P R E P R I N T – ICPWS XV Berlin, September 8–11, 2008

Introduction

Soluble zinc is present in the feedwater of most US domestic BWRs, either from the corrosion of brass condenser tubes and the incomplete removal by filter/demineralizer condensate cleanup systems, or as a consequence of the injection of depleted ZnO (DZO, <1% 64Zn) for shutdown radiation field control. Zinc oxide may be added to PWR reactor coolant to retard PWSCC (primary water stress corrosion cracking) of system materials as well as to control radioactivity buildup. In BWRs, the presence of zinc is related to an increase in the quantity of tenacious crud on fuel cladding, due to the formation of zinc ferrite. In some cases, deposits of zinc silicate have been found within the pores or steam channels of fuel-crud deposits. The potential exists for accelerated corrosion of the underlying cladding if tenacious fuel-crud deposits become consolidated by the precipitation of zinc silicate. This work was undertaken to predict the amount and conditions under which these deposits are formed and to increase our understanding of the mechanisms by which these deposits are formed. The information generated on zinc silicate solubility may also be applicable to PWR operating conditions. The experiment plan involved: (1) the preparation and characterization of crystalline particulates of zinc silicate and ferrite (> 1µm); (2) use of special apparatus to measure the solubility of these phases at 100 ° intervals from 50 to 350 °C

over a range of pH encompassing those relevant to BWR and PWR operating conditions; (3) characterization of the solids recovered after completion of these tests to ensure that the primary phases still existed and to observe any alterations in their physical properties that may have taken place, and to look for new phases that may have formed; (4) conducting independent experiments in an attempt to elucidate the probable mechanism/s by which these phases form under BWR and PWR operating conditions.

Solubility of Zinc Silicate and Zinc Ferrite in Aqueous Solution to High Temperatures

Donald A. Palmer (retired), Lawrence M. Anovitz and Leslie L. Wilson

Chemical Sciences Division, Oak Ridge National Laboratory, P.O. Box 2008 Oak Ridge, TN 37831-6110, U.S.A.

Email: Solution_Chemistry@comcast.net

Crystalline zinc silicate and zinc ferrite were prepared and characterized. The solubilities of these phases were measured using a flow-through apparatus from 50 to 350 oC in 100 o intervals over a wide range of pH. Both solid phases dissolve incongruently, presumably to form ZnO(s) and Fe2O3(s) or an amorphous ferric hydroxide, respectively. The respective concentrations of zinc(II) and iron(III) matched the reported solubilities of these phases, whereas the corresponding Si(IV) and Zn(II) concentrations were at least an order of magnitude below the solubility limits for pure SiO2 and ZnO2, respectively. Therefore, the solubility constants for zinc silicate and zinc ferrite were determined with respect to the known solubility constants for ZnO(cr) and Fe2O3(s), respectively, and the corresponding measured concentrations of Si(IV) and Zn(II). The results of additional batch-type experiments, as well as those reported in the literature, provide insights into the possible mechanisms of formation of zinc silicate and zinc ferrite in the primary circuits of nuclear reactors.

The components of zinc silicate (willemite) and zinc ferrite (franklinite) are the “end-member oxides”, viz., of Zn(II), Si(IV) and Fe(III), which have widely differing solubilities in pure aqueous solution in the general order SiO2 > ZnO > Fe2O3 (or even Fe(OH)3). In such cases these “binuclear” solid phases are expected to dissolve incongruently, leaving the more soluble component in solution while precipitating the less soluble component as its oxide or hydroxide. Obviously the immediate test for incongruent dissolution for these two solid phases is the observation that the zinc:silica and zinc:iron concentration ratios are not 2:1 and 1:2, respectively. To simplify further discussion, our study showed in fact the zinc:silica ratio is <<2:1 and the zinc:iron ratio is 1:<<2. Therefore, the solubility equilibria are anticipated to be:

2 4(s) 2 (l) (s)

4

Zn SiO H (2 )H O 2ZnO +

Si(OH)nn

n n+

−

+ + − (1)

2 4(s) 2

2 3(s) 3 2 (l)

ZnFe O H Zn(OH)

Fe O {or FeOOH or Fe(OH) }+( 1)H O

nnn

n

+−+ +

− (2)

The precipitate may be either amorphous, if formed at low temperatures or if dissolution is rapid, or it may be crystalline if formed at high temperatures and if the oxide/hydroxide solid tends to nucleate readily. Therefore, in order to interpret the solubility data for these incongruent equilibria, the corresponding data for the end-members, ZnO(s) and Fe2O3(s) or other oxide/hydroxide phases of Fe(III), are required to establish the thermodynamic properties of Eqs. (1) and (2) unambiguously.

