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LA-uR--86-133O DE86 010187

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D. R. Janeckyl, W. E. Seyfried, Jr., and M. E. Berndt Department of Geology and G-physics, University of Minneseta. Minneapolis, MN 55435 ‘ now at: INC-7. Los Alamos National Laboratory, Los Alamos, New Mexico 87545

INTRODUCTION

Oxidation-reduction reactions are important in many hydrothermal systems, predominantly in- volving the elements O, H, C, S, and Fe. but also significantly a~ecting a wide variety ~f minor and trace components. Fe-bearing silicate. oxide. and ,sulfide minerals are commonly used in hydrothermal experiments and petrographic studies to define the redox state of a system, We have experimentally investigated reactions of magnetite, hematite, pyrite. and pyrrhotite with 0.5 molal NaCl solutions. at temperatures of 30&425’C and pressures of 400-5@0 bars, using flex;~le cell hydrothermal equipment [Seyfried et a/.. 1986), These experiments allow reversible steady state solution compositions and the rates of equilibration of various redox couples (including HZ-H20, HaS-SOd. and CC)S-CH4) to be determined,

MAGNET1-fE-HEMATITE EXPERIFJIENTS

Magnetite-hematite-solution equilibria is characteristic of mineral assemblages used to control system redox state. or other potential variables, in a hydrothermal experiment to simplify the analyses necessary and to constrain the ex~eriment to geologically reasonable conditions. We have performed ~Yoeriments to evaluate equilibration of solution redox couples at 300’C, 500 bars. with a mind assemblage of magnetite and hematite, and a 3.5 wt.% NaCl solution (pH=-5.4, total dissolved CO, =-6 mM), To confirm reversible redox equilibrium, two experiments were run simultan~ously, one starting with excess 00 dissolved in solution and the other with excess Ha and CH4, Concentrations of Ha in solution for both experiments reached a consistent steady state concentration after ‘2000 hours of reaction (Figure la). 02 initially present in the oxidised experiment was rapidly consumed by the reactions to below analytical detectian limits. In contrast, concentrations of CFi4 did not reach a consistent steady state between experiments (Figure lb). S04 concentrations in solution decreased slightly, while sulfide was produced. in both experimmts (Figure lc), Dissolved Fe increased rapidly to ~ maximum in the reduced experiment, and then decreased to achieve a steady state concentration consistent with the oxidised experiment (Figure Id), The steady state H, concentration observed in the experiments is consistent with ●quilibrium concentrations of H: calculated for magnetite-hematite equilibria using SUPCRT (Helg,eson et d., 1978, 1981 and references therein and data for H, volubility (Himmelblau, 1959: Drumnwnd, 1982: Naumov et al,, 1971) (Figure la), \ he data of Kashima and Sakai (1984) for dissolv~d Ha i~ oagnettte- hematite. solution cxperimants at 300’IC is also consistent with our experimental data set, but does not represent steady state or ●quilibrium con:entlations (Figure la), SOt-H2S reacticms, (!vJluated by distribution of ~pecies in aqueous solution using SOLVEQ (Reed, 1982) and the aqueous sperne dissociation constant data set of Janecky (1982), Jlso indicate a r,fose approach to equilibrium with the magnetit~hematita assemblage. CHf concentrations in both experiments are, however, fir~atly In excess of that which would be in equilibrium with the C0,2 present, ~t the measured concentrations of H.], CH4 in tho oxidised experiment was apparently Benerated by breakdown of a small amount of organic carbon dissolved in the deionized w,lter used, These results are consistent with stlldies of nat~ral systems in which the redox couples HJ-H,10 and S04-H4S, and Fe-bearing minerals onulllhr,ltc, . while CH4-COj does not (Arnorsson and Cmmlaugsson, 1985: Jarwcky and Seyfri~d, 1984), -

SULFIDE-BEARING SYSTEMS

Sulfide is an important reactive component of rcdox rctictions in hydrothermal systcnls at wI. evated temperatures. For example, in high temperature oceanic ridge crest hydrothermal systems, sulfide alteration rnirmrals form within the basalts am.1sulfide deposits are produced whe) hydmttwr mal solutions exit into seawater, To vnderstilnd [he source and re;lction controls on high twnprvature hydrotlwrrnii! solutions, such o~ those nt 21’ N, E~st P,)cific Rise VOI1 Damm et d,, 1985), thr rela tionship~ betwctm solution composition, twnperature, pressure, anA krsalt alteration pathway Inust h under~tood, Two typas of experiments can ba used to examine these rclationship~ 1) SIIIIPIC mincrd raactlon exmwiments fe,~,. magnetita-llerntitite-py rite reactiorr with 3,5 wt, % NaCl solution), w,rf 2) Janecky, D. R. 2

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Figure 1. Concentrations of analyzed species in solut on versus IOK hours of reaction for hematite- magne~ite- solution experiments with initial excessesof 0~ (open circles), and H I and CH4 (closedcircles), Uncertainties for the analysm are shown by vertical lines, if larger than iyrnbol dh ter. Circleswith attached arrows r~present maximum or detection limit concentrations, (a) Dissolved H? data with respect to calculated ●quilibrium With hematile-magnetite and equivalent log f,,1. Open square is data of Kashima and Sakai, (1984). (b) Dissolved CH4 data. (c) Dissolved sulfate and sulfide data. (d) Dissolved Fe data.

