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AECL-10821, COG-93-81 Oxidation of Magnetite in Aerated Aqueous Media Oxydation de la magnetite dans des milieux aqueux aérés

Peter Taylor, Derrek G. Owen

MOL 2 7 Ri 0 f A -, 1Q0, -, * April 1993 avril AECL RESEARCH

OXIDATION OF MAGNETITE IN AERATED AQUEOUS MEDIA

Peter Taylor and Derrek. G. Owen

Whiteshell Laboratories Pinawa, Manitoba ROE 1LO 1993 AECL-10821 COG-93-81 OXYDATION DE LA MAGNETITE DANS DES MILIEUX AQUEUX AÉRÉS

par

Peter Taylot et Derrek G. Owen

RÉSUMÉ

Les équilibres métastables entraînant des phases moins stables que la phase hématite peuvent être considérablement plus oxydants que l'équilibre cal- culé entre l'hématite bien cristallisée et la magnetite. On tire, dans ce rapport, les relations de solubilité et de stabilité généralisées entre les phases magnetite et Fe2O3»xH2O pour décrire les équilibres métastables. Des essais sur des poudres de magnetite synthétiques dans des solutions aqueuses aérées montrent que l'hématite cristalline se forme en l'espace de quelques jours à des températures supérieures à 100°C dans l'eau pure ou dans les solutions contenant des anions (par exemple, Cl-, S0£-, HCOj) qui ne forment pas de complexes très forts avec les oxydes de fer. Toutefois, la présence de phosphate ou de silice dissout, le moyen de formation de l'hématite par la dissolution-précipitation est fortement inhibé et la maghémite est un produit métastable persistant. Far conséquent, on pense que le phosphate ou la silice retarde la tendance à l'établissement de l'équilibre entre la magnetite et l'hématite dans les eaux souterraines aérées conditionnées par la magnetite. On présente ces conclusions dans le contexte du stockage permanent à^j déchets de combustible nucléaire.

EACL Recherche Laboratoires de Whiteshell Pinawa (Manitoba) ROE 1L0 1993 AECL-10821 COG-93-81 OXIDATION OF MAGNETITE IN AERATED AQUEOUS MEDIA

Peter Taylor and Derrek. G. Owen

ABSTRACT

Metastable equilibria involving phases less stable than hematite can be significantly more oxidizing than the calculated equilibrium between well- crystallized hematite and magnetite. In this report, generalized solubility and stability relationships between magnetite and Fe2O3.xH2O phases are derived to describe the metastable equilibria. Experiments with synthetic magnetite powders in aerated aqueous solutions show that crystalline hematite is formed within days at temperatures above 100°C in pure water or solutions containing anions (e.g., Cl", SOJ-, HCO5) that do not form very strong surface complexes with iron oxides. In the presence of dissolved phosphate or silica, however, the dissolution-precipitation route to hematite is strongly inhibited, and maghemite is a persistent metastable product. Thus, phosphate or silica is expected to delay the approach to magnetite-hematite equilibrium in aerated groundwaters conditioned by magnetite. These findings are presented in the context of nuclear fuel waste disposal.

AECL Research Whiteshell Laboratories Pinawa, Manitoba ROE 1L0 1993 AECL-10821 COG-93-81 CONTENTS

Page

1. INTRODUCTION 1

2. THEORETICAL CONSIDERATIONS 2

3. AVAILABLE SOLUBILITY AND STABILITY DATA 3

4. GENERAL EQUILIBRIUM EXPRESSIONS 5

5. EXPERIMENTS 7 5.1 MATERIALS AND METHODS 7 5.2 EXPERIMENTAL STRATEGY . 8

6. RESULTS AND DISCUSSION 8 6.1 CHARACTERIZATION OF SOLID STARTING MATERIALS 8 6.2 PRELIMINARY EXPERIMENTS IN DISTILLED WATER AT 100 TO 2O0°C 9 6.3 IDENTIFICATION OF HIGHLY CRYSTALLINE PRODUCT 10 6.4 EFFECTS OF SOLUTES 11 6.4.1 Weakly Complexing Species 11 6.4.2 Phosphate and Silicate at 150°C 11 6.4.3 Experiments with Phosphate up to 250°C 12 6.5 EXPERIMENTS BELOW 100°C 13

6.6 POSSIBLE INVOLVEMENT OF FLUORIDE 13

7. CONCLUSIONS 13

ACKNOWLEDGEMENTS 14

REFERENCES 14

TABLES 19

FIGURES 20 1. INTRODUCTION

Interconversion reactions between iron oxides have been reviewed recently by Cornell et al. (1989) and Blesa and Matijevi6 (1989). Much of the research on this topic has dealt with low-temperature (<100°C) reactions of colloidal Fe(III) oxides and oxyhydroxides and the crystallization or dissolution of magnetite (Fe3O4) under reducing conditions. The more limited literature on oxidation of magnetite in aqueous media has been discussed by White (1990), in a review of electron-transfer processes involving iron oxide and silicate minerals. Little work has been published on the hydrothermal conversion of magnetite to hematite (o-Fe203), although Swaddle and Oltmann (1980) have investigated the related conversion of maghemite (-y-Fe2O3 or HFe5O8) to hematite. In contrast, the high-temperature (~1000°C) anhydrous magnetite/hematite equilibrium continues to receive detailed attention, because of its importance as a redox buffer (Hemingway 1990, O'Neill 1988). Crystallo- graphic aspects of magnetite/ hematite intergrowth and interconversion are discussed by Davis et al. (1968). The mechanism of dry oxidation of magnetite to maghemite or hematite or both by 02 at intermediate temperatures has also been the subject of several papers (e.g., Sidhu et al. 1977, Gillot et al. 1977, Egger and Feitknecht 1962, Gallagher et al. 1968, Stockbridge et al. 1961). This transformation can also proceed in aqueous systems by electron-transfer to oxidants other than elemental oxygen and in acidic solutions even in the absence of dissolved oxidants (Jolivet et al. 1988, 1990).

We are interested in the oxidation of magnetite to hematite in aqueous media, because Fe(II)/(III) redox equilibria are expected to be important in establishing and maintaining reducing conditions with respect to U02 fuel dissolution in groundwaters in a nuclear fuel waste vault excavated deep in granite, as envisaged in the Canadian Nuclear Fuel Waste Manage- ment Program (Hancox and Nuttall 1991). Reducing conditions are desirable in this context to prevent the oxidative dissolution of U02 (Lemire and Garisto 1989, Paquette and Lemire 1981, Shoesmith and Sunder 1991) and also to reduce the corrosion rate of metallic container materials (King et al. 1992). Magnetite is one of the possible sources of ferrous iron in a granitic waste vault.

In this report, we describe calculations for 25°C and experiments at temperatures from 25°C to 250°C to investigate the possible importance of persistent metastable Fe(III) oxides in the aqueous oxidation of magnetite. Related iron redox equilibria involving aqueous Fe2+ and Fe(III) oxides in soils and natural waters have been discussed by several authors (Lindsay 1979, 1988, Doyle 1968, Langmuir and Whittemore 1971, Schwab and Lindsay 1983). The influence of fine-grained or metastable phases on the Fe304/Fe203 couple in high-temperature anhydrous conditions is discussed by Hemingway (1990). - 2 -

2. THEORETICAL CONSIDERATIONS

Interconversion between magnetite and hematite can be expressed in various ways (Davis et al. 1968), including a simple oxidation (1), an electrochemical half-reaction (2) and incongruent dissolution (3).

