Rare earth element complexation by PO43− ions in aqueous solution

Geochimico et Cosmochimica Copyright 0 1991 Pqamon

0016-7037/91/$3.00 + .oO

Acfa Vol. 55, pp. 2729-2735 Press pk. Printed in U.S.A.

Rare earth element complexation by PO:- ions in aqueous solution ROBERTH. BYRNE,JONG HYEON LEE, and LINDA S. BINGLER Department of Marine Science, University of South Florida, St. Petersburg, FL 33701, USA (Received December 2 1, 1990; accepted in revised form August 1, 199 1)

Abstract-Complexation of trivalent rare earths by PO:- ions has been assessed at t = 25 “C by examining the influence of aqueous phosphate concentrations on the distributions of ‘“Ce and 15’Gd between 0.68 molar NaClO., and tributyl phosphate (TBP). Estimated Ce3+ and Gd3+ phosphate complexation constants appropriate to zero ionic strength are Ce3+ + PO:- c* CePOO;

log LB? 5 11.7

Gd3+ + PO:- c* GdPO:;

log ,_/I? = 12.2.

Our estimates of these formation constants at zero ionic strength are approximately seven to eight orders of magnitude lower than previously reported estimates for lanthanide and actinide PO:- complexation. Linear free energy relationships relating the complexation of Ca’+ and lanthanides ( Ln3+) by a variety of organic ligands, in conjunction with previous direct observations of CaPO; formation, are consistent with the experimentally derived CePO! and GdPO! formation constants reported in this work. Gadolinium speciation calculations indicate that the PO:- ion can effectively compete with the CO:- ion for free Gd3+ in model groundwater at pH 7-9. Rare earth element phosphate complexation is a significant process in natural freshwater systems which are neutral to mildly basic when the concentration ratio [ HPO$-]/[ HCO;] is greater than approximately 1 X 10W3. addition to reaction ( 1) other significant equilibria will inevitably include

INTRODUCTION PREVIOUS ACCOUNTS OF trivalent lanthanide and actinide complexation by phosphate (PO:-) ions have been sharply conflicting. Although reported stability constants for trivalent lanthanide and actinide complexation

M3+ + HPO:- ++ MHPO:; &I, = [MHPO:][M3+]-‘[HPO:-I-’

(2)

M3+ + H2PO; c-) MHzPO:+; M3+ + PO:- c* MPO:; & = [MPO:][M”]-‘[PO:-]-’

n&,

= [MH2PO:+][M3+]-‘[HzPO;]-‘.

(3)

(1) The consequence of omitting equilibria (2) and (3) in the interpretation of experiments conducted at low pH can be quite severe. Lanthanide complexation experiments interpreted solely in terms of equilibrium ( 1) at low pH produce pH dependent & estimates (BYRNE and BINGLER, 1989). In contrast, when the same experiments are interpreted in terms of equilibrium ( 3 ) , excellent agreement is obtained for the constant &3i over a wide range of solution acidities (BYRNEand BINGLER, 1989). These observations lead to the conclusion that previous examinations of lanthanide and actinide complexation by phosphate ions should be interpreted solely in terms of complexation by H2PO; ions and HPO$- ions. There presently exists no direct measurement of lanthanide and actinide complexation by PO:-. In the present work, we have assessed the complexation of trivalent lanthanides by PO:- ions using two distinctly different strategies. (a) In the first case we have examined CePOi formation in solution at low total phosphate concentration (B(PO:-), 5 3 X 10m4 molar) and high pH (pH 5 7). The extent of CeH,PO:+ and CeHPO: formation in these experiments should be greatly reduced relative to the study of MAYER and SCHWARTZ ( 1950) wherein

have ranged between approximately 10 ” and 10” (MAYER and SCHWARTZ, 1950; MOSKVIN, 197 1), alternative assessments of MPO: formation (BYRNEand BINGLER,1989) have indicated that, due to improper modeling of experimental data, previous estimated 4, values should be many orders of magnitude smaller. This controversy is important in a geochemical sense because 4i values as large as 10 I9indicate that PO:- should substantially dominate trivalent lanthanide and actinide complexation by other inorganic ligands in seawater. It is presently thought that CO:- is the dominant inorganic lanthanide complexing ligand in seawater as well as in most fresh waters. Since phosphate concentrations in seawater are highly variable relative to carbonate, lanthanide and actinide complexation models based on PO:- as a dominant ligand would differ markedly from current carbonatedominated models. Metal complexation investigations are typically performed under pH conditions much lower than those characteristic of natural water. While investigations under relatively acidic conditions reduce adsorption and hydrolysis, for a multiprotic acid such as H3P04, experimentalists must be aware that in 2729

