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Langmuir 1998, 14, 1250-1255
Crystal Growth of Pyrite in Aqueous Solutions. Inhibition by Organophosphorus Compounds N. G. Harmandas, E. Navarro Fernandez, and P. G. Koutsoukos* Institute of Chemical Engineering and High Temperature Chemical Processes, ICE/HT-FORTH and Department of Chemical Engineering, University of Patras, University Campus, GR 26500, Patras, Greece Received April 7, 1997. In Final Form: November 28, 1997 The crystal growth of pyrite in aqueous supersaturated solutions at pH 6.50, 25 °C was investigated using the seeded growth, pH-stat technique. Crystal growth started immediately after inoculation of iron(II) sulfide supersaturated solutions with pyrite seed crystals. The crystal growth rates measured from the desupersaturation curves were found to depend strongly on the relative solution supersaturation with respect to pyrite. Kinetics analysis yielded an apparent order of 3.5 ( 0.5 suggesting a surface diffusion controlled mechanism. The presence of organophosphorus compounds such as 1-hydroxyethylidene1,1-diphosphonic acid (HEDP), N,N,N′,N′-ethylenediaminetetrakis(methylenephosphonic acid) (ENTMP) in the supersaturated solutions at concentration levels between 0.1 and 1000 nM inhibited the crystal growth of pyrite by 70-100% with respect to the blank experiments. Moreover, in the presence of the compounds tested the crystal growth of pyrite was preceded by well-defined induction times which increased sharply with inhibitor concentration. At concentration levels as low as 1 mM, crystal growth was completely suppressed. The presence of nitrilotris(methylenephosphonic acid) (NTMP) had no inhibition effect on pyrite crystal growth. The influence of the inhibitors tested on the kinetics of crystal growth was attributed to the blocking of the active growth sites due to adsorption.
Introduction Pyrite is the most abundant iron sulfide on the Earth’s surface, encountered in either ancient or modern sediments,1,2 and is of great geological importance as it is involved in the geochemical cycles of sulfur and iron. Pyrite along with other iron sulfides may also be found in the insoluble scales forming on the carbon steel pipe walls used for the transport of the natural hot brines in geothermal installations. Natural geothermal fluid may contain large amounts of dissolved H2S, a fact that makes it a very corrosive aqueous medium.3 Due to corrosion, Fe2+ ions are released in the bulk fluid creating areas supersaturated with respect to iron(II) sulfides. In this case iron sulfide precipitation may occur as a result of the presence of the dissolved sulfide. Understanding the mechanism of pyrite formation is a prerequisite to any attempt of prevention and/or elimination of undesirable scales of this type. Despite pyrite’s abundance and numerous studies carried out by several researchers,2,4-10 the mechanism of pyrite formation is not yet fully understood, mainly because of the experimental difficulties involved.10 In the present work we have attempted to approach the problem of iron sulfide formation quantitatively through the formulation of the driving force as a function of the solution supersaturation. More specifically we have studied the kinetics of pyrite precipitation in stable aqueous supersaturated solutions (1) Schoonen, M. A. A.; Barnes, H. L. Geochim. Cosmochim. Acta 1991, 55, 1505. (2) Berner, R. A. Geochim. Cosmochim. Acta 1983, 48, 605. (3) Tewari, P. H.; Campbell, A. B. Can. J. Chem. 1979, 57, 188. (4) Sweeney, R. E.; Kaplan, I. R. Econ. Geol. 1973, 68, 618. (5) Schoonen, M. A. A.; Barnes, H. L. Geochim. Cosmochim. Acta 1991, 55, 1495. (6) Schoonen, M. A. A.; Barnes, H. L. Geochim. Cosmochim. Acta 1991, 55, 3491. (7) Rickard, D. T. Am. J. Sci. 1975, 275, 636. (8) Luther, J. W., III Geoch. Cosmoch. Acta 1991, 55, 2839. (9) Berner, R. A. J. Geol. 1964, 72, 293. (10) Schoonen, M. A. A. Ph.D. Thesis, Penn State University, 1989.