4Fe3O4(s) + 02(g) ^ 6Fe2O3(s) (1)

+ Fe3O4(s) + O.5H2O(1) ,=» 1.5Fe2O3(s) + H (aq) + e- (2)

+ 2+ Fe3O4(s) + 2H (aq) =;=«* Fe203(s) + Fe (aq) + H2O(l) (3)

Oxidation of magnetite in aqueous solution could yield various other less stable Fe(III) oxides or oxyhydroxides instead of hematite, e.g., maghemite, goethite (a-FeOOH), lepidocrocite (7-FeOOH), or several other less well defined amorphous or crystalline materials. All of these phases have the general formula Fe2O3-xH2O, with 0 ^ x s 3, and they are hereafter discussed as "Fe(III) oxides." Interconversion between magnetite and one of these phases can be expressed by Equations (4) to (6), by analogy with (1) to (3).

4Fe3O4(s) + 6xH20(l) + 02(g) =t=S 6Fe2O3-xH2O(s) (4)

Fe3O4(s) + (1.5x+0.5)H20(l) ,=* 1.5Fe2O3-xH20(s) + H+(aq) + e- (5)

2+ Fe3O4(s) + 2H+(aq) ;==& Fe2O3-xH20(s) + Fe (aq) + (1 - x)H2O(l) (6)

For each of the three pairs of related reactions (i.e., (1) and (4), etc.)» the Gibbs energies of reaction, ArG°, differ by ArG° of Reaction (7), multiplied by the appropriate stoichiometric coefficients for Fe203 or Fe2O3.xH2O in the reactions (6.0, 1.5 or 1.0):

or-Fe2O3(s) + xH20(l) ^=s Fe2O3-xH20(s) (7)

Thus, any Fe(III) oxide less stable than hematite, in metastable equilibrium with magnetite, will establish a higher redox potential and p(O2), and lower equilibrium Fe2+ activity at given pH, than those expected for the magnetite/hematite assemblage. Some of these relationships have been discussed by Lindsay (1979, 1988).

Some iron hydroxides and oxyhydroxides contain additional anions, either as impurities or as essential structural components (Fox 1988, Ellis et al. 1976, Harrison and Berkheiser 1982). If present as essential, components, i.e., as basic iron(III) salts, such anions must be included in any equilibrium calculations. Fox (1988) has presented evidence that amorphous "Fe(0H)3" precipitated from nitrate solutions is in fact a basic iron nitrate, Fe(0H)2 4(N03)0#6. The following discussion applies only to equilibria between magnetite and true Fe(III) oxides or hydrous oxides and not to basic salts.

The relative stabilities of Fe(III) oxides are commonly compared by their solubilities (Langmuir 1969, Langmuir and Whittemore 1971, Lindsay 1979, 1988, Fox 1988), expressed by Equation (8): - 3 -

3+ 0.5Fe2O3-xH2O(s) + 0.5(3-x)/2H2O(l) ^à Fe (aq) + 30H"(aq) (8) whence the solubility products are given by expression (9):

3 3 Ksp = {Fe +}{OH"} (9) assuming unit activity of water. Here, {X} denotes the activity of species X. One order of magnitude increase in Ksp at 25°C corresponds to

twelve orders of magnitude increase in the equilibrium p(02)for Reaction (4) (again assuming unit activity of water); an increase of 177 mV in the electrode potential of Reaction (5); two orders of magnitude decrease in the equilibrium {Fe2+}, at given pH, for Reaction (6); and

1 an increase of 5.71 kJ-mol- in the ArG° of Reaction (7). Clearly, the redox properties of a magnetite/Fe(III) oxide assemblage depend strongly on the identity of the Fe(III) oxide phase. Conversely, very fine-grained magnetite or metastable reducing phases (e.g., Fe(0H)2 or Fe3(OH)a) could establish transient redox conditions more reducing than the crystalline magnetite/hematite buffer.

3. AVAILABLE SOLUBILITY AND STABILITY DATA

The literature on solubility and stability relationships among the various Fe(III) oxides is complex (Sadiq and Lindsay 1979; Berner 1969; Langmuir 1969, 1971; Langmuir and Whitteraore 1971; Lindsay 1979, 1988; Fox 1988, and references therein). Calculated and measured values of Ksp reported or cited in these references span more than seven orders of magnitude, from 10°7 or even higher for fresh amorphous precipitates (Sadiq and Lindsay 1979, Langmuir 1969, Fox 1988) to 10"44 or even lower for some reports on crystalline hematite and goethite (Langmuir 1971, Langmuir and Whittemore 1971, Sadiq and Lindsay 1979). Of all the Fe(III) oxides, thermodynamic data are most complete for hematite. Current data needed to calculate Ksp for crystalline hematite at 25°C are presented in Table 1, along with other thermodynamic quantities 41 9±o 4 relevant to the following discussion. These data yield Ksp = io- - - for crystalline hematite, in close agreement with the value of 10"41-91 recommended by Lindsay (1979, 1988). The data compiled by Wagman et al. 43 9 (1982; compiled in 1966) lead to Ksp = 10" - for hematite (see also the discussion of hematite by Sadiq and Lindsay 1979). This extremely low calculated solubility product arises mainly from a high value for A£G° of Fe3+, -4.7 kJ/mol. Larson et al. (1968) reported more negative values than were previously accepted for the Gibbs energies of formation of both Fe2+ and Fe3+, and their results have since been endorsed by Tremaine et al. (1977) and Sadiq and Lindsay (1981). Further work on magnetite solubility under reducing conditions by Tremaine and LeBlanc (1980) also supports the results of Larson et al. (1968). - 4 -

The relative stability of hematite and goethite is a topic of particular interest. On the basis of solubility measurements at 85°C, Berner (1969) concluded that fine-grained goethite is thermodynamically unstable relative to hematite plus water under most, if not all, geological conditions. In their solubility assessments, however, Langmuir (1971) and Lindsay (1979, 1988) both concluded that crystalline goethite is marginally stable with respect to hematite plus water. The Gibbs formation energy data derived by Langmuir (1971) indicate that Ksp of crystalline goethite at 25°C is lower than that of crystalline hematite by a factor of 10°-4±0-4, i.e., K = lO-42.3±o.6 for crystalline goethite} for comparison, Lindsay (1979, 1988) recommends K = 10"42-02. Hsu and Marion (1985), however, obtained activity products that varied between 10"39-80 and io-40-83 with ionic strength variations between 0.005 and 0.2 respectively. We hesitate to 42 adopt a Ksp value for goethite, but note that values much lower than 10" for any Fe(III) oxide phase should be viewed with caution.

Langmuir (1971) has discussed particle-size effects on the relative stabilities of goethite and hematite, using surface enthalpy estimates (1250 and 770 erg/cm2 respectively), derived from heats of solution measured by Ferrier (1966). Langmuir calculated that the solubilities of finely divided and crystalline goethite and hematite are related by the expression log!0Ksp(finely divided) - log10Ksp(crystalline) = 53/z for goethite, and 25/£ for hematite, where z. is the hypothetical cubic particle edge in nanometres. This indicates that particle-size effects on solubility become significant for crystallite dimensions below a few hundred nanometres. More recently, however, Hsu and Marion (1985) were unable to detect any dependence of solubility on particle size, in a series of measurements on goethite specimens aged for 9 to 16 years. If the heats of solution measured by Ferrier (1966) include an internal defect energy term as well as a surface- energy term, this would lead to an overestimate of particle-size effects on solubility in the calculations by Langmuir (1971). The influence of grain size and high-energy surface features, as well as surface treatment, on the solubility of gibbsite, A1(OH)3, has been discussed more recently by Hemingway et al. (1991), Wesolowski (1992) and Palmer and Wesolowski (1992).

