Iowa State University Capstones, Theses and Retrospective Theses and Dissertations Dissertations

1989 Solubility of molybdenite and the transport of molybdenum in hydrothermal solutions Xiaoyun Cao Iowa State University

Follow this and additional works at: https://lib.dr.iastate.edu/rtd Part of the Geochemistry Commons

Recommended Citation Cao, Xiaoyun, "Solubility of molybdenite and the transport of molybdenum in hydrothermal solutions " (1989). Retrospective Theses and Dissertations. 9026. https://lib.dr.iastate.edu/rtd/9026

This Dissertation is brought to you for free and open access by the Iowa State University Capstones, Theses and Dissertations at Iowa State University Digital Repository. It has been accepted for inclusion in Retrospective Theses and Dissertations by an authorized administrator of Iowa State University Digital Repository. For more information, please contact [email protected]. INFORMATION TO USERS The most advanced technology has been used to photo­ graph and reproduce this manuscript from the microfilm master. UMI films the text directly from the original or copy submitted. Thus, some thesis and dissertation copies are in typewriter face, while others may be from any type of computer printer. The quality of this reproduction is dependent upon the quality of the copy submitted. Broken or indistinct print, colored or poor quality illustrations and photographs, print bleedthrough, substandard margins, and improper alignment can adversely affect reproduction. In the unlikely event that the author did not send UMI a complete manuscript and there are missing pages, these will be noted. Also, if unauthorized copyright material had to be removed, a note will indicate the deletion. Oversize materials (e.g., maps, drawings, charts) are re­ produced by sectioning the original, beginning at the upper left-hand corner and continuing from left to right in equal sections with small overlaps. Each original is also photographed in one exposure and is included in reduced form at the back of the book. These are also available as one exposure on a standard 35mm slide or as a 17" x 23" black and white photographic print for an additional charge. Photographs included in the original manuscript have been reproduced xerographically in this copy. Higher quality 6" x 9" black and white photographic prints are available for any photographs or illustrations appearing in this copy for an additional charge. Contact UMI directly to order.

University Microfiims International A Bell & Howell Information Company 300 North Zeeb Road, Ann Arbor, Ml 48106-1346 USA 313/761-4700 800/521-0600

Order Number 0008505

Solubility of molybdenite and the transport of molybdenum in hydrothermal solutions

Cao, Xiaoyun, Ph.D.

Iowa State University, 1989

UMI 300N.ZeebRd. Ann Arbor, MI 48106

Solubility of molybdenite and the transport of molybdenum

in hydrothermal solutions

by

Xiaoyun Cao

A Dissertation Submitted to the

Graduate Faculty in Partial Fulfillment of the

Requirements for the Degree of

DOCTOR OF PHILOSOPHY

Department: Geological and Atmospheric Sciences Major: Geology

Approved:

Signature was redacted for privacy.

In Charge of Major Work

Signature was redacted for privacy.

br the Major Departmépt

Signature was redacted for privacy. For the Graduate College

Iowa State University Ames, Iowa

1989 ii

TABLE OF CONTENTS

Page

GENERAL INTRODUCTION 1

CHAPTER I. DISSOLUTION OF MoOg AND COMPLEXING OF 8 MOLYBDENUM IN NaCl SOLUTIONS AT ELEVATED TEMPERATURES AND PRESSURE

ABSTRACT 9

INTRODUCTION 10

EXPERIMENTAL METHODS 15

EXPERIMENTAL RESULTS AND CONSTRAINTS ON SOLUTION MODELLING 19

Deductions Based on Solution pH 25

Deductions Based on Oxidation State 27

ESTIMATION OF MOLYBDENUM SPECIES 29

GEOLOGICAL APPLICATIONS 42

CONCLUSIONS 53

CHAPTER II. SOLUBILITY OF MOLYBDENITE (MoS,) 55 IN HYDROTHERMAL SOLUTIONS

ABSTRACT 56

INTRODUCTION 57

EXPERIMENTAL PROCEDURES 61

EXPERIMENTAL RESULTS 63

GEOLOGICAL APPLICATIONS 71

Transport of Molybdenum 71

Ore Formation 86

CONCLUSIONS 91 iii

GENERAL SUMMARY 93

REFERENCES 95

ACKNOWLEDGMENTS 103 iv

LIST OF FIGURES Page

CHAPTER I.