of high ionic strength at low temperature, using the seeded growth technique.11 The choice of pyrite was done because it is the least soluble and hence the thermodynamically most stable iron sulfide. We have also evaluated the possibility of retarding or preventing pyrite precipitation with the addition of three organophosphorus compounds in the working solutions. These additives were HEDP (1-hydroxyethylidene-1,1-diphosphonic acid), ENTMP (N,N,N′,N′-ethylenediaminetetrakis(methylenephosphonic acid), and NTMP (Nitrilotris(methylenephosphonic acid), products commercially available under various trademarks. These types of phosphonate compounds have been proven to be effective inhibitors of calcium carbonate and sulfate precipitation and have the advantage of enhanced resistance to hydrolytic decomposition even at elevated temperatures.12-16 Experimental Section All experiments were done in a double-walled water-jacketed reactor vessel volume totaling 250 mL. The temperature was kept constant at 25.0 ( 0.1 °C by circulating water from a constant temperature water bath. A magnetic stirrer was used for the stirring of the working solution. The mixing of the solutions in the reactor was done through an opening on the reactor lid, which was sealed airtight immediately after the mixing. To ensure an inert atmosphere above the working solution, a slight overpressure of water vapor saturated nitrogen gas (99.99%, Linde) was maintained. Maintenance of the inert atmosphere precluded sulfide or pyrite surface oxidation. Degassing of the water used for the solution preparation in combination with the slight positive nitrogen overpressure precluded oxidation of pyrite according to (11) Nancollas, G. H.; Mohan, M. S. Archs. Oral Biol. 1970, 731. (12) Klepetsanis, P. G. Ph.D. Thesis, Department Chemical Engineering, Patras University, 1991. (13) Klepetsanis, P. G.; Koutsoukos, P. G. In Proceedings of 11th Symposium Industrial Crystallization, Mersmann, A., Ed.; 1990; p 261. (14) Weijnen, M. Ph.D. Thesis, Delft Technical University, 1986. (15) Dalas, E.; Koutsoukos, P. G. Desalination 1990, 78, 403. (16) Xyla, A. G.; Koutsoukos, P. G. J. Chem. Soc., Faraday Trans. 1 1987, 83, 1477.
S0743-7463(97)00354-5 CCC: $15.00 © 1998 American Chemical Society Published on Web 02/11/1998
Crystal Growth of Pyrite in Aqueous Solutions
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Figure 1. Experimental setup for the FeS nucleation and growth experiments at 25 °C. the reactions which take place in the presence of oxygen at oxidizing conditions17
FeS2 + O2 T FeS2‚O2
(1)
FeS2 + H2O T FeS2‚(OH) + H+ + e-
(2)
pH stability in the supersaturated solutions prior to seeding confirmed lack of oxidation processes. A combination glass|SCE electrode was used for monitoring the working solution pH throughout each experiment. An automatic system for monitoring and control of precipitation processes was utilized, consisting of a personal computer with an expansion card providing access to an external board called “A-Bus”. Specific software was written for the control of experimental variables, electrode calibration, performance of pHstat operations and data file manipulation. A schematical presentation of the experimental set up may be seen in Figure 1. Two different experimental approaches were followed for the crystal growth experiments at 25 °C: (i) Variable pH (Free-Drift) Experiments. Following the preparation of the supersaturated solutions, the automatic system recorded the pH value of the working solution as a function of time. This method was used only in preliminary experiments done in order to define the width of the metastability zone of the (NH4)2Fe(SO4)2-H2S-(NH4)2SO4-H2O system. (ii) Experiments at Constant pH. After the preparation of the supersaturated solutions, the computer controlled automatic system was allowed to maintain the solution pH constant at a set point value (pH ≈ 6.5 for our experiments). In these experiments, the titrator syringes were filled with 0.01 N KOH solution. A decrease in the working solution pH triggered the addition of KOH solution contained in the syringes to keep pH constant. Before each experiment, fresh stock solutions of Fe2+ and S2ions were prepared with triply distilled water, deaerated with N2 gas bubbling for at least 30 min. Mohr’s salt [(NH4)2Fe(SO4)2‚6H2O] was used for the preparation of the iron stock and standard solutions and NaSH‚xH2O for the sulfide stock solution. A small amount of 0.1 N H2SO4 solution (5 mL) was also added in the iron stock solution so as to avoid Fe(OH)x formation which may occur in alkaline environments. Sulfide solutions were standardized by titration with Hg(NO3)2 using dithizone as indicator.18 Crystalline (NH4)2SO4 was used for the preparation of the inert electrolyte stock solution utilized for the adjustment of the solution ionic strength conditions. In our experiments I ) 0.4 M. All working solutions were supersaturated with respect to pyrite but undersaturated with respect to monosulfides. Equal volumes (100 mL each) of iron(II) and sulfide solutions were directly mixed into the reactor vessel under vigorous stirring to reach the desired supersaturation level at the start of each experiment. (17) Fornasiero, V. E.; Eijt, V.; Ralston, J. Colloids Surf. 1992, 62, 63. (18) Karabash, A. G. Zh. Anal. Khim. 1953, 8, 140.