R. H. Byrne, J. H. Lee, and L. S. Bingler

2730

x(PO:-)r 5 0.02 modal and pH = 1.8 + 0.2. This work should directly demonstrate whether or not r/3, is so large that CePOq formation greatly dominates CeH2POF even at pH= 1.8(where[H,PO;]/[PO:-]a 10i4).(b)Inasecond type of study we investigated Gd( III) complexation by phosphate ions over a wide range of pH and total phosphate concentrations. In this investigation three experiments were added to the four data sets used by BINGLER and BYRNE (1989) to examine Gd(II1) complexation by HzPO; and HPO$-. Within the range of conditions (4.0 I pH I 6.3; 2 X 10e4 M 5 Z(PO:-)7 I 4 X 10m3 M) employed in this work, the data indicated that in addition to MH*PO:+, MHPO: , and MPO: it was necessary to consider formation of the species M( HPO,);:

M3+ + 2HPO:-

c* M(HP0,);;

u,& = [M(HP0,):][M3+]-‘[HPO:-I-‘.

(4)

Our previous investigation of Gd3+ complexation by H2PO; and HPO:- indicated that the GdH*PO:+ and GdHPO: species are likely to be of minor importance in natural aqueous systems even at total phosphate concentrations on the order of 1 X 10m4M. In the present work examinations of CePOt and GdPOi formation (Eqn. 1) are used to assess the possible importance of lanthanide and actinide complexation by PO:- ions in natural systems.

Coming model 130 pH meter. A 3 M NaCl solution was used as the electrode filling solution. Distribution coefficients were determined by measuring the net counts (counts above background) of radionuclides in each phase. Constant geometry was maintained for each sample. Gamma spectroscopy analyses were performed with a Bicron 2 X 2 inch well type NaI detector ( 2MW20 / 2 ) and a Tracer Northern multichannel analyzer(TN-1710). -’ ’ Our Ce distribution coefficient data were analyzed using the equation Do/D = 1 + &,[PO:-1,

DoID = 1 +

H2Lb;

W&‘&I, + HLB’I]HPO:-I, + d;NQ-l+

+ Is’,W-IT

(6)

where H*LIBi= [GdH2PO:+][Gd3+]-‘[HzPO;];‘, Ht.,%= [GdHPO:][Gd3+]-‘[HPO:-I,‘,

MATERIALS AND METHODS The procedures used to examine the influence of PO:- ions on the distributions of Ce3+ and Gd3+ between 0.68 M NaC104 and tributyl phosphate (TBP) closely follow the procedures used by BINGLER and BYRNE (1989) to investigate Gd complexation by HPO:- ions. Stock concentrated solutions of NaClO, (Aldrich Chemical Co., anhydrous 99+%) were kept at pH 9.0 for approximately 24 h. Before dilution to 0.68 M, the solutions were filtered with a 0.45 r.~polycarbonate filter (Nuclepore Corp.) which had been acid cleaned and rinsed thoroughly with deionized water. Approximately 125 mL of TBP (Fluka, puriss pa) was shaken vigorously with equal volume of 1 N NaOH (Baker, CO* free) in a funnel and separated for 5 min. The excess NaOH solution was discarded and the TBP was transferred into five 25-mL mastic tubes for centrifugation at 12,000 rpm for 5 min. Equal volumes of TBP and 0.68 M NaC104 were added to a 250 mL jacketed beaker maintained at the experimental temperature (25 + O.OS”C). The two phases were vigorously mixed for 5 min. After 5 min separation, the NaClO, phase was carefully removed and the process was repeated. Subsequent to these procedures, 100 mL of fresh 0.68 M NaClO, solution was combined with the preequilibrated TBP in the jacketed beaker. A solution consisting of 0.03 M NazHP04 (Fisher Scientific) in 0.68 M NaC104 was added to the jacketed beaker to attain the experimental total phosphate concentrations (Z( PO:-)r 5 3 X 10e4M (Ce experiments) and2 X 10m4M i Z(PO]-),< 4 X 10-3M(Gdexneriments). l”Ce or 15’Gd (Isotope Products Laboratories) were next added to the organic-aqueous solutions, and the mixture was subsequently acidified with HClO4 to pH x 4.0. The total concentrations of ‘“Ce and lS3Gd used in this work were approximately 1.4 X lO-9 M and 1.3 X lO-‘o M, respectively. Phase contact and equilibration were accomplished by stirring and bubbling for 20 min with Nz (99.99%, Air Products). Solution pH in these experiments was varied through addition of 0.01 M Na2B40, - 10 Hz0 (J. T. Baker) which hadbeen dissolved in 0.68 M NaC104. Following eouilibration. 20 to 30 min were allocated for phase separation, final pH was measured, and duplicate 1 mL samples of each phase were taken for gamma spectroscopy analysis. Measurements of pH on the free hydrogen ion concentration scale (MCBRYDE, 1969, 1971; BYRNEand KESTER, 1978) were obtained using a Ross combination pH electrode (Orion No. 8 10200) and a