The solubility of ZnO(cr) has been established from our previous studies [1-3] of the generalized equilibria:

+ n2-n 2(cr) (l)ZnO + H Zn(OH) ( 1)H On n+ − (3)

where n = 2, 1, 0 and -1. The solubility constants involving the four dissolved zinc(II) species, which dominate the solution speciation from ca. 0.1 mol·kg-1 H+ to at least 0.01 mol·kg-1 OH-, are known from 25 to 350 °C. These results apply to solutions in the absence of complexing ligands and were treated with an activity coefficient model that ignores ion pairing by assuming that all departures from ideal behavior are accounted for by the activity coefficients. Note that the densities of water used by Bénézeth et al. [3], and in the current work, were taken from Harvey [4].

The solubility of silica, whether as crystalline quartz or amorphous silica, is also well established over the pH range of interest to this report [5-7], given the caveat that the latter phase is metastable and solubilities will decrease over time as the particles “ripen”, particularly at higher temperatures where the kinetics of recrystallization become faster. There are just two solubility equilibria that need to be considered here; a fact that was substantiated by the results of our present study, viz.:

(s) (l)0

2 2SiO +2H O Si(OH)4

-5

(4)

- -2(s) 2 (l) 5SiO + OH 2H O Si(OH)+ (5)

Von Damm et al. [6] measured the solubility of quartz from 0 to 350 °C at saturation vapor pressure according to Eq. (4) and combined their results with available data for quartz from 45 to 900 °C and pressures to 986 MPa.

Rimstidt [5] later confirmed the validity of their Eq. [6], at least to 325 °C and 50 MPa. Fournier and Marshall [8] established a simpler equation for the corresponding equilibrium involving amorphous silica, which is valid from 0 to 350 °C at saturation vapor pressure.

Busey and Mesmer [7] conducted a potentiometric investigation of the hydrolysis equilibrium:

0 -4Si(OH) +OH Si(OH) (6)

in NaCl solutions (to 5 mol·kg-1) to 300 °C. Wesolowski et al. [9] combined their hydrolysis constants with Eqs. (4) and (6) to produce solubility constants for equilibrium (5) for quartz and amorphous silica. These equations apply strictly temperatures of ≤ 300 °C. Amorphous silica rapidly ripens to a crystalline solid at these high temperatures, such that the amorphous solid phase has little relevance above 300 °C.

The reliability of the solubility data for Fe(III) oxides and hydroxides is questionable at best. The experimental database is extensive, but most of the earlier results refer to fine-grained (or colloidal) material that was generally not characterized. There are far more limited results for established crystalline phases, such as goethite (α-FeOOH) and hematite (α-Fe2O3) [10-12] or the metastable forms, lepidocrocite (γ-FeOOH) and maghemite (γ-Fe2O3), which are highly insoluble from mildly acidic to mildly basic conditions and the kinetics of dissolution are very sluggish at low temperatures. Experiments conducted from supersaturation are hampered by the rapid formation of amorphous phases. In this report we have relied on the evaluation of the available data for hematite by Wesolowski et al. [9] for the generalized reaction:

+2 3(cr) 3- 2 (l)0.5Fe O + H Fe(OH) ( 1.5)H On

nn n+ − (7)

Preparation and Characterization of the Solid Zinc Phases

Zinc Silicate

There are many different methods available in the literature for the synthesis of willemite, Zn2SiO4. However, none appeared appropriate for producing well-crystallized, coarse-grained solids and attempts to use fused-salt methods as described below for zinc ferrite produced a variety of contaminated Zn2SiO4 solids. Finally, a method was adopted based loosely on a synthesis described in the literature [13]. A solution containing 250 mL of pre-hydrolyzed 1 mol·L-1 tetraethyl orthosilicate (TEOS), 91.72 g zinc acetate and 5 g of polyoxyethylene (2) sorbitan monooleate (TWEEN 80) was first adjusted to an approximate pH of 10 by addition of concentrated ammonium hydroxide. The sol, which formed immediately, was contained in a large glass evaporating dish which was then placed on a hot plate and heated gradually to dryness. The dry solid powder was transferred to a large platinum boat which was then placed in a muffle oven and heated slowly to 1000 °C. After two days the solid was recovered and washed with

2

large amounts of distilled, deionized water followed by ethanol. The XRD pattern of the dried Zn2SiO4 always showed the presence of unreacted ZnO, although the level of the contaminant was reduced by further heating at 1100 °C for two weeks. The ZnO could not be leached by washing with dilute HCl or ammonia. However, this problem was overcome by first calibrating the XRD analysis by adding known masses of pure ZnO to a know mass of the impure Zn2SiO4 and thereby determining the level of contamination from the respective peak heights of the principal ZnO band. An equivalent number of moles of solid SiO2·H2O were then added to the Zn2SiO4:ZnO mixture to convert the remaining ZnO to Zn2SiO4 and the resulting solid was heated for ten days in a platinum boat at 1100 °C. Every two days the boat was removed and after cooling the powder was gently ground in a motor and compacted back into the boat. The narrow peaks attest to the crystallinity of this material and only pure Zn2SiO4 was detected in the XRD databank. SEM images of this purified material showed that the particles were clearly crystalline with well-developed crystal faces in many cases, although the range of crystal sizes varied considerably.

