 I +
log
log
log
Pod =
1%
P13 +
lois KM,
013
=
lOis 022
+
log
K13;
with log BIo= C log Kio and log poj = c log Z&j.
(12)
i
where the constants Kij (Pt) are given as ( PESHCHEV~TSKI et al., 1962)
Formation constants for mixed-ligand Complexes can be calculated using the equation of BELEVANTSEV et al. ( 1982): log
Pij
=
&
1%
P(i+j)O
+
&
log
K3,tRj
PO(i+j)
K
(10)
-i(aij+lOg6ij)
22
tRj
= [~C130H2-I[C1-l [nCl$-][OH-] =
= 105’08;
PG(OW:-l[Cl-1 [PtC130H2-][OH-]
where log p-u = log
B-XOW:-l[Cl-1
[MC1i(oH)jl - i log [Cl-]
=
103.79.
’
Ko4(R)
= ~RCI~OH~~_lIOH_l
[WOW:-l[Cl-1
Table 2: Observed and Statistically Modeled Mixed-Ligand Formation Constants (O.SMB 5 l.OM, 2S“C) Observed
log Pij P40 @)
13.99
IO881 (PO
19.08
log p22 m
I
= 103.24
(13)
and log K3,( Pd) = 4.8 (this work, Eqn. 6). It is seen in Table 2 that the directly observed and statistically predicted stability constants are in generally good agreement. The statistically predicted constants are smaller than directly observed values by 0.2 to 0.4 log units, possibly due to ligand effects.
The term aij in Eqn. 10 accounts for the well-known promotion of mixed-l&id complex formation, relative to single ligand-type species, through statistical effects. In the case that i + j = 4, the term “II = -0.402 ( BELEVANTSEV et al., 1982 ). The term log 6Ij accounts for “ligand effects” which include steric hindrance, ligand-ligand bonds, and changes in the
23.38
log PI3PO
27.17
IO8PO4(n)
30.41
IO8P.IO (Pd)
11.32
1%PSI@‘dd)
16.12
PICdjCtd
18.72 I
23.04 26.96
15.69
log &a (Pd)
19.66
h? h 8d)
23.23
log PM W)
=
- j log [OH-].
(11)
I
= 104’30’
[PtC12(OH):-][OH-]
K’3(R)
Df*+l
1%
K22 ;
26.4
I
R. H. Byrne and L. R. Kump
1154 LINEAR FREE ENERGY RELATIONS
Linear free energy calculations provide an additional means of estimating log &(Pd). The comparison of PtX:- and PdX:- formation constants of HANCOCKet al. ( 1977) indicated that log &( Pd) = 0.8 log &( Pt). The &, results shown in Table 2 indicate that log &,( Pd) = 0.809 log &,(Pt), and the PO4results shown in Table 2 indicate that log &(Pd) = 0.868 log &(Pt), Using the average of our Table 2 &, and PO4results for Pt and Pd, log &( Pd) = 0.84 + 0.03 log &( Pt ), in addition to the directly observed P3, (Pt) value shown in Table 2, log &,(Pt) = 19.08, the cumulative PdCl,OH 2- fo~ation constant is calculated as log &, (Pd) = 16.0. This value, combined with the observed value for log ,&,( Pd) in Table 1, yields log X;, (Pd) = 4.7, a value that compares favorably to our experimental determination (KUMP and BYRNE, 1989; Eqn. 6). Thus, our log &(Pd) and log KS1(Pd) results are in good agreement with (a) the direct experimental observations of others ( TAIT et al., 199 1; MILK and BUGARCIC, 1984; KAZAKOVA and PTITSYN, 1967), (b) statistical predictions involving the &( Pd) and &f Pd) results of others, and (c) linear free energy relationships employing the &, (Pt) result obtained from the works of PESHCHEVITSKIet al. ( 1962 ) and ELDING( 1978). In contrast to the conclusion of TAIT et al. ( 199 1 ), mixed-ligand complexes PdCIi(OH)j should be significant species under natural conditions. KINETICS OF HYDROLYSIS WOOD ( 1991 ) indicated that if the hydrolysis results of NABIVANETSand KALABINA( 1970) are correct, Pd( II) hydrolysis should be much more extensive than the degree of hydrolysis observed by KUMP and BYRNE ( 1989). WOOD ( 199 1) suggested that equilibrium may not have been attained in the work of KUMP and BYRNE( 1989), and cited the works of KAZAKOVAand PTITSYN ( 1967) and TAIT et al. ( 199 1) as evidence that equilibrium with respect to hydrolysis is attained only after several days. The work of KAZAKOVAand PTITSYN( 1967) does not support this suggestion. KAZAKOVA and PTITSYN ( 1967) indicated that PdCl:- reacts very rapidly