Figure 2. Typical profile of the decrease of total iron concentration with time; 25 °C, pH ) 6.5; initial total iron Fet ) 2.85 × 10-4; initial total sulfide St ) 3.3 × 10-4; No additives present. Dotted line is the tangent to the initial part of the curve, for the calculation of initial rates. The precipitation of FeS is accompanied by the release of protons according to the reaction:
xFe2+ + yH2S + zHS- + uS2- T nFeS(s) + mH+
(3)
where 2x - m ) z + 2u. According to eq 3 the onset of iron sulfide precipitation is accompanied by a decrease in the pH value of the working solution. The evolution of the precipitation reaction may therefore be monitored by the progressive decrease of pH. Thus the master variable chosen for monitoring and controlling the precipitation reaction was the hydrogen ion activity (pH). The combination glass/SCE electrode used was suitable for use in sulfide-containing solutions and was calibrated before and after each experiment with NBS buffer solutions.19 In the preliminary experiments done, the region of metastability was defined as follows: Past the preparation of the solutions supersaturated with respect to pyrite, the solution pH was monitored. If no change was observed for at least 8 h, the solution was considered to be stable. Stable supersaturated solutions were thus prepared and were allowed to attain thermal equilibrium for a few minutes. Next, the solution pH was adjusted to the value of 6.50 and an accurately weighed quantity (ca. 10 mg) of synthetic pyrite seed crystals was added. Decrease of the solution pH by 0.003 pH unit, triggered the addition of standard KOH solution to compensate the pH drop by neutralization of the released protons. During each experiment, the automatic system continuously recorded the solution pH value and the volume of titrants added as a function of time. Small aliquots of the working solution were withdrawn at regular time intervals (more frequently at the early stages of each experiment). The samples were filtered through membrane filters (0.2 µm), and the filtrates were kept tightly sealed for subsequent analyses for iron (by atomic absorption spectroscopy) and sulfide (spectrophotometrically). The profiles of iron concentration decrease as a function of time were used to calculate the precipitation rates using the initial rates, obtained from nonlinear fitting of the experimental points. A typical profile of total iron decrease with time is shown in Figure 2. Past sufficiently long time periods (when no pH decrease was observed and no further alkali additions were done) the experiments were stopped. It should be noted that the duration of the experiments was on the average 4 h depending on the supersaturation. The solids were collected on 0.2 µm membrane filters, dried at 25 °C in a desiccator under an inert, nitrogen atmosphere. After drying the crystals were characterized with powder X-ray diffraction (Philips PW 1840) using Cu KR radiation and scanning electron microscopy (JEOL JSM 5200). (19) Bates, R. G. Determination of pH. Theory and Practice; Wiley-Interscience: New York, 1973.
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Preparation of Pyrite Seed Crystals. The seed crystals used in our experiments were obtained by preparing a suspension of amorphous iron sulfide and refluxing it for 20 days at 95 °C in a slightly acidic and oxidizing aqueous environment created by the addition of HClO4. During the ripening process the amorphous particles gradually converted to spherullitic crystals of about 1 mm in size. Following the ripening process and preliminary drying, the solid was thoroughly washed with 6 N HCl solution to remove possibly remaining monosulfides. Next the solid was dried and left under vacuum in a desiccator, previously purged with nitrogen. Periodic examination of the solid did not show any morphological bulk or surface changes. The morphology of the crystals obtained is shown in Figure 4a. Such a morphology for synthetic pyrite has been also reported by other researchers.4,20 The powder x-ray diffraction (XRD) spectra of the seed crystals matched the reported spectra for pyrite.21 The BET specific surface area of the seed crystals, determined by a multiple point dynamic nitrogen adsorption method, was found to be 6 m2 g-1.