(5)

where [PO:-], = [PO:-] + [ NaPO:-1, D is a “@Ceactivity ratio (organic phase/aqueous phase), Do corresponds to a distribution coefficient at zero total phosphate concentration, and #, is the CePOq formation constant at t = 25°C and 2 = 0.68 M (#, = [CePO9][Ce”]-‘[PO:-];‘). The procedures used to examine Gd3+ complexation by phosphate ions are identical to the procedures used by BINGLERand BYRNE (1989) in investigations of Gd3+ complexation by HzPO; and HPO:-. Our Gd distribution coefficient data were analyzed using the equation

HL& = [Gd(HP0,);][Gd3+]-‘[HPO:-l-,2, 4;

= [GdPO:][Gd3+]~‘[PO:-I.‘,

(7)

D-W’WT = [ H,PO;] + [ NaH,PO!], [HPO:-1, = [HPO:-] + [NaHPO;], [PO:-],

= [PO:-] + [NaPO:-1,

(8)

and [ ] denotes free concentration of each chemical species. The concentrations [H2P0.&, [ HPO:-]r, and [PO:-], in these experiments were determined using the equations Z(PO-),

= [H,PO,] + [HzPO;]T + = [PO:-],(

[HF’O:-IT+ [PO:-1,

1 + K;[H+] + K;K;[H+]* + K;K;K;[H+]‘)

(9)

with K; = [HPO:-],/[H+][PO:-I,, K; =

[H,PO;I,/[H+l[HPO:-I,,

Ki = WS~~I/W+~FMQ;IT,

(10)

and log K; = 11.19 (ATLASet al., 1976) log K; = 6.38 log K',=

(BINGLERand BYRNE, 1989)

1.72 (ATLASet al., 1976).

(11)

Analysis ofthe titration data (D,[H+], Z( PO:-),) was accomplished using nonlinear least-squares fitting techniques (CANTRELLand BYRNE, 1987). Ce3+ complexation analysis minimized the residual sum of squares (RSS) function

C 11 - Dil#‘(l ,

+ 4’,[PO:-l,)}2,

2731

Complexation of REEs by phosphate ions 0.1

and Gd’+ complexation analysis entailed minimization of the RSS function

o.o-

@q 0

-0.1

-

-0.2

-

-0.3

-

-0.4

-

-0.5

-

o-d

*o

og.o

ob bb 0

. n

+ &;[HPO:-I;+

LJ~‘,[P~:-IT))*~ (13)

where organic/aqueous distribution coefficients at zero total phosphate (Dp ) were constrained to be constant for each of the titration

&n&in-a single experiment. The conditional formation constants (&Y, &‘, and &“) expressed in terms of total phosphate ion concen&ions at 25°C in 0.68 M NaClO, were converted to constants expressed in terms of free ion concentrations (Eqns. l-4) using the ion pairing model of MILLERO and SCHREIBER( 1982). Estimation of complexation constants appropriate to ionic strengths rather than 0.68 M requires an assessment of the activity coefficient behavior of complex ions. Whereas for simple free ions, such as Ce”+, Gd3+, PO:-, HPO-, etc., the calculational procedures of MILLEROand WHREIBER( 1982) were used directly, for complex ions, such as GdPOy, GdHPO;, etc., activity coefficients were calculated assuming that y( MPO:) = 1 at all ionic strengths and that NOT and Ca’+ serve as reasonable analogs for other complex species: y(N0;)

= y(MCO:)

= y(MHPO:)

= -/(M(CO,);)

= r(M(HPO4);)

r(Ca*‘) = -y(MH#Op).