The solubility of amorphous ferric hydroxide depends on the age of the material, aged samples having greater structural order or particle size or both, and hence lower solubility. Most reported Ksp values range from about 10-37 for fresh precipitates to 10"39 for aged precipitates or natural materials (Lindsay 1979, 1988; Langmuir 1969; Langmuir and Whittemore 1971; Fox 1988; Sadiq and Lindsay 1979).

Maghemite and lepidocrocite both appear to be intermediate in solubility between the amorphous hydroxides and well-crystallized hematite. Lindsay 40 41 40 61 (1979, 1988) recommends Ksp values of 10- - and 10- - for maghemite 38 8±0 5 and lepidocrocite respectively. Langmuir (1969) calculated Ksp = 10- - - and io-38-7±0-5 for freshly precipitated maghemite and lepidocrocite

1 erg = 0.1 /xJ - 5 -

respectively, based on unpublished data from R.W. Doyle and on estimated values near 10-40 for well-crystallized material. Sadiq and Lindsay (1988) have reported an experimental value of 10-40•36±0•07 for soil maghemite; unfortunately, some typographical errors in their paper make the determination hard to evaluate.

Given the significant uncertainties about relative stabilities of the Fe(III) oxide minerals, and the possible dependence of these phase relationships on particle size as veil as crystal structure, we have chosen to express magnetite/Fe(III) oxide equilibria in terms of the Ksp value of the Fe(III) oxide phase, rather than its mineralogical identity, except for crystalline hematite. This provides a flexible basis for comparing stability and solubility relationships, and it is adaptable to changes and improvements in the available thermodynamic data.

4. GENERAL EQUILIBRIUM EXPRESSIONS

To obtain an expression relating p(02) for Reaction (4) to Ksp for Reaction (8), it is useful to consider Reaction (10):

3 1 i/3Fe304(s) + /2H20(l) + /1202(g) ^> Fe3+(aq) + 30H"(aq) (10)

for which, from data in Table 1, ArG° = +204.4 ± 2.4 kJ/mol, and log10K = -35.8 ± 0.4. Reaction (10) is a linear combination of Reactions (4) and (8); hence, for equilibrium between crystalline magnetite and any Fe(III) oxide, at unit activity of water,

log10Ksp + 35.8 = 1/12 logloP(02) or log10p(02) = 12 log10K6p + 429.8 ± 5.0 (11)

5 (p(02) in bar, 1 bar = 10 Pa) .

Relationship (11) is illustrated in Figure 1. One interesting conclusion from this relationship is that log10p(02) > 0 if log10Ksp > -35.8, i.e., an Fe(III) oxide with a solubility product above this limit would be unstable with respect to oxygen evolution and transformation to magnetite, as well as with respect to other more stable Fe(III) oxides. This emphasizes the wide range of potential metastable or transient redox conditions that might be established by magnetite/Fe(III) oxide assemblages. It may be significant that this limiting value, log10Ksp = -35.8, is close to the upper limit of measured solubilities for amorphous Fe(III) oxide phases.

Similarly, an Fe(III) oxide with log10KBp < -42.7 would yield log10p(02) < -83.1, thus placing the Fe304/Fe(III) oxide boundary outside the stability limit of water with respect to hydrogen evolution. This gives further reason to question any reported Ksp value much lower than 10-42.

The 25°C electrode potential corresponding to any given p(02) value at pH = 0, derived from the thermodynamics of water decomposition, is given by - 6 -

E° = 1.229 + (0.05915/4) log10p(02)

Substituting expression (11) for log10p(02), we obtain the following electrode potential relationship for Equation (5):

E° = 7.585 + 0.17745 log10KBp ± 0.074 . (12)

This relationship is also illustrated in Figure 1. Reactions (6) and (8) can be combined algebraically to give Reaction (13)!

3+ 2+ Fe304(s) + 4H20(l) ^ 2Fe (aq) + Fe (aq) + 80H-(aq) (13) for which ArG° = +577.7 ± 5.3 kJ/mol, and log10K = -101.2 ± 0.9, whence, at unit activity of water, we obtain for Reaction (6)

2+ log10{Fe } = -101.2 - 2 log10Ksp + 28.0 - 2pH

= -(2 log10Ksp + 2pH + 73.2 ± 0.9) . (14)

2+ This relationship between Ksp, {Fe } and pH is illustrated in Figure 2. Note that hydrolysis of Fe2+ (to FeOH+, etc.) at 25°C becomes significant (>1% of total Fe(II)) at pH > 7.5 and predominant at pH > 9.5 (Baes and Mesmer 1976, Tremaine and LeBlanc 1980).

Equations (11), (12) and (14), as illustrated in Figures 1 and 2, thus provide a framework, for general discussion of magnetite/Fe(III) oxide equilibria at 25°C in terms of the relative solubilities -f the different Fe(III) oxides. The principles of this discussion apply to temperatures above 25°C, but a full evaluation of Fe(III) oxide solubility/stability relationships at elevated temperatures is beyond the scope of this paper. The main point of the foregoing calculations and discussion has been to establish the potential significance of metastable Fe(III) oxides in redox equilibria involving magnetite. A useful reference point for calculations above 25°C is provided by the Gibbs energy expressions derived by Hemingway (1990) for Reaction (1).

It is important also to consider the redox potential of a solution of given pH and Fe(II) activity in the absence of magnetite and how this is affected by the form of Fe(III) oxide present. This can be addressed by considering the reaction

2+ 4Fe + 02 + 4H20 --^ 2Fe203 + 8H+ (15)

2 log10K = -(8pH + 4 log10{Fe +} + log10p(02))

= 30.7 ±1.6 for Fe203 = hematite;

= (30.7 ± 1.6) - 4x for any Fe(III) oxide phase x orders of magnitude more soluble than hematite.

Whence an effective p(02) value (or a corresponding redox potential) can be determined for any combination of pH, {Fe2+} and Fe(III) oxide solubility. Equilibria of this sort have been discussed at some length by Doyle (1968) and Lindsay (1979, 1988). - 7 -

5. EXPERIMENTS

Temperatures in the near field of a Canadian nuclear fuel waste vault (i.e., in the vicinity of the fuel containers) are not expected to exceed 100°C (Hancox and Nuttall 1991). We investigated the oxidation of commer- cially available magnetite powders in aerated solutions at temperatures ranging from 25 to 250°C; we ran most experiments at 150°C. This temperature was chosen for experimental convenience, because extensive oxidation of the magnetite occurred within a few days. The purpose of these experiments was to see how rapidly the system approaches a crystalline hematite- magnetite phase assemblage and to determine which groundwater components could hinder the formation of hematite. The situation with metastable phases, described in the preceding section, is simplified above 100°C, with maghemite, goethite and hematite being the only Fe(III) oxides likely to occur; amorphous Fe(III) hydroxides and lepidocrocite are normally encountered only at lower temperatures (Blesa and Matijevic 1989, Cornell et al. 1989, Schwertmann and Cornell 1991).

5.1 MATERIALS AND METHODS

Experiments were performed with two synthetic magnetite powders, of 99.999% and 97% stated purity (metals basis), supplied by Johnson-Mat they. These powders are referred to hereafter as magnetites A and B respectively. Most experiments were performed with magnetite A, because of its higher purity and better defined microstructure (see below). Other materials used were powdered vitreous silica ("fused quartz," GEM-435, -200 mesh, Engineered Materials), aqueous colloidal silica (30% SiO2, Matheson Coleman & Bell), and various reagent-grade inorganic chemicals. Solutions were prepared using double-distilled or deionized water; no differences were observed in the behaviour of magnetite in the tvo waters.

