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THE EFFECT OF SILICA, pH AND TEMPERATURE ON PYRITE PRECIPITATION

R. N. Cunha 1, R. M. Z. Castro 1, M. M. Seckler2

' S. D. F. Rocha 3

1 Universidade Federal de Minas Gerais, Departamento de Engenharia Química Rua Espírito Santo, 35- Centro, Belo Horizonte, Minas Gerais, Brasil

2 IPT- Instituto de Pesquisas Tecnológicas, Prédio 36, Cidade Universitária "Armando de Salles Oliveira", 05508-90 I, São Paulo, SP, Brasil

3 Universidade Federal de Minas Gerais, Departamento de Engenharia de Minas Rua Espírito Santo, 35 -Centro, Belo Horizonte, Minas Gerais, Brasil

sdrocha@demin. ufmg. br

RESUMO

O efeito da sílica na precipitação da pirita foi investigado visando contribuir na interpretação da gênese das foi-mações feníferas bandadas. Ensaios de precipitação em sistema batelada foram realizados na ausência e na presença de sílica. Pirita fo i sintetizada pela reação entre íons ferrosos e íons sul fetos em solução aquosa. i\ v aliou-se a influência das variáveis tempo (O a 550h), pH (6 e 7), temperatura (65 e 80°C) e concentração de sílica presente no sistema (O a I OOOmg/L) na fonnação do precipitado. Os produtos sólidos foram caracterizados por difração de raios-X, microscopia eletrônica de varredura e determinação do tamanho das partículas. Em pH 6 e 65°C, obteve-se um favorecimento da precipitação da pirita, tanto no sistema puro quanto em presença de sílica. Em pH 7 e 80°C, a presença de 80mg/L de sílica também favoreceu a precipitação da pirita em relação ao sistema puro. No entanto, nessa temperatura e em pH 6 um efeito contrário foi observado. Baseado nas concentrações de enxofre solúvel, observou-se que, em pH 7 e 80°C, elevadas concentrações de sílica (500 e IOOOmg/L) podem mascarar o efeito da silica. Para toda a faixa de concentração de sílica avaliada, não foi observada alteração na morfologia das partículas de pirita formadas.

Palavras-chaves: precipitação, pirita, sílica, formações ferríferas bandadas.

ABSTRACT

The effect of silica on pyrite precipitation was studied in order to contribute to lhe investigation to the genesis ofbanded iron formations (BIF) . Batch precipitation experiments were canied out in the presence and absence of silica. Pyrite was obtained by lhe reaction between ferrous and sultíde ions in aqueous solution. The influence of time (O to 500h), temperature (65 e 80°C) and silica concentration (O a I OOOmg/L) on the amount of solid formed and also on the solids morphology was investigated. The solid reaction product was characterized with respect to lhe presence of crystalline phases and size. ln pH 6 and temperature of 65°C, pyrite formation was favored both in pure system and in lhe presence of silica. The presence of 80mg/L of silica in pH 7 and temperature of 80°C also favored the pyrite precipitation in comparison of pure system. However, in Iower pH, 6, lhe opposite effect was observed. The sulfur soluble in solution thus indicates that in pH 7 and temperature of 80°C, high silica concentrations (500 e lOOOmg/L) could mask lhe silica effect. For lhe whole concen!ration range evaluated, it was not observed change on lhe morphology of pyrite particles formcd .

Keywords: precipitation, pyrite, silica, banded iron formation

190

Renata Nepomuceno Cunha, Roberto Machado Zica de Castro, Marcelo Martins Seckler and Sônia Deni se Ferreira Rocha

l. INTRODUCTION

Pyrile (FeS2) is an imporlanl source of sulfur and iron in lhe earth 's crus!. Pyrite deposits are oflen explored due to the gold or copper associated with it, FeS2 itself being a source of sulfur for sulfuric acid and ferrous sulfate production. Synthelic finely divided pyrite finds application in lithium batleries as lhe active cathodic material. II has also been considered for use as anodic depolarizer in the production of electrolytic hydrogen. Pyrite deposits occur as alternated layers of silica and iron minerais such as magnetite, hematite, siderite and pyrite, and conslitute the so-called banded iron formations (BlF). This alternation of layers suggests that BIF would be associated to some cyclic natural phenomenon such as the earth's rotation (Castro, 1994) or translation (Trendall and Blockley, 1970). These formalions derive from the pre-cambrian period, when both hydrosphere and atmosphere were of a reducing character. ln order to betler elucidate the genesis of BIF, the formation of pyrite in reducing environments has been extensively studied (Brener, 1970, 1983 and Braterman et ai, 1983).