A BET analysis of the starting material gave an average surface area of (0.68 ± 0.10) m2·g-1. Note that following completion of the final 350 °C experiment the surface area of the zinc silicate charge recovered from the pressure vessel was (0.71 ± 0.06) m2·g-1 indicating that no measurable change had occurred in particle size during all the experiments carried out from 150 to 350 °C.

Similarly, XRD patterns obtained for solid samples removed from the inlet and outlet ends of the columns after completion of the 50 and 350 °C experiments showed sharp peaks corresponding only to willemite and certainly no zincite could be detected within the ca. 5% composition detection limit of the XRD. Moreover, in SEM images of this material there was clear evidence of crystal faces indicative of some degree of recrystallization and there also appeared to be a more even distribution of particle sizes. Nevertheless, the overwhelming evidence is that during all these experiments the solubility of crystalline willemite was measured.

Zinc Ferrite

Zinc ferrite was prepared from a slurry of fine-grained ZnO and Fe3O4 in water in a 1:1 mole ratio that was heated in a platinum crucible to about 110 °C in air to drive off the water then to 1000 °C for two days. This produced a pure ZnFe2O4 which consisted of small particles (total surface area of ca. 27 m2·g-1). This red-brown solid was well mixed in a NaCl:Na2BB4O7·H2O flux, then placed in the platinum boat again and heated to 1000 °C for three days, which allowed the fine particles to recrystallize to well-defined, large crystalline

particles. After cooling the dark-brown solid was repeatedly washed with dilute HCl and finally with distilled, deionized water until the washings were pH neutral and no chloride could be detected with an AgNO3 test.

The XRD pattern of this material clearly demonstrated that it was pure zinc ferrite with very sharp peaks. A BET analysis of the same material gave an average surface area of (0.23 ± 0.1) m2·g-1. The SEM images provided striking evidence of the crystallinity of this solid.

Samples of the ZnFe2O4 solid charge were taken from the pressure vessel during the 150 °C series of experiments after a leak forced us to dismantle the entire apparatus and again after completion of the final 350 °C experiment. As in the case of the corresponding Zn2SiO4 experimental series, XRD patterns of the samples taken from both ends and the middle of the charge in the pressure vessel showed only the presence of the original phase. Apparently, the hematite that we postulate must form during the incongruent dissolution of zinc ferrite, must be at a sufficiently low mole percentage as to undetectable by XRD (<5%). This is not unreasonable given the low solubility of the spinel that would produce minor amounts of hematite.

Experimental Technique

Apparatus

Two types of flow-through experiments were performed to determine the solubility of Zn2SiO4 whereas only one of these methods was employed in the study of ZnFe2O4 solubility. These apparatus are described in previous EPRI reports [14-16], which focused on measurements of the solubility of cuprous and cupric oxides utilizing a high-temperature, flow-through apparatus at temperatures ≥ 100 °C and a low-temperature apparatus for temperatures < 100 °C. Only the former apparatus was used in the Zn2SiO4 study although it was modified to allow 50 °C measurements to be carried out. However, as we were able to prepare larger amounts of the Zn2SiO4 powder, both apparatus could be run simultaneously reducing not only the overall time required to complete these measurements, but also allowing more pHs to be investigated at 50°C at slower flow rates with longer reaction vessels. The common flow rate of the feed solution for the former apparatus was 0.1 mL·min-1 through the 35 cm long x 1.27 cm diameter Zircaloy 4 vessel, corresponding to a retention time for the solution assuming plug-flow conditions of two to three hours. For the low-temperature where flow rates were typically 0.06 mL·min-1 through three PEEK (PolyEtherEther-Ketone polymer) columns connected in series (total length of 75 cm with a

3

diameter of 0.075 cm) the contact time of ca. 33 hours. However, no variations of the measured zinc and silica concentrations were detected when faster flow rates were.