All but a few experiments were run in polytetrafluoroethylene(PTFE)-lined 3 3 stainless steel vessels (20-cm capacity), using 0.5 g of Fe3O4 and 5 cm of solution. When the vessels were sealed, this gave about a threefold excess of Fe304 over oxygen, according to Equation (1). A few experiments were performed with excess oxygen, either by reducing the quantity of Fe304 to 0.1 g or by using a larger vessel (420-cm3 capacity, with 100-cm3 water). The 20-cm3 vessels were heated in a forced-air convection oven, and the 420-cm3 vessels were heated individually, as described by Taylor et al. (1979). Solid products were identified by X-ray diffraction (XRD), using a computer-controlled Rigaku Rotaflex diffractometer, with a 12-k.W rotating-anode Cu Kxx source and a diffracted-beam monochromator. Most products were also examined by scanning electron microscopy (SEM), using an ISI DS-130 instrument.

5.2 EXPERIMENTAL STRATEGY

Over 95% of the experiments performed fall into the following four main categories:

(i) Preliminary experiments were run with both types of magnetite in distilled water for 1, 6 or 20 d at 100, 150 or 200°C (all 18 permutations of starting material and conditions). - 8 -

(ii) Experiments were then run with magnetite A for 6 d at 150°C to investigate the effects of various inorganic solutes (NaCl, Na2S04, NaH2PO4, Na2HP04, NaHCO3, Na2CO3, Na2B4O7, NaOH, H2S^3, and SiO2). Most of these substances represent common grour.vWater solutes, and others were included to provide a range of solution pH and surface-complexation strength. (ill) Further experiments, with magnetite A in distilled water, were run at lower temperatures (22 to 90°C) and for longer durations d). (iv) Further experiments to investigate the effects of dissolved silica and phosphate were extended to temperatures up to 250°C. A few additional runs that do not fall into these four main categories are described where appropriate. Altogether, about 80 experiments were run.

6. RESULTS AND DISCUSSION

6.1 CHARACTERIZATION OF SOLID STARTING MATERIALS Magnetite B is a fine-grained powder (0.1 to 0.5 jtm), with irregular particle morphology. In contrast, magnetite A consists of distinctive, porous, briquette-like particles up to a few tens of micrometres in size, as shown in Figure 3; these particles in turn form aggregates up to about 100 /im across. The particles appear to have been formed by sintering a fine-grained powder (0.2 to 0.5 nm) in a controlled atmosphere at low pressure. The distinctive microstructure of magnetite A proved useful for distinguishing between solid-state and dissolution-precipitation mechanisms of oxidation. The powdered vitreous silica powder consisted of irregular shards, 5 to 20 yum in dimension.

Both of the magnetite starting materials yielded well-defined XRD patterns, with no peaks other than those expected for an iron oxide spinel phase. Calculated unit-cell parameters, a0, were 8.3936 ± 0.0017 (2a) and 8.3890 ± 0.0032 A* for powders of type A and B respectively, which may be compared with literature values of 8.396 A for magnetite and 8.3515 A for maghemite (Swanson et al. 1967, Schulz and McCarthy 1988). For comparison, we have obtained a0 = 8.394 A for a series of freshly reduced, (but air- exposed) magnetite powder; 8.3969 ± 0.0033 A for a series of freshly reduced, sintered magnetite disks; and 8.3512 ± 0.0015 A for a series of three maghemite samples prepared by oxidizing magnetite B in air at 200°C for 6 d. There is thus some evidence of partial bulk oxidation of both magnetite starting materials, particularly Type B, but the cell-parameters are much closer to magnetite than maghemite values.

1 A =0.1 nm - 9 -

According to the work of Stockbridge et al. (1961), a thin (~2.5 nm) surface layer of 7-Fe2O3, formed by ambient air oxidation, was probably present on the magnetite powder surfaces. No attempt was made to remove such a layer prior to our experiments; it is unlikely to impede magnetite oxidation significantly at higher temperatures, based on known kinetics of the magnetite-maghemite conversion (Sidhu et al. 1977, Gillot et al. 1977, Egger and Feitknecht 1962, Gallagher et al. 1968). To estimate the extent of prior oxidation, both magnetites were subjected to thermogravimetric analysis, heating in air at 5°C/min up to 500°C. Magnetites A and B experienced 95% and 54%, respectively, of the expected weight gain (3.45 wt.%) for complete oxidation of Fe3O4 to Fe2O3. This is consistent with the higher stated purity and (probably) lower specific surface area of magnetite A. 6.2 PRELIMINARY EXPERIMENTS IN DISTILLED WATER AT 100 TO 200oC The XRD patterns of reaction products included sharp hematite peaks (Morris et al. 1981) in every case except the 1-d runs at 100°C, in which hematite was barely detectable. With magnetite A, the variation in cell parameters of the spinel component of the oxidation products was only slightly greater than the uncertainty of the measurements. Much more pronounced variation was observed with magnetite B. Details of this variation will be reported elsewhere; the main observations are as follows: (a) At 100°C, the cell parameter declined from the initial value of 8.3890 to 8.363 ± 0.005 Â (average for three experiments). This indicates substantial oxidation of the magnetite towards maghemite; broad, weak features identifiable as maghemite superlattice reflections were also detected (see Schulz and McCarthy 1988). (b) At 150°C and 200°C, the cell parameter usually did not change significantly. In three cases, it increased significantly to values near 8.396 À, indicating restoration of the magnetite component to near the ideal stoichiometry, and conversion of the maghemite component to hematite. Accurate quantitative analysis of the product phases by XRD was not possible, because of the large particle-size of the high-purity magnetite, and consequent difficulty in quantifying the effects of X-ray absorption on diffraction peak intensities. The intensity (peak height) of the principal hematite peak (d = 0.270 nm) reached 30 to 40% of that of the principal magnetite peak (0.253 nm) after 20 d at 150°C or 6 d at 200°C, suggesting near-complete oxygen consumption in these cases; the hematite peak was a little less intense after 6 d at 150°C. Based on these preliminary runs, 6 d at 150°C was chosen as a convenient experimental condition for investi- gating the effects of various solutes on hematite formation.

Scanning electron microscopy of products from reactions involving magnetite B was not very illuminating, because both the starting material and the product consisted of submlcrometre particles. There was some indication of faceted particles in the products from longer runs at higher temperatures, suggesting a dissolution/precipitation process. This material was not investigated further, however, because magnetite A yielded much better defined crystalline products. -10-

In all the reactions with magnetite A, except the 1-d, 1OO°C run, there was clear evidence of crystal growth by a dissolution/precipitation process. Visually, most reaction products consisted of both fine-grained red precipitate (mainly hematite, based on the XRD data discussed below) and larger black, particles resembling the original magnetite. Micrographs of the product from the 6-d, 150°C run are shown in Figure 4. Figure 4(a) shows that the fine-grained material is almost monodisperse, with ~0.2-/^m particle size, and crystal facets just visible. A few larger crystals (~1 urn) are visible in the upper right corner of this figure. Elsewhere in the sample, localized groups of much larger crystals (5-10 inn) were observed; an extreme example is shown in Figures 4(b). The larger aggregates were coated with a thick layer of crystalline precipitate, replacing and/or obscuring the original magnetite microstructure. 6.3 IDENTIFICATION OF HIGHLY CRYSTALLINE PRODUCT The microstructures described above are characteristic of reaction products obtained from magnetite A after either 6 or 20 d at all three reaction temperatures. Evidently, a significant fraction of the hematite detected by XRD was formed by a dissolution-precipitation process. The formation of quite large iron oxide crystals (sometimes >5 fim) is unusual under these rather mild conditions of temperature and pH. Because these large crystals represent only a small volume fraction of the total precipitated product, and only about one third of the magnetite can be oxidized by the available oxygen in the 20-cm3 vessels, it is not immediately obvious whether these crystals represent growth of hematite or recrystallization of magnetite.