For low pressures, temperatures below I 00°C and under a reducing environment, hydrogen sul fi de and dissolved iron react to form amorphous ferrous monosulfide (mackinawite) with chemical formulae varying between FeS0•89 to FeSu. This precursor, in the presence of elemental sulfur or polysulfide ions, transforms into pyrite (Osseo-Asare and Wei, 1996, 1997). For non reducing conditions, i.e. with little oxygen present, amorphous FeS lransfonns into hexagonal pirrotite (similar to mackinawite in structure and composition), greigite (Fe3S04) and finally pyrile. The complete oxidation of amorphous FeS results in the formation of the ferric oxide called lepdocrocite (y-FeOOH). The conversion of mackinawite to greigite is likely to be promoted in slightly reducing conditions. Apparently, the Ostwald rule of stages applies to the precipitation of iron sulfides, i.e. the first compound formed is thennodynamically the least stable (Morse et ai., 1988).

The main chemical species involved in an aqueous systems containing Fe and S are summarized in Table I. Reaction mechanisms involving theses species have been proposed for the formation of pyrite. Robert et ai. ( 1969) suggested that pyrite formation under sedimentary conditions occurred by a direct reaction between ferrous and disulfide ions, eq. (I). However, the concentration of disulfide ions in natural waters is very low due to the fonnation of polisulfides S2

-,

(Murowchick, 1986), which would preferentially be the sulfur source for pyrite in sedimentary formations (Rickard, 1975; Goldhaber and Kaplan, 1974), eq. (2). More recently, Luther (1991) suggested that not only the specie Fe(HSt as a1so solid iron monosulfide directly reacts with the polisulfides ions to nucleate pyrite.

ble I. Chemica1 "libria in the Fe-S-H,O ~ -------

H+ + Hs- ~ HzS Fe2+

. Hs- _____:,.

Fe(HSt + ~

Fe3+ + Hs- _____:,. 2Fe2+ + so + H+ ~

so + Hs- _____:,. st + H+ ~

so + 2St _____:,. st ~

Hz _____:,. 2H+ ~

z+ s2- _..::..F s Fe <•ql + 2(aql~ e z<s> Fe2+(aq) + S2

X (aq) + HS-(aq) '""'FeS2(s) +S2-x-l (aq) + H\aq)

(1) (2)

The diagram Eh-pH for the system Fe-S-H20 at 25°C is presented in Figure 1. It can be observed that pyrite is stab1e under a board range of pH and in reducing conditions. ln high v alues of Eh the oxides/hydroxides predominate.

ln this work pyrite precipitation in the presence of si1ica, a very important component of BIF, has been investigated.

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XXII ENTMME I VII MSHMT- Ouro Preto-MG, novembro 2007.

Eh (Volts) 1.8

1.6 Fc' 3 (aq) Fc(OH)'2 <,~ 1

1.4

1.2 Fe(OH)3 Fe02. 11aol

1.0

0.8

0.6 Fc +2(>~1 0.4

0.2

0.0

-0.2

-0.4 Fe +21ao\

-0.6

-0.8

-1.0 Fe

-1.2 o 2 4 6 8 10 12 14

pH

Figure 1- Eh-pH for the Fe-S-H 20 system at 25"C and concentrations of 0.00 I M of iron and 0.00 I M of S.