(14)

One may note that the choice of NOT as a model ion for MCO; and M(COa); activity coefficient behavior is consistent with the choice r(MCOf) = r(M(CO,)r) = 0.5 made by CANTRELLand BYRNE ( 1987) at 0.68 M ionic strength. RESULTS

AND DISCUSSION

The CePOi complexation constants, &, obtained in five experiments are shown in Table 1. Combined Ce 3+data sets, obtained at nearly constant total phosphate ( 3 X 10m4M), are shown in Fig. 1. Gd complexation results obtained in seven experiments at 25’C and 0.68 molar ionic strength are given below:

B P M 0

-0.6

-

-0.7

-

0

Experiment

#l

*

Experiment

#2

.

Experiment

#3

0

Experiment

#4

b

Experiment

#5

.

I.

-0.8 4.0

4.5

A.

5.0

. .

I.

I

5.5

6.0

“4;

= (7.7 f 1.0) x 106

4;

= (6.1 + 0.4) x 108.

(15)

The least well-defined complexation constant obtained in the Gd analyses (&‘, = 43 ? 7) has a relative uncertainty of

Table 1. Conditional CeP04o comconstants obtained at plexation 25oC and 0.68 M ionic strength.

pH range 4.66-6.35 4.44-6.50 4.47-7.00 4.39-6.41 4.47-6.66

Do

LPI’

9.23 9.29 9.55 9.39 9.68

(1.67fO.l4)x108 (1.79fO.l7)x108 (2.56f0.15)~108 (1.75*0.08)x10* (2.83kO.22)xlOS

average:

(2.12f0.54)~108

I. 7.0

7.5

FIG. 1. Ce( III) distribution data (D/Do) as a function of pH for five experiments (total PO, = Z(PO:-)r 5 3 X 10m4M).

+16%. Since superfluous terms in the complexation model should exhibit large relative uncertainties, the observation that all of the Gd complexation constants are relatively well defined indicates that all of the complexation constants given in Eqn. ( 15) are required for a complete description of the data. In the case of Ce complexation, we attempted to ascertain whether additional terms were warranted in the model ( Eqn. 5 ) . Since our experimental procedures produced higher [ CePOy ]/ [ Ce3+] ratios than were obtained for [ GdPOi]/ [ Gd3+] in our analyses, we investigated the possibility that an estimate of the constant 4; = [Ce( PO,):-] [ Gd3+]-’ [PO:-] ?2 could be extracted from the distribution data. Simultaneous analysis of the five combined Ce complexation experiments, which included the possibility of both CePOt and Ce( P04):- formation, yielded a ,j3’, value very similar to the average result shown in Table 1: J;(Ce3+)

&(Ce3’)

= (7.38 f 0.32) X lo3

6.5

PH

= (2.0 f 0.3) X 108.

The result obtained for Ce( PO,):&‘,

I.

(16)

formation is

= (4.6 + 2.4) X 10t5.

(17)

This result is rather poorly defined relative to the other formation constant results. The lower 95% confidence limit for & in the analysis was zero. Our 4; result is probably most useful for its upper 95% confidence limit, log L/355 16. The upper and lower 95% confidence limits obtained for 4’,(Ce3’) in our analysis indicated that 8.13 I log ,&I’, I 8.42. Calculated phosphate complexation results appropriate to 25°C and zero strength are shown in Table 2, column 4. Formation constant estimates at zero ionic strength are approximately seven to eight orders of magnitude lower than previous 41 estimates for lanthanide or actinide ions. Direct observations thereby support the contention (BYRNE and BINGLER, 1989) that previous examinations of Ce3+ complexation (MAYER and SCHWARTZ, 1950) by PO:- were improperly modeled. Investigations of metal-phosphate complex formation, conducted at low pH, must account for the pres-

ence of species such as MHPO: and MH2PO:+. In the apparent absence of any useful previous work on trivalent metal complexation by PO:- ions, it is helpful to consider our results in light of previous examinations of Ca2+

2732

R. H. Byrne, J. H. Lee, and L. S. Bingler Table 2. Ce3+ and Gd3+ phosphate complexation constants (t=25oC ) expressed in terms of total ligand concentrations (column 2, 1=0.68 M), free concentrations (column 3, 1=0.68 M), and activities (column 4, I=0 M).