with OH -, and modeled their titration data in terms of Eqn. 4 by assuming rapid replacement of Cl- by OH- followed by slow processes corresponding to polymerization. The data of KUMP and BYRNE ( 1989) were obtained through rapid titration in which the replacement of Cl- by OH- occurred in minutes to seconds followed by slower changes which we interpreted as polymerization. The work of TAIT et al. ( 199 I) did not produce evidence that the reaction shown in Eqn. 4 is a slow process. TAIT et al. ( 199 1) observed slow hydrolysis, as have all experimentalists investigating Pd*+ hydrolysis, but did not distinguish between the possibility of rapid monomeric reactions and slow polymeric reactions. TAIT et al. ( 1991) appear to assume that all Pd2+ hydrolysis reactions are slow based on the slow kinetics of Pt2+ hydrolysis (WV et al., 1990). It must be noted, however, that the kinetics of Pt*+ complexation is typically many orders of magnitude slower than the kinetics of Pd*+ complexation ( ELDING, 1973; BAESand MESMER, 1976). PURE HYDROLYSIS OF PALLADIUM Although investigation of Pd2+ hydrolysis in chloride solution is demanding, investigation of Pd2+ hydrolysis in noncomplexing media (e.g., perchlorates) appears to be even more challenging. Pd*+ has strong affinities for a wide variety of ligands, making concerns over solution impurities of extreme importance. Pd*+ has strong tendencies toward polymerization in perchlorate solutions over the range of pH in which the tmnsitions between Pd*‘, PdOH+, and Pd(OH): occur ( NABIVANETSand KALABINA,1970). Table 3 compares previous quantitative assessments of Pd’+ hydroxide complexation constants, &j( Pd) and &j( Pd) where
Poj(Pd)=
[Pd(OH)jl [pd2+l[oH_]j
lPd(OH)jl ~“‘(Pdf = [Pd(OH)j_~][OH-]
.
(14)
The stepwise stability constant behavior in Table 3 is very unusual. Whereas it is generally expected that the ratio of stepwise stability constants is much smaller than one (e.g.,
Table 3: Formation Constants for Palfadium Hydroxide Species ,
and
1155
Discussion: Comment log K20/K,O = -0.9, Table 1) the ratio of K&Ko, obtained by IZATT et al. ( 1967) is as large as five. According to the results shown in Table 3, addition of OH- to Pd2+ actually increases the affinity of palladium for OH- : PdOH+ has a larger affinity for OH - than does Pd’+ . The stepwise hydrolysis constant results of NABIVANETSand KALABINA ( 1970) are especially unusual in that Ko2/KoI > 1 and Ko, /I& < 10 -I’. Given the results which have been obtained for PdClI formation ( KzoJKlo _NK301Kzo N Km/ Kjo), and the absence of any explanation for the unusual stepwise stability constant behavior shown in Table 3, we feel that some of the data in Table 3 may be greatly in error. We suspect that the data of IZATT et al. ( 1967) are in error because the authors assumed negligible polynuclear species formation. IZATT et al. ( 1967) used total Pd concentrations on the order of one millimolar and greater, and equilibrated their solutions for five hours prior to each measurement. In contrast, NABIVANETSand KALABINA ( 1970) carefully accounted for the substantial influence of polymerization in their solutions. We suspect that the principal failing of the NABIVANETS and KALABINA ( 1970) experimental procedures was an inability to measure the shape of the Pd( II) solubility curve at total Pd( II) concentrations much lower than 1 X 10e5 molar. If the apparent solubility minimum observed by NABIVANETS and KALABINA( 1970) were somewhat lower than 1 X 10e6 molar, the equilibrium constants PO,and PO2would be smaller, and likely consistent with our observation of the dominance of PdCl:- in experimental seawater solutions even at pH = 8.3. It is interesting, and important, to note that a lower minimum in the Pd( II) solubility curve of NABIVANETSand KALABINA( 1970) would not affect their estimate for &: using the Fig. 1 solubility curve of NABIVANETSand KALABINA ( 1970), it is seen that total dissolved monomeric Pd concentrations are equal to 1 X 10m3molar at both pH = 0.5 and pH = 13.6. Because the free Pd*+ concentration in a solubility equilibrium with Pd( OH)201 decreases by a factor of 1026.2(i.e., 10ZApH)between pH = 0.5 and pH = 13.6, it follows that [Pd”] = 10e3- 10-26.2 at pH = 13.6 and [ Pd(OH):-] = 10e3. Noting that the [OH-] concentration
6
7
8
9
10
PH
FIG. 1. Predominance diagram for palladium chloride and hydroxide species in seawater (S = 35, 25”C, 1 bar) based on the formation constants recommended in the text. Numbers on graph indicate the i, j values for the species PdCli(OH )pi’.