Determination of the Metastable Zone and Seeded Crystal Growth Experiments. The driving force for the formation of pyrite in aqueous supersaturated solutions is the difference between the chemical potentials of the salt at equilibrium, µ∞, from the corresponding value in the supersaturated solution, µs:
(4)
and
RT ln ∆G ) 2
(Fe2+)(S22-) o Ks,FeS 2
(5)
In eq 5 the ratio
ΩFeS2 )
(Fe2+)(S22-) o Ks,FeS 2
(6)
is the supersaturation ratio with respect to pyrite, parentheses denote the activities of the corresponding ions o is the thermodynamic solubility product of and Ks,FeS2 pyrite (8.511 × 10-26 M2).22 The relative supersaturation with respect to pyrite, σFeS2 is defined as 1/2 -1 σFeS2 ) ΩFeS 2
Aqueous Equilibria for the Computation of Solution Speciation in (NH4)2Fe(SO4)2-H2S-(NH′4)2SO4-H2O Systems log Ko
equilibrium H2S(g) S H2S(aq) HS- + H+ S H2S (aq) S2- + H+ S HSNH3 + H+ S NH4+ SO42- + H+ S HSO4Fe2+ + SO42- S FeSO40 Fe2+ + OH- S FeOH+ Fe2+ + 2OH- S Fe(OH)20 Fe2+ + 3OH- S Fe(OH)3Fe2+ + 4OH- S Fe(OH)42Fe2+ + 2H+ + 2S2- S Fe(HS)2
(7)
The calculation of supersaturation was done using HYDRAQL,23 a speciation calculation program taking into account all chemical equilibria and the mass and charge balance of the specified aqueous system. The equilibria and the corresponding values for the stability constants used in the calculations are summarized in Table 1. The solution of the high-order simultaneous equations was done by minimizing the free energy of the system. Extended Debye-Hu¨ckel equations were used for the calculation of the activity coefficients. Computations of the relative supersaturation were done considering within the equilibria the presence of disulfide ions. Longer chain (20) Wang, Q.; Morse, J. W. Mar. Chem. 1996, 52, 99. (21) JCPDS Card No. 42-1340. (22) Lindsay, W. L. Chemical Equilibria in Soils; Wiley-Interscience: New York, 1979. (23) Papelis, C.; Hayes, K. F.; Leckie, J. O. HYDRAQL A program for the Computation Of Chemical Equilibrium Composition of Aqueous Batch Systems including Surface Complexation Modeling of Ion Adsorption at the Oxide/Solution Interface, Technol. Report No 306, Stanford University, 1988; p 130.
0.99 7.02 13.90 9.24 1.99 2.20 4.50 7.40 10.00 9.60 36.75
Thermodynamic Solubility Products of the Various Iron Sulfides Forming in Aqueous Media name
Results and Discussion
∆µ ) µs - µ∞
Table 1
formula
amorphous FeS troilite R-FeS pyrrhotite Fe0.98S mackinawite FeS greigite Fe3S4 pyrite FeS2 marcasite FeS2
expression
RFe2+ RS2RFe2+ RS2RFe2+ RS2RFe2+ RS2R3Fe2- R4S2RFe2+ RS22RFe2+ RS22-
1.44 × 10-17 6.17 × 10-17 2.70 × 10-19 2.88 × 10-18 2.99 × 10-55 8.51 × 10-26 8.65 × 10-26
Table 2. Experimental Conditions and Corresponding Rates of Precipitation Measured for the Experiments of Seeded Pyrite Growth in the Absence of Any Additive (Blank) expt no
Fet × 104 (M)
St × 104 (M)
σpyr
rate (mol m-2 s-1)
EF-L EF-A EF-B EF-D
1.5 2.0 2.5 3.0
1.65 2.20 2.75 3.30
1.12 × 107 1.6 × 107 2.0 × 107 2.4 × 107
2.3 × 10-8 4.5 × 10-8 1.5 × 10-7 2.1 × 10-7
polysulfides were excluded as their concentrations were not appreciable in our experimental conditions.24 Preliminary experiments were done in order to determine the supersaturation limit at which the supersaturated solutions remained stable for sufficiently long time periods. This limit was found to be at a supersaturation ratio of 5.7 × 10 14 ((10%) with respect to pyrite at 25 °C, pH 6.5. The rates of iron sulfide crystallization, Rp, showed a marked dependence on the solution supersaturation with respect to pyrite, σFeS2. The experiments for the study of crystal growth of iron sulfide on pyrite seed crystals were done in stable supersaturated solutions. The initial conditions for the experiments and the results obtained are summarized in Table 2. The powder X-ray diffraction spectra of the seed crystals before and after the crystal growth experiments, shown in Figure 3, showed that the phase grown on the seed crystals was exclusively pyrite. As may be seen in the scanning electron micrographs shown in Figure 4, the morphology of the grown material is strikingly different from that of the seed crystals, suggesting preferential growth in one direction leading to the formation of prismatic (needlelike) crystals. Needlelike morphology for pyrite forming at high temperatures (250-300 °C) and at relatively low supersaturations has also been reported.25 The rates of pyrite crystal growth on the pyrite seed crystals were measured as initial rates from the profiles of total iron-time curves. These curves (as, e.g., in Figure 2) showed a plateau indicating a slowdown of the precipitation process which continued afterward without (24) Murowchick, J. B. Ph.D. Thesis, Penn State University, 1984. (25) Murowchick, J. B.; Barnes, H. L. Am. Mineral. 1987, 72, 1241.