species CePO$ GdPO&’ GdHP04+ Gd(HP04)2GdH2PO$+

QLP'

@LP

l%TLP

8.33 8.19 3.87 6.89 1.63

8.84 9.30 4.11 7.38 1.71

11.73 12.19 5.91 9.91 2.14

complexation by PO:-. Following the expectation that lanthanide and calcium ions should exhibit coherent linear free energy relationships, we have compiled all stability constants in Critical Stability Constants, Vol. 1-6 (MARTELL and SMITH, 1974, 1977,1982; SMITHSONMARTELL, 1975,1976, 1989 ) , which depict Ce 3+, Gd 3+, and Ca2+ complexation by organic ligands under generally comparable conditions (t = 20-25”C, Z = 0.1 M). Data for Ce3+, Gd3+, and Ca2+ complexation by inorganic ligands under comparable conditions are comparatively scarce. The correlation diagrams obtained by plotting Ce 3+and Gd 3+stability constants against Ca2+ stability constants (t = 20-25”C, Z = 0.1 M) are shown in Fig. 2. Through least-squares regressions of the data shown in Fig. 2, we obtained the following relationships:

energy) estimate and indicates that at 25°C and zero ionic strength log &y( Gd3+) g 12.2. The somewhat poorer agreement between direct vs. linear free energy based estimates in the case of CePOt formation suggests that, due to the neglect of species such as CeHPOa, the direct assessments of &(Ce3’) should be regarded as upper bound values. Due, apparently, to the minor role of CeHPO: formation at the low total phosphate concentrations employed in cerium complexation experiments, we were unable to obtain welldefined estimates for &I ( Ce3’) in this work. Had we been able to account for both CePOy and CeHPO: formation in our experiments, we expect that our resulting log &( Ce3+) estimate would be somewhat less than 11.7. Thus, direct measurements should be taken as indicating that log fiy(Ce’+) I 11.7. The formation constant results shown in Tables 2 and 3 are dependent on the protonation constants K; , K2, and K; chosen to describe phosphate speciation in our 0.68 M perchlorate solutions. The selected result log K2 = 6.38 is directly appropriate to the experimental solutions since it was directly determined in 0.68 M NaC104. The influence of K’, on our calculations is quite small. Since pH > 4 in our analyses, H3P04 accounted for no more than 0.5% of the total phosphate in our experimental solutions. In contrast, the results obtained for L/3’,are directly proportional to the magnitude of the constant (K;) used to calculate relative

log LPI(Ce3’) = (2.71 f 0.37) + (1.29 + 0.05) log&(Ca2+) log J,(Gd3+)

(18)

= (2.87 t 0.37) + (1.42 + 0.05) log &,(Ca2’).

(19)

Because of the dominantly electrostatic nature of alkaline earth and rare earth complexation (MOELLER, 1963), it was expected that these equations, describing complexation by organic ligands, could be used to relate Ce3+, Gd3+, and Ca2+ complexion by inorganic ligands. The CaPO; formation constant at 25°C and zero ionic strength in the MILLERO and SCHREIBER ( 1982) recalculation ofthe complexation results of CHUGHTAI et al. (1968) is log &(Ca2’) = 6.46. Employing the ion pairing model of MILLER0 and SCHREIBER (1982) with y(CaP0;) = 0.782, log J1(Ca2’) at 25°C and Z = 0.1 M is then calculated as log JI(Ca2+)

= 5.44.

-I 0

2

4

6

8

10

12

14

loizLP,(Ca*+) 25

*

3,

8.

I.

I.

I.

I.

. 20

+-

(20)

Combining Eqns. ( 18) through (20), we obtained the estimated CePO: and GdPOt complexation constants shown in Table 3. Comparison of the Ce3+ and Gd3+ complexation results shown in column 4 of Table 2 and column 3 of Table 3 indicates that our direct observations of CePOz and GdPOi formation are generally consistent with expectations based on comparative Ca2+ vs. Ln3+ complexation results and previous direct observation of CaPO; formation. In the case of GdPO$’ formation, our direct formation constant estimate is in excellent agreement with our indirect (linear free

0

0

2

4

6

8

IO

12

14

logLP,(Ca'+)

FIG. 2. Linear energy relationships (Ca’+ vs. Ce3+and Ca’+ vs. Cd”) for a variety of organic ligands at I = 0.1 M and t = 20-25°C. Open circles and filled circles, respectively, represent aminocarboxylic acids and iminodiacetic acids. Open squares represent mono and dicarboxylic acids. Other organic ligands, represented by crosses, include phenols, pyridinecarboxylic acids and carbonyl ligands.