Magnetite crystals, both natural and synthetic, are most commonly octa- hedral, and less frequently cubic or icosahedral, whereas hematite sometimes assumes more complex euhedral habits (Booy and Swaddle 1978, Kolb et al. 1973, Palache et al. 1944, Sapieszko and Matijevic 1980, Schwertmann and Cornell 1991). Hematite formed by aqueous recrystallization of maghemite is reported to be rhombohedral in habit (Swaddle and Oltmann 1980). Growth of large (several micrometres) euhedral hematite crystals usually requires strongly alkaline conditions (Kolb et al. 1973, Schwertmann and Cornell 1991). In Figure 5, we compare hematite crystals grown from 1 freshly precipitated "Fe(0H)3" in ~l mol.L" NaOH at 200°C with some of the large crystals obtained by oxidation of magnetite in aerated water at 150°C. The similar crystal habits of the pure euhedral hematite and the large crystals formed in the magnetite oxidation experiments indicate that the latter crystals are, indeed, hematite.

These large crystals were only obtained in reactions with magnetite A, and not magnetite B. Their formation is probably a consequence of compara- tively slow release of dissolved Fe2+ from within the porous magnetite particles, resulting in low hematite nucleation rates, but extensive growth of those nuclei that do form. In a few experiments at or below 150°C, SEM revealed some acicular crystals as well as euhedral hematite, suggesting the formation of small quantities of goethite or one of the other Fe(III) oxyhydroxides; the contributions of these crystals to the XRD patterns were too weak for positive identification. There was no indication of such solids competing with hematite or - 11 - maghemite as solubility- and potential-determining phases on magnetite. Examples of twinned, lath-shaped crystals, closely resembling aged goethite particles described by Hsu and Marion (1985) are shown in Figure 6. 6.4 EFFECTS OF SOLUTES 6.4.1 Weakly Complexing Species Oxidation of magnetite for 6 d at 150°C in the following solutions was indistinguishable either by SEM or XRD, except for. some variations in hematite crystal size and habit, to oxidation in pure water: 0.1 mol.L*1 1 NaCl or Na2S04; 0.01 mol-L" Na2B407, Na2CO3, NaHCO3 or NaOH; examples of these reaction products are shown in Figure 7. Crystals with hexagonal- bipyramidal and rhombohedral habits are fairly common in the products from the latter four weakly alkaline solutions, consistent with the hexagonal symmetry of hematite. The crystals formed in these solutions rarely exceeded 2 fim in dimensions. We conclude that the presence of weakly complexing anions and pH variations in the neutral to weakly alkaline range have little effect on the course of magnetite oxidation at 150°C, beyond some modification of hematite crystal size and habit.

1 The product recovered from 0.01 mol-IT H2SO4 had a grossly altered, heterogeneous microstructure, as shown in Figure 8; the only phase detected in this product by XRD was poorly crystallized hematite. The nature of the reaction in this aggressive medium was not investigated further; the absence of any residual magnetite peaks in the XRD pattern indicates that substantial incongruent dissolution (Equation (3)) had occurred in this case. Incongruent dissolution of Fe2+ from magnetite at lower temperatures under moderately acidic conditions, leaving a residue of maghemite rather than hematite, has been described by Jolivet and Tronc (1988) and White (1990). 6.4.2 Phosphate and Silicate at 150°C Phosphate solutions were included in this study, because phosphate is known to bind strongly to iron oxide surfaces, and hence can influence oxide dissolution and precipitation (Borggaard 1983, Cornell et al. 1989, Hingston et al. 1972, Parfitt et al. 1975). Silica is known to have a marked effect on both the sorption and recrystallization behaviour of iron oxides and hydroxides (Cornell et al. 1987, 1989, Schwertmann and Cornell 1991, Vempati and Loeppert 1989, Zachara et al. 1987). In particular, traces of silica impurity strongly inhibit the hydrothermal conversion of maghemite to hematite (Swaddle and Oltmann 1980). Figure 9 shows micrographs of products recovered from reactions (0.5 g magnetite A, 5 cm3 solution, 20-cm3 vessel, 150°C, 6 d) in which either phosphate or silica was present. Expanded portions of some representative XRD patterns are shown in Figure 10.

1 Addition of either NaH2P04 or Na2HPO4 (0.01 moLL" ) reduced the formation of hematite to levels undetectable by either XRD and SEM. The product retained the microstructure of the original magnetite A (Figure 9a), but the XRD patterns indicated substantial conversion of magnetite to maghemite, i.e., high-angle shoulders or distinct secondary peaks appeared on the magnetite peaks (Figure 10). In some cases, weak magheraite super- - 12 - lattice peaks were also observed (Schulz and McCarthy 1988). Evidently, conversion of magnetite to maghemite occurred by a solid-state reaction 1 (Sidhu et al. 1977). At a concentration of 10° mol.L" Na2HP04 or NaH2P04, a small amount of hematite was formed, roughly one tenth of that obtained in pure water or weakly complexing solutions, and no large hema- 4 1 tite crystals were observed. At a concentration of 10" mol'L" Na2HP04 or NaH2P04 (1 #mol.g(Fe304)), there was no measurable difference from the pure system in either the quantity of hematite formed or the product morphology (Figure 9b).

Introduction of silica, either as powdered vitreous SiO2 or as an aqueous colloid, had similar effects to phosphate. Addition of 18 £iL of the colloidal SiO2 was sufficient to virtually eliminate hematite formation, and maghemite was again the principal oxidation product detected by XRD (Figure 10). As with the reactions in 10'3 or 10"2 mol-Ir1 phosphate solutions, the original magnetite A microstructure was retained, with no detectable solution-grown oxidation products. When only 5 /iL of colloidal silica was added, the quantity of hematite produced was about one tenth of that in the pure system, while 1 nL had no measurable effect on the reaction. In the former case, the hematite appeared as localized patches of platy crystals, similar to those shown in Figure 9c, while in the latter case the now-familiar euhedral habit was observed. Platy or disk-shaped hematite particles have been described previously by Sapieszko and Matijevic (1980).