One crucial aspect of the high-temperature apparatus is that a 0.6% high-purity HNO3 solution was constantly supplied from a second HPLC pump via a platinum mixing T into the hot sample line at the exit of the pressure vessel. The addition of acid ensured that the dissolved metal did not precipitate in the platinum sample line as the solution cooled prior to reaching the sampling valve. The nitric acid solution was constantly checked to ensure that no detectable metal ions were in solution and the flow rates of both the feed solution and the nitric acid were monitored during each sampling event by recording the weight loss of the two solution reservoirs. The flow rate of the acid was generally maintained at one third that of the feed solution to guard against any back diffusion of the acidified solution into the reaction vessel. Moreover, the actual samples collected thereby contained ca. 0.15 to 0.2% HNO3 which was the matrix used for the analytical standards.

In the case of the low-temperature apparatus, the collection syringe was partially immersed in the 50 °C water bath to remove any temperature gradient. Note that addition of a known mass of 0.6% nitric acid to the collection syringe prior to each sampling event or after collection of the sample gave identical Zn and Si analyses.

Analytical Methods

Three techniques were employed for analyses of zinc, silica and iron, namely, standard flame emission atomic absorption (AA) using a Perkin Elmer 3110 instrument, graphite furnace atomic absorption (GFAA) spectrophotometry using a Varian AA2407 instrument and inductively coupled plasma mass spectrometry (ICPMS) using either a Thermo Finnigan Element2 High Resolution instrument or a Thermo Fisher Scientific X-SeriesII Quadrupole ICP-MS fitted with a MCN-6000 autosampler and sample introduction system. The former ICPMS instrument provided significantly improved sensitivity over the latter, but it experienced a number of long outages such that the second instrument was employed during those periods of time. Zinc was generally analyzed by AA at 214.2 nm where it gave a linear response over a very narrow range of 0.1 to 1.0 ppm requiring substantial dilutions of some samples. The strength of the emission peak for zinc made it very difficult to analyze by GFAA. Therefore, most of the samples were analyzed for zinc by ICPMS. Iron was analyzed by all three techniques depending on the concentration range and availability of the instruments. Silica was analyzed exclusively by GFAA at 251.6 nm.

Generally the analytical results of between 8 and 29 samples were averaged at each condition depending on the reproducibility of the analyses. This number of samples does not include those taken after each change of conditions where the trend in the analytical results clearly showed that the system had not re-equilibrated.

Chemicals

The nitric acid solution that was used to stabilize the sample and standard solutions was obtained from J.T. Baker (ULTRX II Ultrapure reagent) was diluted with distilled, deionized (Barnstead, NANOpure four stage system) water and stored in acid-washed polypropylene containers. Pure trifluoromethanesulfonic acid (commonly known as “Triflic acid”, CF3SO3H) was prepared by vacuum distillation of commercial (Alfa Aesar, 99%) reagent and diluted with distilled, deionized water. Note that this strong acid was used in these solubility measurements because it has little, if any, tendency to complex metal ions and it is thermally stable to high temperatures in dilute solutions. The zinc, iron and silica standard solutions were prepared from 1000 ppm AA/ICP calibration solutions (Aldrich Chemical Co.) by dilution with 0.2% HNO3. The pH buffers, ammonia (J.T.Baker, ULTREX Ultrapure reagent), sodium dihydrogenphosphate (99.999%, Aldrich Chemical Co.), and boric acid (99.9%, US Borax) were used without further purification. The ammonia stock solution was standardized potentiometrically against a standard HCl solution.

All dilutions and solution preparations were performed on a mass basis. The sample solutions were collected either in disposable polypropylene-polyethylene syringes (Aldrich Chemical Co.) or high-density polyethylene Nalgene narrow-mouth bottles which had been presoaked for at least 24 hours in 0.2% HNO3, then washed five times with deionized water immediately prior to use. Blank tests, in which the bottles and syringes were exposed to 0.2% HNO3 over extended periods of time, showed no detectable release of zinc, iron or silica. The NaOH stock solutions were prepared from 50% (by weight) NaOH, stored under argon, and standardized by mass titrations against potassium hydrogenphthalate.

Results and Discussion

The basic need prior to discussion of the results is the calculation of the pH of these solutions at the temperature and ionic strength of each experiment. Fundamental to this approach is the definition of pH as -log10(hydrogen ion molality). Implicit in these calculations is the use of an activity coefficient model, and in view of the generally low ionic strengths employed, the Meissner model was chosen with NaCl as the model electrolyte

4

following Archer’s equations [17]. Moreover, although application of the activity coefficient model to the determination of pH at infinite dilution is straightforward, for the dissolved metal ions their activity coefficients must be calculated iteratively by first estimating the metal speciation in solution and then applying the appropriate activity coefficient correction to give the molality of the species at infinite dilution. Note that the usual assumption was made that the activity coefficient of a neutrally charged species is unity.