Table 4: Proposed Hydrolysis Constants for Palladium (OJM sI 5 l.OM, ZS’C!)
log PO4
26.5
at pH 13.6 is approximately IO-‘.*, the formation constant log PO4is then calculated as log PO4= log ‘p;E;;‘-l
- 2 log [OH-]
= 26.2 + 0.4 = 26.6.
(15)
This result is in good agreement with the reported result of NABIVANETSand KALABINA( 1970). We suspect this result may be the only truly quantitative formation constant which can be directly derived from their work. A final extension of this analysis of the hydrolysis of Pd can be made by assuming that the ratios of Kocj+l,/Koj for stepwise Pd2+ complexation are similar for OH- and Cl-. If Kocj+l,fKoj = 0.1, then /304= K~,(O.1K,,)(O.O1Ko,)(O.OO1Ko,)
= 1026.5
where 1026.5is an average of our estimate and the estimate OfNABIVANETSand KALABINA( 1970). The cumulative OHcomplexation constants, thus obtained, are shown in Table 4. It is interesting to note that the recent results of WOOD ( 199 1) shown in Table 3 are in much better agreement with these results than the Bo3and PO2results of NABIVANETSand KALABINA( 1970). A MODEL
FOR THE SPECIATION OF PALLADIUM IN SEAWATER
Through the use of a combination of experimental results, linear free energy relationships and statistical modeling, we have developed an internally consistent set of formation constants for PdCli(OH)j species in solutions at intermediate ionic strength. Recommended log fiio results are found in Table 1. Our recommended log p3I (Pd) result is the directly observed value in Table 2, log &, (Pd) = 16.1. Noting that the statistically predicted log 83, (Pt ) and log ,822(Pt ) results underestimate observed results by about 0.35 log units (Table 2), our recommended p22and /?I3values for Pd are log f12*(Pd) = 20.0 and log &(Pd) = 23.6. Suggested formation constants for Pd( 0H)j are shown in Table 4. Figure 1 shows the predicted speciation of Pd(l1) in seawater using our recommended constants. The species PdCl$- dominates throughout the normal range of seawater pH, 7.4 I pH 5 8.4. The next most abundant species are the three-chloride complexes PdC&OH 2- and PdCl 5 . Mixedligand complexes, PdC&OH 2- and PdC120H :-, become dominant above pH 8.7. Interpretation of palladium hydrolysis experiments is presently controversial. Indeed, in our earlier work ( KUMP and BYRNE, 1989), we were careful to report our experimental data in terms of the observed extent of hydrolysis
R. H. Byrne and L. R. Kump
1156
without specifying the relative importance of single-ligandtype hydrolyzed species (Pd( 0H)j) vs. mixed-ligand hydrolyzed species (PdCli( OH)j). Herein we have extended our analyses, by using statistical and linear free energy calculations, to elucidate the potential importance of mixed-ligand species. Clearly, these arguments are subject to considerable uncertainty and cannot bc substituted for direct observations. However, even order of magnitude uncertainties are dwarfed by the range of hydrolysis constants which have recently been reported for the species Pd(OH)i and Pd(OH); (Table 3). Our hydrolysis observations in seawater ( KUMP and BYRNE, 1989), plus the arguments advanced in the present work, indicate that PdC&OH ‘- is the dominant, hydrolyzed form of Pd within the normal pH range of seawater. Nevertheless, we know of no direct experimental observations which preclude the possibility that we have underestimated the significance of Pd(OH); due to very slow formation kinetics for this species. Although we feel there is little reason to suspect that the kinetics of Pd( OH)! formation is substantially slower than the formation kinetics of PdC130H2- or PdC12(OH):-, we are currently initiating long-term experiments to investigate this possibility.
CONCLUSION In this paper we have quantitatively summarized previous Pd( II) chloride complexation results and quantitatively compared these results with the results of TAIT et al. ( 199 1). We have produced recommended formation constants for the species PdC130H2- and other mixed species, and have shown that our recommended PdCljOH *- data are in agreement with (a) direct experimental evidence including the work of TAIT et al. ( 199 1 ), (b) statistical predictions based on the well-known characteristics of mixed-ligand complexation, and (c) linear free energy relationships using experimentally observed Pt( II) complexation characteristics. Finally, we have produced recommended Pd( OH)j hydrolysis constants which are consistent with Pt( OH):- hydrolysis data and which are in better agreement with the recent data of WOOD ( 199 I ) than the data of NABIVANETSand KALABINA (1970). We feel that our narrative produces a coherent picture of Pd( II) complexation and provides a framework within which future experimental investigations of both Pd( II) and Pt( II) complexation can be assessed.
Acknowledgments-This work has been partially supported hy grants from the Marine Chemistry and Geology and Paleontology Programs of the National Science Foundation. We express our gratitude to Drew Tait, Scott Wood, and an anonymous reviewer for their thoughtful critiques of the manuscript.
Editorial handling: G. Faure
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