2. EXPERIMENTAL

A ferrous sultide (FeS) suspension was prepared by first rapidiy adding 20 mL of a 0,39 M Na2S.5H20 solution into a beaker containing 20 mL ofa 0,1 M FeCI2 soiution. Thereafler I mL ofa 0,1 M NaOH solution and 3mL water were added and the flask was sealed with piastic film. FeS precipitated instantaneously. Ali solutions had previously been deionized and deaerated by bubbling with nitrogen gas for half an hour. Pyrite precipitation was accomplished by mixing 0,084 g of elemental sulfur, 3 mL of sílica solution, I O mL of a FeS suspension and by filling the flask to 50 mL with KH2P04 - NaOH buffer solutions wilh pH values of 6, 6,5 or 7. The reaction was allowed to proceed for 240 h (experiments with pH 6) and for 380 h (pH 7). The silica solutions were prepared with commercial silica in order to obtain lhe following initial conccnlralions in the reac lion media: O, 80, 100, 200, 500 and I 000 mg Si02/L. The fl asks were sealed with plastic film and placed in an oven at 60 or 80 °C. The flask contents were manually agilated once a day. The suspension was hot tiltered through Millipore 0,22 J..llll membranes. ln the filtrate, the total sulfur concenlration in the Iiquor was dclennined by a turbidimetric method involving sulfur oxidation to sulfate with water peroxide and subsequent sultàte precipitation with barium chloride solution. The precipitate was washed with water, carbon disullide (CS2) and acetone and weighed. CS2 was used to remove any non-reacted elemental sulfur. The silica content in lhe precipitale was determined. The solid product was characlerized by X-ray diffraction and scanning electron microscopy (SEM) coupled with energy dispersive spectrometry (EDS). Its size distribution was also determined with a Coulter-Counter. The amount of unreacted FeS in lhe product was determined by washing the precipitate wilh a hot 0,6 N HCI solution. Since FeS dissolves, weighing lhe dry precipitate before and after washing gives the amounl of FeS which was originally present in the solid.

3. RESULTS

3.1. System without silica

Precipitation of pyrite at 80"C and pH 6,5 did not proceed substantially after 144 h, since the amount of precipitate formed was constant thereafter (Figure 2) . During the reaction, a yellowish color developed, which is characteristic of polisultides. Soluble sulfur concentration initially decreased and then increased (Figure 2), indicating that the precipitation of pyrite dominated the process in its early stages and that !ater the formation of polisulfides was dominant. The amount of precipitate fonned did not increased, despite it was soluble sulfur present into solution. The

192

Renata Nepomuceno Cunha, Roberto Machado Zica de Castro, Marcelo Martins Seckler and Sônia Denise Ferreira Rocha

pyrite formed may be deposited on the FeS particles, thus hindering the progress ofthe reaction. lncreasing lhe pH from 6 to 7 caused a reduction in lhe amounl of precipilate fonned (Figure 3) and a corresponding increase in lhe amount of soluble sulfur (Figure 4 ). These resulls are consistent with data obtained by Wilkin el ai. ( 1996b ). These authors used polisulfides as pyrile precursor and found for pH 7 a conversion into pyrile of 75 wt% of lhe iron initially present after 218 h. For pH 8, a conversion of only I O wt% was obtained by lhose authors . Robert et ai ( 1969), as well as Rickard (1975) also found that the precipitation is fasler for lower pH values. Figure 3 also shows lhal a lower temperature favors a larger amount of precipita te.

0,181

_. I 800

~

700 '§,

Ol E à) c :§ o "õ.. 0,16 ~ "ü 600 -E QJ .....

-\ QJ

c.. o - c o o c o

..... :::J 5oo E o E 0,14 .9-Cll :::J

Vl QJ

400 :g o Vl

0,121

IÍ .......,

I 300

70,5 143,5 240 310

time (h)

Figure2- The effect oftime on the amount ofprecipitate and on the dissolved sulfur concentration (80°C e pH 6,5).

X-ray diffraction showed that pyrite was the main phase formed in ali experiments. Pyrite particles investigated by SEM had framboidal morphology (Figures 5 and 6). According to Wilkin and Sarnes ( 1996a), the framboidal formation results from 4 consecutive processes: (i) nucleation and growth of ferrous monosulfide particles; (ii) their transformation into greigite (Fe3S4}; (iii) the aggregation of uniformly sized greigite microcrystals to give a framboidal morphology; (iv) transformation ofgreigite into pyrite. Wilkin and Sarnes (1996c) also propose that framboidal habit is associated to slightly oxidizing conditions and to high supersaturations. Particle sizes were within the range of 2 to 20 J..lffi (Figure 7).

3.2. System with silica

ln ali experiments, pyrite was the only phase detected with X-ray diffraction. SEM with EDS revealed that sílica was associated with iron and sulfur and that no individual crystalline sílica particles were present. The co-precipitated sílica was amorphous . The amount of precipitate formed increased only slightly with the sílica concentration in the liquid phase for values up to 500 ppm When 1000 ppm sílica was present , however, the amount ofprecipitate formed was four times higher (Figure 8). Correspondingly, the concentration of soluble sulfur decreased slightly for liquid phase sílica concentrations up to 500 ppm and showed a marked drop for I 000 ppm.