The following equilibria must be considered for the dissociation of water and the hydrolysis constants of the various pH buffers used. In the low and high pH regions, where CF3SO3H and NaOH were used (usually > 10-4 mol·kg-1), respectively, the pH was determined directly from the free H+ molality, , in the first case and in the latter

case from the OH

+Hm

- molality, , converted to pH using the corresponding dissociation constant of water, K

-OHm

w, corrected where necessary for those consumed by the dissolution of the solid.

The values of Kw, were derived from the equation of Marshall and Franck [18].

H2O(l) H+ + OH- (8)

w H H OH OHK m mγ γ+ + − −= (9)

where +Hγ and -OH

γ are the molal activity coefficients of the ions. The Marshall and Franck equation provides estimates of Kw at finite temperatures (T in kelvin) and pressures (water densities, ρw, g·cm-3 [4]):

log10Kw=-4.098-3245.2/T+223620/T2- 3.984x107/T3

+ (13.957-1262.3/T+856410/T2)log10ρw (10)

For feed solutions containing ammonia at pressures at or slightly above the saturation vapor pressure, the hydrolysis of aqueous ammonia must be considered:

+4NH NH3(aq) + H+ (11)

The temperature dependence of the hydrolysis constant, Ka(11), is given by the expression (in terms of t in °C) which is based largely on the experimental work of Hitch and Mesmer [19].

log10Ka(11) = -10.0106+0.0329798t–7.86471x10-5t2 +9.68690x10-8t3 (12)

Ammonia forms complexes with Zn2+ which become progressively weaker with increasing temperature, in contrast to the reverse trend for the anionic forms of the other pH buffers used. Therefore, ammonia buffers were only used at ≥250 °C.

The boric acid, sodium borate buffer (in fact, solid boric acid and sodium hydroxide solutions were used to make these buffer mixtures) was used only at low temperatures to fix the pH because at higher temperatures complexation of borate with Zn(II), Fe(III), and especially Si(IV), in solution may become significant. Nevertheless, because high concentrations of boric acid could still complex these hydrolyzed species even at low temperatures, only low concentrations of boric acid were used. The dissociation of aqueous boric acid is described by the equation:

-3 (aq) 4B(OH) B(OH) + H+

)

+

(13)

Polymeric borate species do not need to be considered at such low concentrations of total borate, < 5x10-3 mol·kg-1. The following equation the dissociation constant of B(OH)3(aq) is [20]:

log10Ka(13) = -36.2605+3645.18/T+11.6402log10T +(16.4914-0.023917T) log10ρw (14)

The equilibrium involving the di- and mono-hydrogenphosphate buffer (actually a mixture of NaH2PO4 and NaOH) is represented by,

- - 2-2 4 4 2 (lH PO + OH HPO + H O (15)

where the equilibrium is written in this pseudo-isocoulombic form to minimize the non-linear temperature dependence of the equilibrium constant. The expression for this base hydrolysis constant at infinite dilution [21] is:

log10Ka(15)=-246.045+17156.9/T+37.7345lnT -0.0322082T-897579/T2 (16)

Equation (10) allows conversion of Eq. (16) to its protolytic form, which corresponds to:

- 2-2 4 4H PO HPO + H (17)

Zinc Silicate Solubility

As indicated in the Introduction, the experimental results confirm that Zn2SiO4 did not dissolve congruently to give a 2:1 ratio of zinc and silicate in solution. Moreover, the solutions were strongly undersaturated at all temperatures (50 – 350 oC) with respect to amorphous SiO2 or even quartz by as much as two to three orders of magnitude, see Fig 1 for example. The experimental points in Fig. 1a obtained in three different pH buffering media are consistent verifying that no significant complexation occurred. The dashed curves in Fig. 1 were obtained by averaging the pH-independent values then applying the hydrolysis constant for Eq. (5) [7]. Except for one inexplicable value in Fig. 1a, the dashed curve agrees well with the high pH values, which was also the case for the intermediate temperatures.

5

pH(50°C)5 6 7 8 9 10 11 12

log 10

(mS

iγ)

-6

-5

-4

-3

-2

-1

0

crystallinequartz

amorphousSiO2

100ppb

1ppm

pHN

10ppm

1000ppm

100ppm

(a)

pH(350°C)3 4 5 6 7 8 9 10 11

log 10

(mSiγ)

-7

-6

-5

-4

-3

-2

-1

0

crystalline quartz

amorphousSiO2

100ppb

1ppm

pHN

10ppm

(b)

100ppm

1ppt

Figure 1: Logarithm of silica activity as a function of pH at (a) 50 oC and (b) 350 oC: dot-dot-dashed line, amorphous SiO2; solid line, quartz; dashed line, for Zn2SiO4(cr); ▲, NaOH(aq); ∇, B(OH)3-NaOH(aq); ○, NaH2PO4-NaOH(aq); □, CF3SO3H-NH3(aq).