Crystal Growth of Pyrite in Aqueous Solutions
Figure 3. XRD patterns of the seed crystals and the solid precipitated in the blank experiments.
Langmuir, Vol. 14, No. 5, 1998 1253
Figure 5. Kinetics plot for seeded pyrite precipitation at 25 °C in the absence of additives, pH ) 6.5, I ) 0.4 M.
Figure 6. Molecular structure of the tested organophosphorus compounds.
aggregates could lead to the second part of the concentration-time profile, which is similar to the first part. The kinetics data were fitted in the semiempirical equation n RP ) ksσFeS 2
Figure 4. (a) Scanning electron micrograph of the pyrite seed crystals used in the experiments. (b) Scanning electron micrograph of the solid precipitated in the blank experiments
attaining a clear second plateau at least for the duration of the experiments reported here. Since the solutions employed were undersaturated with respect to monosulfides, the break in the concentration-time profiles did not correspond to any precursor or intermediate phase. Since however the pyrite particles have high negative values of ζ potential (ca. -40 mV),17 it is very likely that aggregation of the suspended seed crystals may take place (concentration coagulation),26,27 thus causing reduction of the available surface area with subsequent decrease of the active growth sites. Next, crystal growth of the
(8)
for σ >1. In eq 8 k is the crystal growth rate constant, s is a function of the active growth sites on the seed crystals, and n is the apparent order of the crystal growth process. The logarithmic plot of the crystal growth rate of pyrite as a function of the relative supersaturation according to eq 8 is shown in Figure 5. A value n ) 3.5 ( 0.5 was obtained from the linear regression of the log-log plots. This relatively high value suggested a surface diffusion controlled mechanism. It is interesting to note that similar values for the apparent order for precipitation have been reported for the spontaneous precipitation of pyrite in aqueous supersaturated solutions at 80 °C.28 Moreover, high apparent orders (>2) have been reported for a number of sparingly soluble salts the crystal growth of which is clearly controlled by surface diffusion.29 (26) Adamson, A. W. Physical Chemistry of Surfaces, 5th ed.; J. Wiley: New York 1990; p 531. (27) Voyutsky, S. Colloid Chemistry, MIR Publ: Moscow, 1978; pp 316-322. (28) Harmandas, N. G.; Klepetsanis, P. G.; Koutsoukos P. G. In 5th Semiannual Report, Joule II, EC Programme No JOU2-CT92-0108, Corrosion and Scaling in Geothermal Systems; Ignatiadis, I., Ed.; Orleans 1995.
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Table 3. Experimental Results for the Inhibitor Testing Experiments induction time τ (min)
expt no.
relative inhibition (%)
inhibitor concentration (M)
100 88 88 86
1 × 10-6 1 × 10-7 1 × 10-8 1 × 10-9
100 89 85 70
1 × 10-7 1 × 10-8 1 × 10-9 1 × 10-10
EF-E EF-14 EF-F EF-G
stable 200 170 130
1.5 × 10-7 ENTMP N/A 1.75 × 10-8 1.80 × 10-8 2.15 × 10-8
EF-H EF-I EF-JB EF-K
stable 150 105 80
HEDP N/A 1.60 × 10-8 2.25 × 10-8 4.75 × 10-8
EF-11
0
NTMP 1.95 × 10-7
EF-B (blank)
∼0
rate (mol m-2 s-1)
1 × 10-5
Inhibition of Pyrite Crystal Growth. All inhibitor testing experiments were done at the same initial supersaturation ratio with respect to pyrite (ΩPYR ) 9.55 × 1013) and over inhibitor concentration range between 0.1 and 1000 nM. The organophosphorus compounds ENTMP, HEDP, and NTMP shown in Figure 6 were used. Inhibitor concentrations down to 0.1 nM retarded considerably the crystal growth of pyrite. One of the organophosphorus compounds tested (NTMP) did not show any inhibiting function. On the contrary, it seemed to accelerate precipitation. The results obtained for the kinetics of pyrite crystal growth in the presence of the organophosphorus inhibitors are summarized in Table 3. The relative inhibition was calculated from eq 9
relative inhibition )
R o - Ri × 100 Ro
Figure 7. Induction times preceding the onset of pyrite formation measured as a function of inhibitor concentration. Dotted lines correspond to stable supersaturated solutions for time periods exceeding 48 h.