basalt reaction with evolved seawa!er compositions (Seyfrled and Janecky, 1985), Magnetitt+pyrite and magnetite-hematite-pyrite experirwnts at 3501C and 4001’C approach re- versible equilibrium wiih solution. although slowly at 350’C (Figure 2), In contrast, at 400 C Fe- bearing assemblages of pyrite-pyrrhotite-magnetite appear tc equilibrate metastablv with pyrrbotlte and magnetite, within the pyrh fdd (Figure 2), At 31511C, however. the pyrite- pyrrhotite- m~gnetlte ass~mblage equilibrates with pyrite and magnetite (Figure 2), Mineralogical relationships observed In the charges at the completion of the experiments are consistent with these relationships, Basalt-solution ex eriments at 375 and 409”C. appeai tu equilibrate or closely a~p~o,~ch e.qull~- bration with F*oxide [and sulfide) assemblages (Figuro 2. A! 350’C, equilibration was .~ppatently slow and at the end of the experiment sulfide remained slig1lt!y below pyrite saturation (Figlire 2) At 4251C, sulfide cw~~entrations are abovo pyrrhotite and/or ma~petite stah;lity with pyrite, evvn though -- magnetite was present, Sinlilarly, at 400’9C and water-rock ~aticr of 0.5, the solution colnposltmn IS at metastabla ●quilibrium with rragrwtite-pyrl hotite and heaagona! oyrrhotite c[ystals wwe observed with several ragged margins and overgrowths of smaller pyriu? crystals (!3eyfried and Janecky, 1985), Water-rock ratio is also important in determining the stability of sulfide minerals, with those 400 C experiments at ratios greater than ‘2 Lwking observable pyrite and indicating undersaturatlon with pyrite (Figure 2),

CONCLUSIONS AND APPLICATION

The exper;mwtal data on solutmn compositions and mmerals for simple bufier assemblages and ● bti~i]lt-.. “solution interaction provide a very important check on rodox quilibria Ciagrarns [or the Fc-O- -1 .------:- .I-- -.;0:--1 ●-”:nm u*II k=uam~ *k- ~nnumntinnal limit- Janecky, D. R.

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Figuru 2, Stability fields for Fe oxide and sulfide minerals with respect to Hll,,,,l and hqS acr~vi:ies In e~pef. imental solutions at 400 bars (Solid Squares — EPR ferrobasalt-solution interaction: Solid Circles -- MAR olivine-normative basalt-solution interaction: Open Squares — magnetite-pyrite experiments at 500 bar? op~n Diamonds — magnetite-pyrite-hematite experiments at !;00 bars, Open Circles and Open Triangles — pvfll Q- pyrrhotite-magnetite experiments) and ●rid-member hydrothermal solutions at 21’ N EPR (X) (Craig ●t J/ l)CIO Welhan, 1980; Welhan and Craig. 1983) In the 400”C diagram, pyrite-magnetite-hematite relation~hlps .~Ic JI\O

shown for 500 bars pressure, the rnctastable pyrrhotitc magnetite boundary is shown, and basalt -~CIUIIm CQ. perlments with water-rock ratios of O S dnd 5 are l~heled Thermochemical data for mineral (JiSSr)lUtit>~~~A{ll~n\ from Helgeson et ●l. (1978, 1981, and referencestherein), for H? volubility from Hirnlnel’~l~u(1959) ~rurnrn,>nd (1982), and Naumov ●t d (1971) and for H2S(,,,,I dissociation from Ellis and ~~ge~,lba~h (1971]

for most solution equilibria computations (Figure 2). The experiments Jlso indicate that k ilrtlcs of -- Fe-oxide redox reactions are important, at least on nrr cxperirnental time scale (Fi~~le 1 ,]i~(! J,il,ecky and Seyfried, 1986). The results of the buffer assemblage and water-rock reaction experiments may be applied !(;, ~nder- standing the redox reactions in natural systems, For example, 211N, EPR, end-member hydro!llcrlnJl solutions are apparently undersaturated with all sulfide and Fe oxide minerals at vent conditions ~350 C and 260 bars) (Janecky and Seyfried, 1984 and Figure 2 . In col~trast, experimental I.)asdlt rr,lctlor] with evolved seawater at 350-4CQ C results in growth o) pyrite and magnetite, whil~ In soIne c,)st~

pyrrhotlte was observed reacting to form pyrite (Seyfried and Jar~ecky, 1985), The t~,pf!rirllt~llt,ll II; II I~. rJl assembla es are consl~tent wit+ green stone phas~ assemblages and theoretical predlctlorl~ ( FIY,U(:l )) 2111N, EP h . end-rnen~ber hydrothermal solutions do Jppcar to approach equilibrium with m,]grlt~[ltu JIId pyrite and agreement with experimental results at 40011C irnd ’400 bars, but are significantly sep~r;ltcd from this assemblage and the experiment~l results at the other temperatures examined (F, gurc 2) Jmnecky, D. R. 4

ACKNOWLEDGEMENTS: The experimental wcwk was supported by NSF grants OCE-8018644, OCE-8315116. OCE-8400676 and OCE-8542276. Preparation of this paper was supported by funding from USDOE Ofice of Basic Energy Science to Los Afamos National Laboratory.

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