Based on the assumed incongruent dissolution of zinc silicate, the silica analyses showed that the dominant equilibrium over a wide range of pH is:

02 4(cr) 2 (l) (s)Zn SiO 2H O 2ZnO + Si(OH)+ 4 (18)

temperature, °C50 100 150 200 250 300 350

log 10

Ks(4

/18)

-6.5

-6.0

-5.5

-5.0

-4.5

-4.0

-3.5

-3.0

-2.5

-2.0

-1.5

Figure 2: Logarithm of solubility constant for Eqs.

(4) and (18) as a function of temperature in oC: medium-dashed line, amorphous SiO2 [9]; solid and short-dashed lines, quartz [6] and [9], resp.; dashed-dot line, zinc silicate Eq. (19); ●, this study.

Equation (19) was derived from a weighted fit of the log10K(18) values at 150, 250 and 350 oC:

log10Ks(18) = -(1.098±0.394) + (2068±218)/T (19)

The value obtained at 50 oC was omitted from the fit because the corresponding zinc molalities were far below the solubility limits of ZnO(cr) on the acidic side of the solubility minimum; indeed they were almost independent of pH. This may be due to slow kinetics of dissolution at low temperatures, but the fact that the corresponding silica analyses are higher than those expected from a simple extrapolation of Eq. (19) seems inconsistent with this explanation. The only feasible explanation offered at this time is that the released Zn2+ cations were re-adsorbed onto the zinc silicate crystal surfaces [9], noting the high surface to solution ratio used at this temperature. Presumably, at higher temperatures where Zn2+ adsorption would be even stronger, the lower solution to solid ratio and the higher flow rates would serve to saturate the surface sites. This rationalization is purely speculative, so that the dissolution equilibrium (18) can only be defined between 150 and 350 oC from this study. Interestingly, a similar phenomenon was observed for the corresponding zinc ferrite study (see below).

As alluded to above, the dissolution equilibrium at high pH,

2 4(cr) 2 (l) (cr) 5Zn SiO 2H O 2ZnO +Si(OH) H− ++ +

(20)

is defined by the Ks(18) and the corresponding values of Kw [18] and Ka(6) [7] where:

log10Ka(6) = -18.4014+2346.69/T+2.57979lnT (21)

Again, following the assumption that the dissolution of )Z proceeds according to Eq. (1), an hypothesis that is generally supported by the agreement between the measured zinc concentrations and the reported solubilities of ZnO(cr) [1-3] according to Eq. (3):

2 4(crn SiO

10 slog (3, 2) 4.0168 4527.66 /K n T= = − + (22)

10 slog (3, 1) 6.3065 2278.86 / 0.0035905K n T T= = − + + (23)

10 slog (3, 0) 5.6431 3139.58 / 0.009940K n T T= = − − (24)

s10

w10

log (3, 1) 4.5097 3319.75/0.0043268 8.2466log

K n TT ρ

=− =− −− +

(25)

Equation (22) is strictly only valid to 300 oC [1], but at such high temperatures the Zn2+ species contributes insignificantly to the dissolved zinc inventory in the present system.

Zinc Ferrite Solubility

For the equilibria involved in the incongruent dissolution of zinc ferrite in aqueous solution, Eq. (2) where n = -1 to 2, the nature of the precipitated ferric oxide or hydroxide (amorphous or crystalline) phase will depend primarily on temperature and to a lesser extent on pH and probably on the history of the experiment (i.e., whether Fe2O3(cr) preexists from previous high-temperature experiments and on the duration of these experiments). Regardless, of

6

the nature of the solid iron(III) phase present, the relationship between the dissolved zinc(II) species on the one hand and the dissolved iron(III) species on the other will be governed by the independent hydrolysis equilibria for each metal with the reasonable assumption that at the low concentrations of these metals in solution, no multinuclear metal hydroxide species will form.