(9)
where Ro and Ri are the crystal growth rates measured without and in the presence of inhibitors, respectively. It should be noted that crystal growth was preceded by an induction period, corresponding to the formation of the critical nucleus. The induction times and the precipitation rates measured were affected markedly by the concentration level of the inhibitors. The induction times measured increased with the inhibitor concentration reflecting a progressively increased difficulty in the formation of critical nucleus. The crystal growth rates measured were inversely proportional to the induction time. Plots of the induction times as a function of the inhibitor concentration for HEDP and ENTMP for the formation of pyrite are shown in Figure 7. The morphology of the solids resulting from these experiments may be seen in the micrographs shown in Figure 8. In the case of ENTMP the solid grown on the pyrite seed crystals was similar to that found in the blank experiments. The crystals, however, obtained in the presence of HEDP were markedly different from the morphological point of view. As may be seen, the presence of this compound suppressed the growth of the newly formed crystallites toward a direction perpendicular to the seed crystal surface (needlelike morphology) resulting in granular crystallites of smaller size. Habit modification and crystal size reduction by HEDP have also been reported for gypsum.12,14 According to the results of the present study, the seeded growth of pyrite is a surface (29) So¨hnel, O.; Garside, J. Precipitation; Butterworth-Heinemann: Oxford 1992; p 95 and pp 307-352.
Figure 8. (a) Scanning electron micrograph of the solid precipitated in experiments in the presence of ENTMP. (b) Scanning electron micrograph of the solid precipitated in experiments in the presence of HEDP.
diffusion controlled process. Thus, the inhibition effectiveness of the compounds tested may be attributed to their adsorption on the seed crystal surface and subsequent blocking of the active crystal growth sites. HEDP and ENTMP molecules have longer chains and a linear configuration as compared with NTMP which is of tetrahedral geometry. It may therefore be speculated that the first two compounds may approach the crystal surface more easily and block more active sites than the third one
Crystal Growth of Pyrite in Aqueous Solutions
which is probably stereochemically hindered. This may explain the inability of NTMP to retard precipitation in the system studied and may also explain the slightly better performance of ENTMP compared with HEDP as demonstrated by the induction times and precipitation rates measured in each case, respectively. Moreover, since all compounds tested contain the same type of ionizable groups, it may be suggested that the ability of a compound to inhibit crystal growth depends not only on its ability to form surface complexes but also on the geometry of these complexes relative to the surface of the substrate. It is for this reason that emphasis should be given to molecular modeling studies in order to obtain a better picture of the energetics of the adsorption process on the surface of the crystals. Conclusions The crystal growth of iron sulfide was studied in stirred, stable, supersaturated solutions inoculated with synthetic
Langmuir, Vol. 14, No. 5, 1998 1255
pyrite seed crystals at constant pH ) 6.5, 25 °C. Pyrite was grown exclusively on the seed crystals introduced. The pyrite overgrowth was accompanied with a change in morphology from spherical to needlelike. The kinetics analysis suggested that the mechanism of crystal growth is governed by surface diffusion processes. Three organophosphorus compounds (NTMP, ENTMP, HEDP) were evaluated as inhibitors of pyrite crystal growth. NTMP failed to inhibit pyrite crystal growth. The other two compounds inhibited drastically pyrite growth (70-100% relative inhibition) in the concentration range between 10-6 and 10-10 M. HEDP modified the morphology of the pyrite crystals grown on the inoculating seeds favoring granular appearance. Acknowledgment. The authors wish to express their thanks to the European Commission for financial support, Contract No. JOU2-CT92-0108. LA970354C