Chemical analysis revealed that sílica incorporation in lhe precipitate increased with sílica concentration in the liquid phase (Table II) . The solubility of amorphous sílica is, approximately, 350 p.p.m at 65"C, 420 ppm al and is practically independent of pH in the range of interest here (ller, 1957; Saumann, 1957). Therefore, in the experiments with 500 ppm and 1000 ppm of sílica was co-precipitation possible. The increase in Si content in the precipitate from 4.2wt% to

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XXII ENTMME I VIl MSHMT - Ouro Preto-MG, novembro 2007.

24 wt% and the incrcased mass of precipitate formed in these conditions indicate that co-precipitation indeed took place. For ali other concentration values, Si was mos! likely adsorbed on lhe surface of pyrite parti eles.

Pyrite particle sizes increased substantially with Si content in lhe reacting. The mean particle size D50 for the pure system and for 500 ppm Si were respectively 3,8 and 7,0 !liTI. For 1000 ppm Si, D50 was 9,8 !liTI, but this measurement represented not only pyrite, but also included co-precipitated silica particles (Figure 7). It has been reported (Schenck, 1967) that monomeric silica promotes lhe oxidation of ferrous ions to ferric ions in oxidizing media for the pH range of 6,5 to 7. Ferric ions react promptly lo form pyrite (Osseo-Asare and Wei , 1996) according to the reaction:

2 3+ 2Sz- ~ 2+ S Fe (aq) + n(aq) ~ Fe (aq) + Fe 2(s) (3)

Although ferric ions would eventually lead to the formation of oxidized phases such as FeOOH, in environments with reducing conditions, where sulfides are the predominant species, these iron oxihydroxides would be consumed to form ferrous ions again (Wilkin and Barnes, 1996c):

(4)

0,16

0,14 - 65°C - I'Z:.J2! ao0 c ~ Q) 0,12 ..... cu ..... 'õ.. 0,10 'ü Q) ..... 0,08 Q_ .._ o ..... c

0,06

:J o 0,04 E <{

0,02

0.00

5 6 7 8

pH

Figure 3- The effect of pH on lhe amount of precipitate formed.

194

Renata Nepomuceno Cunha, Roberto Machado Zica de Castro, Marcelo Martins SeckJer and Sônia Denise Ferreira Rocha

.-- 1800 '§, I• 65oc j ~ 1500 • o ~80°-º.J ~ 1200 c: Q,) o 900 c: 8 ....

ª 600 :::J

300 1

• VI Q,)

:c • :::J • o (/) o

5.5 6 6.5 7 7.5

pH

Figure 4- The effect of temperature and pi I on sulfur concentration (pH 6, 240 h and pH 7 ,384 h) .

. ...... ,~.,.,.~\ ~~ ·- -~· q I •····• ~-;, ·• ~ /' fT# . • • ' '.'J:..l . • .. : ... ' ,· ····'!':· ,. ... '

,. ~·'li ·' . ·. I "'"li . 1 . j,. '

~~r;.·.· ...••.•• --. ~. .. . ; '~ ~- . . :. . .':til-···· #~ { ,. ~~ . ( ~.' . . ' f ' ~· ; . ti' '. -~ . o F,-~. ~ . , r . ...--, "" •

~ '(~ ... .... ' "' ; ~·· .\' C'. ~-·•: t~' ·- ~.,t.'

- - -• ..-._ ~~' e e .:'I t~~

Figure 5-Precipitate formed at pure system (pH 6,0) and temperature of(a) 65°C e (b) 80°C (reaction time of240 h)

r # r ....... , ... . • . ;: . , .,_.,, . . • •.. , . 1 • ,.·.,) . ~~ ..... "' . I' • •\_ " : . ~ .. ~~- . (• . . . .. ..... .