Complexation

Table 3.

CePO$

and GdPOda complexation

constants obtained energy relationships

Species

of

from the linear free depicted in Fig. 2.

logLpI(I= 0.1 M)

logLPIU= 0 M)

CaP04-

5.44

6.46

CeP04O

9.73

11.33

GdP040

10.59

12.19

PO:- and HPO:- concentrations. As such, speciation calculations employing our L@‘rresults should be performed in a manner consistent with the choice for K; in 0.68 M NaC104. Alternatively, the dependence of Ce3+ and Gd3+ phosphate speciation calculations on K3 can be eliminated by noting that the equilibrium constant for the reaction

M3+ + HPO:- * MPO: + H+

(21)

can be written as ,V,, = &,/K; = [MPO:][H+]/[M’+][HPO:-I,.

(22)

The conditional constants calculated for Ce3+ and Gd3+ in 0.68 M NaC104 at 25’C are log P’,,(Ce3’) = -2.86

(23)

log /3;,(Gd3+) = -2.40.

(24)

The GdPOi complexation results (Table 2) can be used to assess the relative importance of REE phosphate and carbonate complexation in natural waters. Total phosphate concentrations in seawater are generally confined to values less than or equal to 3 micromolar. Approximately one-third of the total phosphate in seawater is present as free HPO:( MILLERO and SCHREIBER,1982). If log K; (Eqn. 11) is on the order of 11.19, it follows that the free PO:- concentration in seawater (pH < 8.3) is less than or approximately equal to 1 nanomolar. The result in Table 2, log Lpl(Gd3’) = 9.30 at I = 0.68, then indicates that for normal seawater [GdPO!] I 2[Gd3’]. Since carbonate complexes account for approximately 90% or more of the total inorganically complexed Gd( III) in seawater (BYRNE et al., 1988), it appears that phosphate complexation is a relatively minor feature of REE complexation in seawater. Although REE complexation by PO:- ions appears to be of secondary importance in seawater, this is not necessarily the case for REEs present in groundwaters. In Fig. 3, we have constructed speciation schemes for Gd(II1) employing the mode1 groundwater of WOOD ( 1990) wherein: ZCl- = 2 X lO-4 M ZFZO:-

ZSO:- = 1 X 1O-4M

= 1 X 1O-6 M

ZCO:-

= 1 X 1O-4 M

= 1 X 1O-6 M

ZNOS

= 1 X 1O-4 M.

The speciation scheme shown in Fig. 3 was constructed using the formation constants and dissociation constants shown in Table 4 corrected to a model groundwater ionic strength equal to 0.002 M. The speciation results shown in Fig. 3 indicate that MPO: complexes are likely to be an important aspect

2133

REEs by phosphate ions

of lanthanide and trivalent actinide speciation in some groundwaters. In examining Fig. 3, wherein GdPOt is a dominant species between pH 7 and pH 9, it should be noted that carbonate and phosphate concentrations in groundwaters are highly variable. For groundwaters in contact with carbonate-rich substrates, total carbonate concentrations can be approximately an order of magnitude higher than that employed in our model calculation. Phosphate concentrations in polluted groundwater systems ( WEDEPOHL,1978; WOOD, 1990) and porewaters ( JONASSONet al., 1985) can be as much as a factor of one hundred times greater than the concentration assumed in our calculations. Thus, within the range of conditions found in natural systems it is possible for lanthanide complexation to be dominated by either phosphate or carbonate. It should also be noted that since the formation constants for phosphate complexes, like carbonate complexes, are likely to be strongly dependent on atomic number, the relative importance of phosphate and carbonate complexation in natural waters is likely to vary across the 15 member series of elements. Due to the substantial variability of carbonate and phosphate ion concentrations in natural waters, it is difficult to formulate general rules with respect to the relative importance of rare earth phosphate and carbonate complexation. As a reasonably straightforward means of addressing this problem, it can be noted that an equilibrium constant, K, for the reaction, GdCO : + HPO:- ++ GdPO y + HCO ; , at zero ionic strength can be calculated as

K= [GdWiI[HCOiI

= Is?GWCOI)

[GdCO:][HPO:-] 10

12.19

x

&Kj(HPO:-)