Addition of fused silica powder, even as much as 50 mg, did not completely eliminate hematite formation, presumably because the time for sufficient silica to dissolve and then react with the magnetite surface is significant. However, 10 or 50 mg of fused silica reduced the quantity of hematite formed by about a factor of 10, and platy crystals were again observed (Figure 9c). Addition of 1 mg or 3 mg of Si02 did not affect significantly either the quantity or the crystal habit of the hematite formed in 6 d at 150°C. 6.4.3 Experiments with Phosphate up to 25O°C Further experiments on the inhibition of dissolution-precipitation processes by phosphate were extended to higher temperatures, with excess oxygen. In the most extreme case, magnetite A (0.5 g) was heated in a solution of 3 1 3 3 Na2HP04 (10- mol.L" , 100 cm ) in an air-filled 420-cm autoclave for 10 d at 250°C. In this case, the magnetite microstructure was retained, as shown in Figure 11, but XRD indicated complete conversion to hematite. Thus, whereas the phosphate still strongly inhibits the solution-mediated transformation at this temperature, it does not prevent the solid-state oxidation and transformation of magnetite, via maghemite, to hematite. The solid-state conversion of maghemite to hematite, however, is very slow below 200°C. At such temperatures, maghemite can be regarded as either a metastable or long-lived transient phase. Gallagher et al. (1968), for example, state that maghemite remains unchanged indefinitely below 220°C, although they have observed partial solid-state conversion of magnetite to hematite at temperatures as low as 180°C. Schwertmann and Cornell (1991) recommend a temperature of 250°C (for 2 h) for synthesis of maghemite from either lepidocrocite or magnetite. - 13 -

6.5 EXPERIMENTS BELOW 1OO°C Magnetite oxidation experiments were run for up to 20 d at 75CC and 60 d at both 22 and 50°C, using both magnetite A and B, in both pure water and 1 mol-L"1 NaCl. Additional experiments were run with magnetite A for 6 d and 20 d in water at 70, 80 and 90°C. Hematite was barely detectable by XRD (based on a single weak peak at d = 0.270 nm) on magnetite A after 6 d at 90°C, 20 d at 80°C, and 60 d at 50°C; maghemite was detected on magnetite B after 20 d at 75°C. No solution-grown reaction products could be clearly identified by SEM in any of these experiments. Thus, the combination of SEM and XRD is not sufficiently sensitive to monitor magnetite oxidation at temperatures much below 100°C, except perhaps for very long experiments, extending from months to years. 6.6 POSSIBLE INVOLVEMENT OF FLUORIDE Because the magnetite oxidation experiments were performed at elevated temperatures in PTFE-lined vessels, it is possible that fluoride, leached from the PTFE, influenced the dissolution/precipitation processes. Fluor- ide can substitute for OH in oxyhydroxides and on hydroxylated oxide surfaces, and it also forms strong complexes with Fe3+. As part of a previous investigation of cesium aluminosilicate phase relationships, using the same vessels, we measured fluoride concentrations up to 10'5 mol-L"1 after 6 to 10 d at 200°C, in solutions with initial NaOH concentrations up to 10'3 mol.L"1 (Taylor et al. 1989). Because most of our experiments with magnetite were conducted at 150°C, and the reactions were unaffected even by phosphate at a concentration of 10'4 mol-L*1, it is unlikely that the reactions were significantly affected by released fluoride. The difficulty of sustaining dilute hydroxide solutions (-10-4 mol.L"1) is usually a more significant problem than fluoride release per se in experiments of this kind.

7. CONCLUSIONS

Comparison of the solubilities of different Fe(III) oxide phases provides a means for comparing the transient redox conditions that may be established by these phases in association with magnetite. These range from moderately oxidizing for some of the least stable, amorphous Fe(III) oxides to strongly reducing for crystalline hematite. Therefore, the nature of the Fe(III) oxide phase - both its mineralogy and its crystallinity - must be considered when estimating magnetite/Fe(III) oxide redox buffer conditions. Experimental oxidation of synthetic magnetite powders has shown that, in pure water or solutions of salts that do not form strong surface complexes with iron oxides, magnetite is oxidized in periods of days to weeks at 100cC, with well-crystallized hematite as a major product. Thus, redox and solubility equilibria approaching the theoretical magnetite/hematite buffer should be attained rather quickly under such conditions. Dissolved silica or phosphate, however, strongly inhibit the formation of hematite by a dissolution-precipitation route, and permit the formation and persistence of maghemite at temperatures up to 200°C. Further solid-state transformation of maghemite to hematite has not been observed below 200°C, so metastable magnetite/maghemite redox equilibria may persist for significant periods in the presence of dissolved silica or phosphate. Other metastable Fe(III) oxides may become important as magnetite oxidation products at T £ 100-C, but the combination of XRD and SEM is inadequate to investigate such reactions, at least for experimental durations up to 20 d at 75°C or 60 d at 50°C.

ACKNOWLEDGEMENTS

We thank J.J. Cramer, R. Dayal, L.H. Johnson, R.J. Lemire, D.S. Mancey, H.W. Nesbitt and D.W. Shoesmith for helpful discussions. The Canadian Nuclear Fuel Waste Management Program is jointly funded by AECL and Ontario Hydro under the auspices of the CANDU* Owners Group.

REFERENCES

Baes, CF. Jr. and R.E. Mesmer. 1976. The Hydrolysis of Cations. John Wiley & Sons, New York.

Berner, R.A. 1969. Goethite stability and the origin of red beds. Geochimica et Cosmochimica Acta 33, 267-273.

Blesa, M.A. and E. Matijevic. 1989. Phase transformations of iron oxides, oxohydroxides, and hydrous oxides in aqueous media. Advances in Colloid and Interface Science 29, 173-221.

Booy, M. and T.W. Swaddle. 1978. Hydrothermal preparation of magnetite from iron chelates. Canadian Journal of Chemistry 56, 402-403.

Borggaard, O.K. 1983. Effect of surface area and mineralogy of iron oxides on their surface charge and anion-adsorption properties. Clays and Clay Minerals 31. i 230-232.

Bray, W.C and A.V. Hershey. 1934. The hydrolysis of ferric ion. The standard potential of the ferric-ferrous electrode at 25°• The equilibrium Fe+++ + Cl" = FeCl++. Journal of the American Chemical Society 56, 1889-1893.

Cornell, R.M., R. Giovanoli, and P.W. Schindler. 1987. Effect of silicate species on the transformation of ferrihydrite into goethite and hematite in alkaline media. Clays and Clay Minerals 35, 21-28.

Cornell, R.M., R. Giovanoli and W. Schneider. 1989. Review of the hydrolysis of iron(III) and the crystallization of amorphous iron(III) hydroxide hydrate. Journal of Chemical Technology and Biotechnology, 46, 115-134.

CANDU (ÇANada Deuterium Uranium) is a registered trademark of AECL - 15 -

Davis, B.L., G. Rapp Jr. and M.J. Walavender. 1968. Fabric and structural characteristics of the martitization process. American Journal of Science 2.66, 482-496. Doyle, R.W. 1968. The origin of the ferrous ion-ferric oxide Nernst potential in environments containing dissolved ferrous iron. American Journal of Science 266, 840-859.

Egger, K. and W. Feitknecht. 1962. Oxidation of Fe304 to 7- and or-Fe2O3. Differential thermal analysis (DTA) and thermogravimetric (TG) determination of the reaction rate of synthetic forms of Fe3O4. Helvetica Chimica Acta 45, 2042-2057. Ellis, J., R. Giovanoli and W. Stumra. 1976. Anion exchange properties of 0-FeOOH. Chimica 30, 194-197. Ferrier, A. 1966. Influence of the state of division of goethite and ferric oxide on their heats of reaction. Revue de Chimie minérale 3, 587-615. Fox, L.E. 1988. Solubility of colloidal ferric hydroxide. Nature 333, 442-444. Gallagher, K.J., W. Feitknecht and U. Mannweiler. 1968. Mechanism of oxidation of magnetite to -y-Fe2O3. Nature 217, 1118-1121. Gillot, B., A. Rousset, J. Paris and P. Barret. 1977. Oxidation kinetics study of finely divided magnetites substituted for aluminum and chromium. In Proceedings of the 8th International Symposium on the Reactivity of Solids. Plenum, New York, pp. 125-129. Hancox, W.T. and K. Nuttall. 1991. The Canadian approach to safe permanent disposal of nuclear fuel waste. Nuclear Engineering and Design 129, 109-117. Harrison, J.B. and V.E. Berkheiser. 1982. Anion interactions with freshly prepared hydrous iron oxides. Clays and Clay Minerals 30, 97-102. Hemingway, B.S. 1990. Thermodynamic properties for bunsenite, NiO, magnetite, Fe3O4, and hematite, Fe203, with comments on selected oxygen buffer reactions. American Mineralogist 75, 781-790. Hemingway, B.S., R.A. Robie and J.A. Apps. 1991. Revised values for the thermodynamic properties of boehmite, AlO(OH), and related species and phases in the system Al-H-0. American Mineralogist Z6, 445-457. Hingston, F.J., A.M. Posner and J.P. Quirk. 1972. Anion adsorption by goethite and gibbsite. I. The role of the proton in determining adsorption envelopes. Journal of Soil Science 23, 177-192. Hsu, P.H. and G. Marion. 1985. The solubility product of goethite. Soil Science HO, 344-351. - 16 -