• \ ... •' ."& ....... .. • ,._.• > ~- l \r..., ,., ,• O ,.,., ,. . _. ~-· )' ..... ... ,) ......... I ..;N . -~ \ ·P-~':?.~ . .. • . .. ·

. •• ••. ,. - • 'JrJ· • ~ •.• • '!) ~J •• ?, -)7' ··' ·, ··ci!.'.1 ,. _.., :, ·1), '1J ; ' 4 (ii.C· l .... t ... . ' . • • . .·· .J • ' •. ••• ,. ,, : • . '

. · ·~ ' ,.1, IJ ti;..! •1·• •• • • .J ""' I-~ • 1' 4!.a>' ·> • ; • ·\ • • P- J.... . ') ' ) ~ ~ . ,,.

h .. ~· \'t ·,,.~.,;-'t•• J L.· \

. ...... ,~ :-.. t:J', • . r ~ ' ..,"' ·~ .Y'f3 :> • l~J,. l . ..~.:-~: f • . ~~-~ ,,:ocz. 99" ~ .. J ... r::~ • .• ~'J . v \. ~ ' • >; ' ~ . .. . et3~· a. . •

~ ~ . ···'"f • ... . ·--.·' . r t'-. ·'' -~·, ' ~ •• ., r , "' . . .,.ti.. ' ..

. \ ,.. r. ') ~ ;' .)) """'-. • .~ ·,. .. . . . " . ........ ~ , .. ~

~ . . ' ·; .. ,. ...if,. 1 ..J', -y .. f

•, ,- ·-..J <# ) ~ I I

'-" ),, ~ > . ' t .. 'lo. - ) •. \, ' \\,../ ' .

'- .-~· j. • .. u~ . ,' '-..ti- ~ ,. ... • / ~ i 1 :y; .. } J,• ...) r"l . :J , 11 ~ -~

~·:l} u ,~,·4P~~ -.~: .~ ~Olf66.) .~:. .,- - ..... \..A·'---

Figure 6- Precipita te formed at pllre system at pH 7 ,O and temperature of 65°C (Reaction time of 384 h)

..

XXII EN TMM E I VII MSHMT - Ouro Preto-MO, novembro 2007.

120

Q) 100

N 'iii Qj 80 "O c :::J

~ 60 o Q)

-~ .§

40 :::J E :::J

ü 20

o

I I i '

I I I I I i I

r--- · .. ·- . !

I ff.IT 1----::-:-- - -;., -. -. --~//. 1 v '

I I I ' . I 1---- - pH 6.wc i 1 : ' ' ' ! -" •, I _,.:

I ./ , ' I / · I ! --pH 7 , 65"C I I I

I I • • "" I I

/ I -- - pH 6.80"C , j

I _,r; i I . v / ~ I I

i - - - • pH 7 . sooc I . - - · , --- -

,• I i ,_.· / / !

I r

/

I I

I .·)// i I ' I I

10 100

Figure 7- The effect of pH and lemperature on lhe precipita te size distribution in absence of silica.

1

-- 0.8 --0) ,.._.... Q) -ro

:g_ 0.6 -'ü Q) '-o..

...... ~ 0.4 c ::::1 o E ro 0.2

o o

•

Particle Diameter (~m)

... • ...

•

300 600 900 1200

silica concentration (mg/1)

Figure 7- The effect of silica concentration on the amount ofprecipitate formed at pH 6.

Table II- Silica content ofthe precipitate formed at ~H 6,5 and 80°C, after 240 h ofreaction.

Dissolved silica concentration (mg/L)

80 100 500 lO OO

Silica content in the precipitate (mass%)

<O. I 0.4 4.2 24.5

195

196

Renata Nepomuceno Cunha, Roberto Machado Zica de Castro, Marcelo Martins Seckler and Sônia Denise Ferreira Rocha

CONCLUSIONS

Pyrite synthesis from ferrous ions, elementary sulfur and sulfide ions is feasible and may be easily implemented. The extent of conversion of pyrite from ferrous monosulfate is higher at low pH values in the range 6 to 7 as well as at low temperatures in the range of 65 °C to 80°C, both in lhe absence and in lhe prcsence of sílica in lhe reacting media . Precipitation is complete after 144 h for a pH of 6,5 and 80°C. Soluble sílica present during pyrite precipitation incorporates in the product. For liquid phase Si concentrations in the range 80 to 500 ppm incorporation is likely to follow a surface adsorption mechanism. The Si content in the product increases with liquid phase concentration from <0, I to 4,2 wt%. For I 000 ppm Si in the liquid phase, substantial co-precipitation of amorphous sílica took place. The conversion of ferrous sulfide to pyrite was substantially enhanced in lhe presence of sílica ai I 000 ppm Pyrite parti eles synthesized in the presence of sílica were larger than in the absence of this admixture, but their morphology was not affected.