10’0.33 =

=

107.82

x

102.35

(24)

1012.35

where ,& = [GdPOi][Gd3+]-‘[PO:-]-‘, c/3: = [GdCOf] X [Gd’+]-‘[CO:-]-‘, e(HC0;) = [HCO;][H+]-’ and e(HPO:-) = [HPO$-][H+]-’ x [co:-]-‘, X [PO:-]-I. It can then be noted that in mildly acidic to mildly alkaline solution, where free HCO; and HPO:- ions are dominant forms of carbonate and phosphate, REE phosphate complexation will be significant compared to carbonate complexation when the ratio ZPO-/ZCO:is approxi-

80

60 s

50

7 5

40

B

30

* 20

0 4

5

6

7

8

9

IO

11

PH FIG. 3. Gadolinium speciation calculated for the model groundwater of WOOD ( 1990) as a function of pH. Only species whose fmction of total Gd is larger than 5% are explicitly shown.

2734

R. H. Byrne, J. H. Lee, and L. S. BingIer Table 4. shown

for

phosphoric

Equilibrium

Gd complexation constants (25oC and zero ionic strength) are a variety acid

of

dissociation

inorganic constants

ligands,

along

with

carbonic

at 25oC and zero ionic

and

strength.

Source

Relationship

log[GdF*+]-log[Gds+]-log[F-]=4.10 log[GdFs+]-log[Gds+]-21og[F-I=662

Wood (1990)’ Lee and Byrne (1991a) Lee and Byrne (1991a)

log[GdCl*+]-log[Gd3+]-log[CI-]=0.33

log[GdS04+]-log[Gd3+]-log[SO$-]=3.66

Wood

log[Gd(SO&-]-log[Gd3+]-21og[SO&]=5.20

Wood

log[GdOH*+]-log[Gd3+]+log[H+]=-8.00

Baes

(1990)* (1990)2

Mesmer (1976) Baes and Mesmer (1976)

log[Gd(OH)a+]-log[Gds+]+21og[H+]=-16.42 log[GdNOsa+]-log[Gds+]-log[NOs-]=1.23 log[GdHCOsa+]-log[Gds+]-log[HCOs-]=2.57 log[GdCOs+]-log[Gds+]-log[COs*-]=7.82 log[Gd(COs)a-I-log[Gds+]-2logICOs*-]=13.33 log[GdHsPO$+]-log[Gds+]-log[HsPOd-]=2.74 log[GdHP04+]-log[Gds+]-log[HP0&=5.91

and

Wood (1990)3 Lee and Byrne (1991b) Lee and Byrne (1991b) Lee and Byrne (1991b) This work This work This work

log[Gd(HPO&-]-log[Gds+]-2log[HPO.?-]=9.97 logIGdPO&‘]-log[Gds+]-log[PO$-]=12.19

This work

log[H+]+log[HCOs-I-log[HaCOs]=-6.35 log[H+]+log[COsa-I-log[HCOs-I=-10.33 log[H+]+log[H2PO&I-log[HsP04]=-2.15 log[H+]+log[HP042-]-log[H2P04-I=-7.20

Smith Smith Smith Smith Smith

log[H+]+log[PO$-I-log[HPO&]=-12.35

and and and and and

Martell Martell Martell Martell Martell

(1976) (1976) (1976) (1976) (1976)

1. Derived from Mironov et al. (1982). 2. Derived from Powell (1974). 3. Derived from Eu data of Choppin and Strazik (1965).

mately 1 X 10m3 or greater. This condition seems likely to be satisfied in many natural freshwater systems. SUMMARY Direct measurements of CePO! and GdPOi complexation constants indicate that REE and trivalent actinide phosphate complexes (MPO:) are minor species in the oceanic water column. Our MPO$’ complexation results, applied to model freshwater systems, indicate that REE complexation by PO:- ions is likely to be an important process in many freshwater systems. As such, the results of this work in conjunction with previous investigations ( CANTRELL and BYRNE, 1987 ) indicate that models of REE and trivalent actinide behavior in the environment should consider pH, carbonate ion concentration, and phosphate ion concentration as key solution variables. Acknowledgments-This

work was supported by grant No. OCE8400548 from the National Science Foundation. The authors wish to express their gratitude to Gregory R. Choppin, Scott A. Wood, and an anonymous reviewer for their constructive criticism of this work. Editorial handling: F. J. Miller0

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