Jolivet, J.-P. and E. Tronc. 1988. Interfacial electron transfer in colloidal spinel iron oxide. Conversion of Fe3O4—y-Fe2O3 in aqueous medium. Journal of Colloid and Interface Science 125» 688-701.

Jolivet, J.P., E. Tronc, C. Barbe and J. Livage. 1990. Interfacial electron transfer in colloidal spinel iron oxide. Silver ion reduction in aqueous medium. Journal of Colloid and Interface Science 138. 465-472.

King, F., B.M. Ikeda and D.tf. Shoesmith. 1992. Nuclear waste: Can ve contain it? Chemtech 214-219.

Kolb, E.D., A.J. Caporaso and R.A. Laudise. 1973. Hydrothermal growth of hematite and magnetite. Journal of Crystal Growth 19, 242-246.

Langmuir, D. 1969. The Gibbs free energies of substances in the system Fe-02-H20-C02 at 25°C. U.S. Geological Survey Professional Paper 650-B, B180-B184.

Langmuir, D. 1971. Particle size effect on the reaction goethite = hematite + water. American Journal of Science 271, 147-156 (and correction (1972)), ibid. 272, 972.

Langmuir, D. and D.O. Whittemore. 1971. Variations in the stability of precipitated ferric oxyhydroxides. In Nonequilibrium Systems in Natural Water Chemistry (J.D. Hem, Symposium Chairman), Advances in Chemistry Series, American Chemical Society, Washington, D.C., 106, 209-234.

Larson, J.W., P.O. Cerutti, H.K. Garber and L.G. Hepler. 1968. Electrode potentials and thermodynamic data for aqueous ions. Copper, zinc, cadmium, iron, cobalt, and nickel. Journal of Physical Chemistry 72, 2902-2907.

Latimer, W.M. 1952. The Oxidation States of the Elements and Their Potentials in Aqueous Solutions, 2nd edition. Prentice-Hall, Inc., Englewood Cliffs, NJ, p. 223.

Lemire, R.J. and F. Garisto. 1989. The solubility of U, Np, Pu, Th and Tc in a geological disposal vault for used nuclear fuel. Atomic Energy of Canada Limited Report, AECL-10009.

Lindsay, W.L. 1979. Chemical Equilibria in Soils. John Wiley & Sons, New York, Chapter 10.

Lindsay, W.L. 1988. Solubility and redox equilibria of iron compounds in soils. In Iron in Soils and Clay Minerals (J.V. Stucki et al., editors). D. Reidel Publishing Company, Dordrecht, Holland, Chapter 3.

Morris, M.C., H.F. McMurdie, E.H. Evans, B. Paretzkin, H.S. Parker, N.C. Panagiotopoulos and C.R. Hubbard. 1981. Standard X-ray Diffraction Powder Patterns, U.S. National Bureau of Standards Monograph 25, Section 18, p. 37 (JCPDS Card 33-664). - 17 -

O'Neill, H. St. C. 1988. Systems Fe-0 and Cu-O: Thermodynamic data for the equilibria Fe-"FeO", Fe-Fe3O4, "FeON-FejO,, Fe304-Fe203, Cu-Cu20, and Cu2O-CuO from emf measurements. American Mineralogist 73, 470-486.

Palache, C, H. Berman and C. Frondel. 1944. The System of Mineralogy of James Dvight Dana and Edward Salisbury Dana. Three volumes. 7th edition John Wiley, New York, NY, p. 528 (hematite) and p. 699 (magnetite).

Palmer, D.A. and D.J. Wesolowski. 1992. Aluminum speciation and equilibria in aqueous solution: II. The solubility of gibbsite in acidic sodium chloride solutions from 30 to 70°C. Geochimica et Cosmochimica Acta 56, 1093-1111.

Paquette, J. and R.J. Lemire. 1981. A description of the chemistry of aqueous solutions of uranium and plutonium to 200°C using potential-pH diagrams. Nuclear Science and Engineering 79, 26-48.

Parfitt, R.L., R.J. Atkinson and R. St. C. Smart. 1975. The mechanism of phosphate fixation by iron oxides. In Proceedings of the Soil Science Society of America 39, 837-841.

Robie, R.A., B.S. Hemingway and J.R. Fisher. 1979. Thermodynamic Properties of Minerals and Related Substances at 298.15 K and 1 bar (105 pascals) Pressure and at Higher Temperatures. U.S. Geological Survey Bulletin 1452 (reprinted with corrections), U.S. Govt. Printing Office, Washington, D.C., 456 pp.

Sadiq, M. and W.L. Lindsay. 1979. Selection of standard free energies of formation for use in soil chemistry. Colorado State University Experiment Station Technical Bulletin 134.

Sadiq, M. and W.L. Lindsay. 1981. Standard free energy of formation of some chemical species: discrepancies and selections. Arabian Journal for Science and Engineering 6, 95-104.

Sadiq, M. and W.L. Lindsay. 1988. The solubility product of soil maghemite. Soil Science 146, 1-5.

Sapieszko, R.S. and E. Matijevic. 1980. Preparation of well-defined colloidal particles by thermal decomposition of metal chelates. I. Iron oxides. Journal of Colloid and Interface Science 74, 405-422.

Schulz, D.L. and G.J. McCarthy. 1988. X-ray powder data for an industrial maghemite (7-Fe2O3). Powder Diffraction 3, 104-105.

Schumb, W.C., M.S. Sherrill and S.B. Sweetser. 1937. The measurement of the molal ferric-ferrous electrode potential. Journal of the American Chemical Society 59, 2360-2365.