Plots of the logarithm of Zn(II) molality versus pH at 50 and 350 oC are shown in Fig. 3. Clearly the zinc concentrations are significantly below those estimated for crystalline ZnO at all four temperatures (Figs. 1a-1d). The dashed curves were derived from isothermal fits of the experimental data using Eqs. (22-25) with a variable offset parameter to preserve the known hydrolysis behavior. The values for the acidic solutions at 50 and 150 oC are inconsistent (as was observed above for Zn2SiO4) and were not included in the fits, which are based on the following equations:

pH(50°C)3 4 5 6 7 8 9 10 11 12

log 10

(mZnγ)

-9

-8

-7

-6

-5

-4

-3

ZnO(cr)

1ppb

10ppb

100ppb

1ppm

10ppm

pHN

(a)

pH(150°C)2 3 4 5 6 7 8 9 10 11

log 10

(mZnγ)

-8

-7

-6

-5

-4

-3

1ppb

10ppb

100ppb

1ppm

10ppm

pHN

(b)

pH(250°C)3 4 5 6 7 8 9 10 11

log 10

(mZnγ )

-8

-7

-6

-5

-4

-3

1ppb

10ppb

100ppb

1ppm

10ppm

pHN

(c)

pH(350°C)3 4 5 6 7 8 9 10 11

log 10

(mZnγ)

-8

-7

-6

-5

-4

-3

1ppb

10ppb

100ppb

1ppm

10ppm

pHN

ZnO(cr)

(d)

Figure 3: Logarithm of zinc activity as a function

of pH at (a) 50, (b) 150, (c) 250 (d) 350 oC: solid line, ZnO(cr) [1-3]; dashed line, ZnFe2O3(cr); ▲, NaOH(aq); ∇, B(OH)3-NaOH(aq); ○, NaH2PO4-NaOH(aq); □, CF3SO3H-NH3(aq); ♦, CF3SO3H.

2Zn s s

s s 1

{ (3, 2) (3, 1)(3, 0) (3, 1) / } ( )

m K n h K n h

2K n K n h p p= = + =+ = + = − − + t

(26)

where h is the hydrogen ion molality, t is the temperature in degrees Celsius, p1 = (1.09±0.12) and p2 = (1.6±0.5)x10-3.

These experimentally based solubilities of ZnFe2O3(cr) from Eq. (2) in terms of zinc molalities are derived relative to a Fe(III)-bearing solid phase, the morphology of which depends on many factors. It is not surprising that in these experiments where the sequence was from low to high temperatures with undetectable amounts of Fe(III) containing solids being formed that extremely fine-grained, sub-micron sized amorphous material would exist at low temperatures with probably coarser-grained crystalline particles finally persisting at the higher temperatures. This behavior appears to be borne out by the results of the iron analyses shown in Fig 4.

First, the range of pH depicted in Fig. 4a-4d is shown to encompass the experimental results obtained in this study and does not necessarily imply that the models [9,11] were intended to incorporate this entire pH range. Second, the solubility curve for crystalline hematite [9] shown in Fig. 4a was obtained by extrapolation of higher-temperature data. Third, our experimental data for the Fe(III) solid phase cannot be considered as being a reliable indicator of the solubility of that phase, because of the inability to identify the phase present.

Despite the stated limitations of the results depicted in Fig. 4, with the exception of those obtained at 350 oC (Fig. 4d), which appear to be suspect for reasons unknown, they are quite interesting. Certainly the hydrolysis speciation, mainly involving:

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pH(50°C)3 4 5 6 7 8 9 10 11 12

log 10

(mFeγ )

-11

-10

-9

-8

-7

-6

-5

Fe2O3(cr) [9]

0.1ppb

100ppb

0.01ppb

10ppb

1ppb

pHN

(a)

pH(150°C)2 3 4 5 6 7 8 9 10 11

log 10

(mFeγ )

-10

-9

-8

-7

-6

-5

-4

Fe2O3(cr) [9]

pHN

1ppb

100ppt

10ppb

100ppb

1ppm(b)

pH(250°C)3 4 5 6 7 8 9 10 11

log 10

(mFeγ)

-10

-9

-8

-7

-6

-5

-4

pHN

1ppb

100ppt

10ppb

100ppb

1ppm(c)

Fe2O3(cr) [9]

Fe2O3(cr) [10] at 300 oC

Fe2O3(cr) [11]

pH(350°C)3 4 5 6 7 8 9 10 11

log 10

(mFeγ)

-10

-9

-8

-7

-6

-5

-4

pHN

Fe2O3(cr) [9]

1ppb

100ppt

10ppb

100ppb

1ppm(d)

Figure 4: Logarithm of Fe(III) activity as a

function of pH at (a) 50, (b) 150, (c) 250 (d) 350 oC: solid line, Fe2O3(cr) [9]; short-dashed line, Fe2O3(s); dotted line, Fe2O3(cr) at 300 oC [10]; long-dashed line, Fe2O3(cr) [11]; ▲, NaOH(aq); ∇, B(OH)3-NaOH(aq); ○, NaH2PO4-NaOH(aq); □, CF3SO3H-NH3(aq); ♦, CF3SO3H.