REFERENCES

Saumann, Seitr. Silikose, Forsch, v 37, 4 7, 1955 apud II e r. R. R. The Colloid Chemistry of Sílica. ln: Colloidal-Silica­Concentrated Sois; The occurrence, Dissolution and Deposition of Sílica; Water-Soluble Silicates, Cornell University Press, New York, 1957. Serner, R.A. Sedimentary Pyrite Formation. American Journal of Science, v 268, p 1-23, jan, 1970 Serner, R.A. Sedimentary Pyrite Formation -an update. Geochimica et Cosmochimica Acta, v 48, p 605-615, dez, 1983. Bratern1an, P.S.,Cairms-Smith,A .G. and Sloper, R., "Photoxidation of Hidrate Fe2

+ Significance for Sanded lron Formation", Nature, v 303, pl63-164, 1983. Castro, L.O. Genesis ofSanded Iron Formations. Economic Geology, v 89,n 6, p 1384-1397,out, 1994. Goldhaber, M.B. & Kaplan, I. R. The sulfur cicie in the sea. vol 5, cap 17, Wiley, 1974. ller. R. R. The Colloid Chemistry of Sílica. ln: Colloidal-Silica-Concentrated Sois; The occurrence, Dissolution and Deposition ofSilica; Water-Soluble Silicates, Cornell University Press, New York ,1957. Luther, G.W. Pyrite Synthesis via Polysulfide Compounds. Geochimica et Cosmochimica Acta, v 55, p 2839-2849, junh, 1991. Morse, J.W.; Cornwell ,J.C.; Rickard, D. The chemistry of hidrogen sulfide and iron sulfide system in natural waters, Earth Sei, 24, pl-42, 1988 apud Luther, G.W. Pyrite Synthesis via Polysulfide Compounds. Geochimica et Cosmochimica Acta, v 55, p 2839-2849,junh, 1991. Murowchick, I.S .& Sarnes,H.L. Formations of cubicc FeS, American Mineral,v7l,l986 apud Krauskopf, K. S . Introdução à Geoquímica~ São Paulo,Polígono,v I, 1972 , 294 p. Osseo-Asare, K. & Dawei Wei ,D. Aqueous synthesis of finely divided pyrite particles. Colloids and Surfaces A: Physicochemical and Engineering aspects, v 121, p 27-36, 1997. Osseo-Asare, K. & Dawei Wei,D. Particulate pyrite formation by the Fe3

+1HS. reactions in aqueous solutions effects of

solution compositions. Colloids and Surfaces A: Physicochemical and Engineering aspects , v 118 , p 51-61,1996 Rickard D.T., Kinetics and mechanism ofpyrite formation at low temperatures. American Jou'tnal of Science, v 275, p 636-652, junh, 1975. Robert, W.M.; Walker, A .L.; Suchanan, A.S. The chemistry ofpyrite formation in aqueous solutions and its relation to the depositional environments. Mineral Deposita, v 4, 1969. Schenk, J.E. & Walter W.J.J. Chemicallnteractions ofDissolved Sílica with lron (II) and (III). Jour. A .W.W.A, 1967. Trendal, A.F. & Slockley, J.G. The lron Formation of Precambriam Hamersley Group, Western Australia Geological Survey Bulletin, 119, 1970, 366p. Wilkin ,R.T. & Sarnes, H .L. Pyrite formation by reactions of iron monosulfides with dissolved inorganic and organic sulfur species. Geochimica et Cosmochimica Acta, v 60, p 4167-4179, 1996c. Wilkin, R. T.; Sarnes, H.L.; Srantley, S.L. The size distribution offramboidal pyrite in modem sediments: An indicator of redox conditions. Geochimica et Cosmochimica Acta, v 60, n 20, p 3897 - 3912, 1996b. Wilkin, R.T. & Sarnes, H.L. Formation Processes of framboidal pyrite . Geochimica et Cosmochimica Acta, v 61, n 2, p 323-339, 1996a. Willey, Mar. Chem, v2, 239, 1974 apud ller. R. R. The Colloid Chemistry of Sílica. ln: Colloidai-Silica-Concentrated Sois; The occurrence, Dissolution and Deposition of Sílica; Water-Soluble Silicates, Cornell University Press, New York ,1957.