Schwab, A.P. and W.L. Lindsay. 1983. Effect of redox on the solubility and availability of iron. Soil Science Society of America Journal 47, 201-205. - 18 -

Schwertmann, U. and R.M. Cornell. 1991. Iron Oxides in the Laboratory: Preparation and Characterization, VCH, Weinheim, Germany, 137 pp. Shoesmith, D.W. and S. Sunder. 1991. An electrochemistry-based model for the dissolution of U02. Atomic Energy of Canada Limited Report, AECL-10488. Sidhu, P.S., R.J. Gilkes and A.M. Posner. 1977. Mechanism of the low temperature oxidation of synthetic magnetites. Journal of Inorganic and Nuclear Chemistry 39, 1953-1958. Stockbridge, CD., P.B. Sewell and M. Cohen. 1961. Cathodic behavior of iron single crystals and the oxides Fe3O4, gamma-Fe2O3, and alpha- Fe203. Journal of the Electrochemical Society 108, 928-933. Swaddle, T.V. and P. Oltmann. 1980. Kinetics of the magnet!te-maghemite- hematite transformation, with special reference to hydrothermal systems. Canadian Journal of Chemistry 58, 1763-1772. Swanson, H.E., H.F. McMurdie, M.C. Morris and E.H. Evans. 1967. Standard X-ray Diffraction Powder Patterns, U.S. National Bureau of Standards Monograph 25, Section 5, p. 31 (JCPDS Card 19-629). Taylor, P., S.D. DeVaal and D.G. Owen. 1989. Stability relationships between solid cesium aluminosilicates in aqueous solutions at 200°C. Canadian Journal of Chemistry 67, 76-81. Taylor, P., T.E. Rummery and D.G. Owen. 1979. Reactions of iron monosulfide solids with aqueous hydrogen sulfide up to 160°C. Journal of Inorganic and Nuclear Chemistry 41, 1683-1687. Tremaine, P.R. and J.C. LeBlanc. 1980. The solubility of magnetite and the hydrolysis and oxidation of Fe2+ in water at 3008C. Journal of Solution Chemistry 9, 415-442. Tremaine, P.R., R. von Massow and G.R. Shierman. 1977. A calculation of Gibbs free energies for ferrous ions and the solubility of magnetite in H20 and D20 to 300°C. Thermochimica Acta 19, 287-300. Vempati, R.K. and R.H. Loeppert. 1989. Influence of structural and adsorbed Si on the transformation of synthetic ferrihydrite. Clays and Clay Minerals 37, 273-279. Wagman, D.D., W.H. Evans, V.B. Parker, R.H. Schumm, I. Halow, S.M. Bailey, K.L. Churney and R.L. Nuttall. 1982. The NBS tables of chemical thermodynamic properties. Selected values for inorganic and Cx and C2 organic substances in SI units. Journal of Physical and Chemical Reference Data H, Supplement 2 (392 pp.) Wesolowski, D.J. 1992. Aluminum speciation and equilibria in aqueous solution: I. The solubility of gibbsite in the system Na-K-Cl-0H- Al(0H)4 from 0 to 100°C. Geochimica et Cosmochimica Acta 56., 1065-1091. - 19 -

White, A.F. 1990. Heterogeneous electrochemical reactions associated with oxidation of ferrous oxide and silicate surfaces. Reviews in Mineralogy 23, 467-509. Zachara, J.M., D.C. Girvin, R.L. Schmidt and C.T. Resch. 1987. Chromate adsorption on amorphous iron oxyhydroxide in the presence of major groundwater ioiu». Environmental Science and Technology 21., 589-594.

TABLE 1 GIBBS ENERGIES OF FORMATION (25"Cï USED IN THIS PAPER

a Substance A£G°, kJ/mol Source

Fe304(s) -1012.6 ±2.1 Robie et al. (1979); Hemingway (1990) a-Fe2O3(s) -744.3 ±1.3 Hemingway (1990) Fe2+(aq) -91.2 ±2.1 Larson et al. (1968); see text Fe3+(aq) -16.8 ±2.2b Larson et al. (1968); see text OH-(aq) -157.33±O.O9 Robie et al. (1979)

H20(l) -237.14±0.08 Robie et al. (1979) a Rounded values b Assuming minimal uncertainty (<5 mV) in E° for the Fe2+(aq)/Fe3+(aq) couple (Latimer 1952, Bray and Hershey 1934, Schumb et al. 1937) FIGURE 1:Calculated EquilibriumConstantsat25°CforReaction s (4)(p(0Scaleand(5(E° 2 5 S I—1 o d) o bar] -20 -80 - -40 -60 - horizontal dashe d linesrepresen t thestabilitylimit s ofwateratP=1 bar(0.1MPa). point labelle d Hrepresentsequilibriu m betweenmagnetit e andcrystallin hematite;the Scale), asFunctiono f KAccordingtoEquations(11)and(12 ) Respectively.The spf p KofFe(III)oxid sp ^- 0.0 - 1.0 - 1.2 - 0.2 -0.2 0.4 0.8 0.6 I o a. Cfl oltsi SHE at i o • i - 21 -

CM 0) u.

FIGURE 2: Calculated Relationships Between {Fe2+}, pH and K at 256C for Magnetite/Fe(III) Oxide Equilibria According to Equation (14) - 22 -

FIGURE 3: Scanning Electron Micrographs of Magnetite A: (a) Low- Magnification, Showing Typical "Briquette" Particle Shapes; (b) High-Magnification, Showing Microstructure of One of the More Porous Particles - 23 -

FIGURE 4: Scanning Electron Micrographs of Products Obtained from 3 3 Oxidation of Magnetite A (150°C, 6 d, 5 cm H20, 20-cm Vessel): (a) Monodisperse Hematite Precipitate with a Few Larger Crystals; (b) Aggregate of Large, Solution-Grown Hematite Crystals FIGURE 5: Scanning Electron Micrographs of Hematite Crystals: (a) Product of "Fe(OH)3" Precipitated 1 from Fe(N03)3 and Excess NaOH, Ripened at 200°C for 6 d in 1 mol.L" NaOH; (b) and (c) Product from Oxidation of Magnetite A, Duplicate Experiment Under Same Conditions as Figure 4 - 25 -

FIGURE 6: Twinned Lath-Shaped Crystals, Possibly Goethite, Obtained in Trace Quantities in One Experiment with Magnetite A (150°C, 6 3 3 5 cm H2O, 20-cm Vessel). - 26 -

FIGURE 7: Scanning Electron Micrographs of Products Obtained by Oxidation of Magnetite A (150°C, 6 d, 5 cm3 Solution, 20-cm3 Vessel) in the Following Solutions: (a) 0.1 mol-L"1 NaCl (Product Obtained 1 1 in 0.1 mol-L- Na2S04 tfas Similar); (b) 0.01 mol-L" NaHC03 1 (Products Obtained in 0.01 mol-L- Na2B407, Na2C03, and NaOQ Were Similar) - 27 -

FIGURE 8: Scanning Electron Micrographs of Two Different Portions of the Product of Oxidation/Dissolution of Magnetite A in 0.01 mol-L"1 3 3 H2S04 (5 cm Solution in 20-cm Vessel, 150°C, 6 d) S3 oo

FIGURE 9: Scanning Electron Micrographs of Products Obtained by Oxidation of Magnetite A (150°C, 3 3 2 1 6 d, 5 cm Solution, 20-cm Vessel) in the Following Solutions: (a) 10- mol-L" Na2HP04 4 1 (b) 10- mol.L- Na2BP04; (c) Water Plus 10 mg Powdered Vitreous SiO2 - 29 -

c o

40 41 42 43 44 45

Diffraction angle (26), degrees

FIGURE 10: Expanded Portions of X-ray Diffraction Patterns of Magnetite Oxidation Products (5 cm3 Solution, 20-cm3 Vessel, 150°C, 6 d). (a) Deionized waterj note sharp magnetite 400 peak and strong hematite 113 peak. Slight asymmetry of peaks is due to the 3 1 Cu Ka X-ray doublet, (b) 10- mol-L" Na2HPO4. Note additional high-angle shoulder on magnetite peak and greatly reduced hematite peak. All other XRD peaks in the patterns were consistent with these features. - 30 -

FIGURE 11: Scanning Electron Micrograph of Product, Identified by XRD as Pure Hematite, Obtained by Oxidation of Magnetite A (250°C, 3 3 10 d, 100 cm 10-3 moi.L-i Na2HP04, 420-cm Vessel) AECL-10821 COG-93-81

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