03Fe(OH) H Fe(OH)4

+ −+ (27)

appears to be consistent with the model of Wesolowski et al. [9], and with the data from Yishan et al. [10] at 300 oC and Diakaniv et al. [11] at 250 oC (Fig. 4c). Consequently, the short-dashed curves, which were obtained by applying an offset correction to the model of Fe2O3(cr) [9], become systematically closer to the model from 50 to 250 oC, consistent at least with a trend from amorphous to crystalline particles. Ironically, the first two experiments conducted at 250 oC gave iron(III) analyses that conform to the model of Wesolowski et al. [9] whereas the remainder show remarkable agreement with the data from Yishan et al. [10], albeit that the latter are reported at 300 oC. Nevertheless, the mean deviation of all the results in Fig. 4d from the model is only (0.8±0.5), which is not large considering the inherent difficulties of these measurements. This statement must be taken literally when the currently available solubility data in the literature for hematite exhibit variations of several orders of magnitude, attesting not to the ineptitude of the investigators, but to the experimental challenges that these measurements entail, although improvements in analytical techniques, principally ICPMS, have helped the situation.

As with the results for willemite, Zn2SiO4, those for franklinite, ZnFe2O4, provide compelling evidence for incongruent dissolution reactions and show systematic behavior with pH principally, and temperature, for the formation of zinc and ferric oxide (hydroxide) solid phases, respectively. Ideally, more solubility experiments should have been performed, particularly at 350 oC, where the results are few and, for iron(III) in particular, ambiguous. However, research budgets and time constraints prohibited a continuation of this research.

Solid Phase Formation Mechanisms

Only one experiment was performed to investigate a possible mechanism by which zinc silicate may form in an operating BWR. The aim of this test was to determine whether zinc silicate could be formed simply from a homogeneous nucleation of dissolved silica and zinc species (Zn2+ and/or its hydrolyzed products). This possibility appeared to be a reasonable, because of the high solubility of silica and the low solubility of zinc silicate with respect to that of quartz. Stoichiometric amounts of ZnO powder and pure quartz were added to each end of a gold boat into which was carefully added ca. 19 ppm ammonia. The boat was inserted into a horizontal pressure vessel which was then heated to 350°C. After 28 days the vessel was allowed to cool slowly to room temperature. The extracted gold boat was found to contain a single mass of white, highly crystalline

8

material. The filtered solid was washed with water and ethanol, then vacuum-dried. XRD scans showed only peaks due to zinc silicate which must have formed via dissolution of the silica and zinc oxide through a homogeneous re-precipitation process. Of course, this finding does not exclude other possible paths for zinc silicate formation in the operating reactors.

A series of experiments were initiated to investigate mechanisms by which zinc ferrite may form under reactor operating conditions. The initial hypothesis was that the most likely mechanism would involve the rapid adsorption of zinc ions (possibly as Zn2+ or hydrolyzed equivalents) onto the surface of magnetite particles followed the slow substitution of zinc(II) for iron(II) at the surface of the magnetite grains and then the even slower diffusion of zinc into those grains. Indeed, an experiment similar to that described for Zn2SiO4 failed to produce any ZnFe2O4 by a dissolution/re-precipitation route, which was not surprising in view of the insoluble nature of iron(III) oxide at near neutral pH. In addition, some tests were done with fine hematite particulates and zinc oxide suspended in water at 250 and 300 oC, without success. Solid-state conversions of a mixture of ZnO and Fe2O3 have only been recorded at extreme temperatures on the order of 1200 °C such that this possibility can be immediately ruled out.

Unfortunately, none of the tests with magnetite and zinc oxide in water, or Zn-containing solutions, were successful in producing detectable amounts of crystalline franklinite, ZnFe2O4. Surface adsorption studies in our laboratories to ca. 250 oC [9], involving mainly rutile and magnetite have clearly quantified the adsorption characteristics of these mineral surfaces for metal ions with increasing pH and temperature. Japanese workers [22-24] have established that M(II) ferrites can be formed from Fe(OH)2 particles exposed to M(II)-containing solutions (M = Mg, Fe, Co and Mn) under oxidizing conditions even at near ambient conditions. Therefore, whether under PWR and BWR conditions, adsorption of Zn2+ ions onto Fe(OH)2 particulates occurs with radiolytic oxidation, or adsorption onto existing magnetite particles on the fuel-cladding followed by substitution into the magnetite lattice, provide pathways for the observed accumulation of franklinite in the crud cannot be distinguished or confirmed from the present study. However, in the absence of other evidence, the adsorption mechanism does appear to be the most probable in our opinion.

Acknowledgements

Financial support for this work was provided by EPRI, Palo Alto, California, USA under the supervision of Dr. Bo Cheng. Special thanks are

also due to Ms. T.J. Keever who carried out all the ICPMS analyses during the course of the project.

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