Figure 1. Phase relations in the system Mo-Fe-O-S at 700 K. 17 M0O2 is the stable phase with respect to M0O3 under conditions where Eg2 is buffered by the assemblage 2^20% - FegO^

Figure 2. Plot of log[Mo(tot)] vs. log[NaCl(tot)] 21 illustrating dependence of MoGo solubility on chloride concentration at 300 °C with fQ2 controlled by MH. Solid line represents a least- squares fit to the data

Figure 3. Solubility of M0O2 in NaCl solutions at 350 °C and 22 fQ2 controlled by MH. Solid line represents a least-squares fit to the data

Figure 4. Plot of log[Mo(tot)] vs. log[NaCl(tot)] 23 demonstrating dependence of M0O2 solubility on chloride concentration at 400 °C and fQ2 controlled by MH. Solid line represents a least-squares fit to the data. Note the increase in slope of the solid line with respect to temperature in the range from 300 to 400 °C

Figure 5. Solubility of M0O2 as a function of NaCl 24 concentration in NaCl-H20 solutions at 450 °C with fQ2 buffered by MH. Solid line represents a least- squares fit to the data

Figure 6a. A representative plot of log u vs. log( [Na"'']/[H'^] ), 35 where u-y[MoT]+[Cl-]+[HCl]+2[FeT]-[Na+]-[NaOH], demonstrating the linearity of equation (17). When y-1, the value of x is given by the slope of the solid line which has a value of 1.9. Thus, an integer of 2 is selected for x

Figure 6b. A representative plot of log v vs. log( [H"*"] [Cl"]), 36 where v-x[MoT]+[Na+]+[NaOH]-[CI*]-[HCl]-2[Fe], demonstrating the linearity of equation (18). When x-2, a value of 1 can be determined for y from the slope of the solid line

Figure 7. Plots of y vs. x derived from equations (18) and 38 (17) respectively at 300, 350, 400, and 450 °C. Solutions of the system of equations (17) and (18) at each temperature can be found at the intercept of the two curves y=f(x) and x=f(y), which is shown V

in parentheses

Figure 8. Solubility of MoOg, normalized to fQ2> in HCl-H^O 43 solutions at 450 °C and 0.5 Kb. The shape of the solubility curve suggests a stepwise complexing of chloro molybdenum species. An oxidation state of 3+ is assumed for aqueous Mo species (see text for discussions). The value of n designates the ligation number, and a maximum value of 3 can be derived from the solubility curve. Constructed from the solubility data of Kudrin (1986)

Figure 9. Relative stability of aqueous molybdenum species 47 as a function of log fo2 and pH at 300, 350, 400, and 450 °C. Only those species which are determined from experiments are considered here. At 350 °C and above, solid lines represent boundary between an ion pair and other molybdenum species, and dashed lines between HMoO^' and other molybdenum species. P-0.5 Kb, a^j^^ - •= O.IM

Figure 10. Relative stability field of aqueous molybdenum 48 species as a function of log fQ2 and pH extrapolated to 600 °C. All other conditions and notations are the same as in Figure 9

Figure 11. Activity-activity diagrams depicting the relative 49 stability among aqueous Mo(6+) species as a function of log[Na+], pH, and temperature. Solid lines represent boundary between an ion pair (NaHMoO^® or Na^MoO^^) and bimolybdate, and dashed lines between an ion pair and molybdic acid

Figure 12. Phase relations in the system NaCl-H20 at high 51 temperatures and 1.0 Kb. There exists a large unmixing region above 630 °C, within which two phases are coexisting, one rich in H2O, and the other rich in NaCl. Adopted from Bodnar and Sterner (1987)

CHAPTER II.

Figure 1. Phase relations in the system Fe-O-S, illustrating 59 the buffering capacity of the assemblages Fe^O^- Fe203-FeS2 (MHP) and Fe30^-FeS2-FeS (MPP). MHP buffer has higher values of fQ2 and fg2 than MPP buffer does

Figure 2. Plot of logMo(tot) vs. logNaCl(tot) demonstrating 65 dependence of molybdenite solubility on NaCl concentration and temperature in MHP-buffered vi

solutions

Figure 3. Solubility of molybdenite as a function of NaCl 66 concentration and temperature in MPP-buffered solutions. In both Figures 2 and 3, the solubility increases with decreasing temperature under the experimental conditions

Figure 4. Plot of logMo vs. total sulfur concentration in 67 some run solutions buffered either by MHP or by MPP. There is a negative correlation between Mo and S concentration in the solutions

Figure 5. Comparison of molybdenite solubility in sulfur- 69 free and sulfur-bearing systems. HM denotes the buffer assemblage hematite-magnetite. Solid lines are solubility curves at 400 °C, and dashed lines at 350 °C. At constant temperature, the solubility in three different systems decreases in the following order: HM > MHP > MPP

Figure 6. Calculated molybdenite solubility as a function of 77 temperature with fQ2 controlled by HM at 0.5 Kb. ^H2S " O'Ol M, 3^2+ = ^ci- " pH =• 5.5. See text for calculation model

Figure 7. Calculated solubility of molybdenite as a function 78 of temperature with fQ2 and fg2 controlled simultaneously either by MHP or by MPP buffer at 1.0 Kb. pH - 5.5, and 8^^+ - agi. - 0.1 M. Note the retrograde behavior of solubility in the lower temperature range

Figure 8. Calculated solubility of molybdenite as a function 81 of pH in MHP-buffered solutions at 300, 350, and 400 °C, and 1.0 Kb. - ^ci. - 0.1 M. The measured solubility at relevant temperatures is also plotted for comparison. Triangle represents the solubility data at 350 °C and circle at 400 °C. Arrow indicates the direction that the measured solubility approaches the calculated value. Neutral pH for the run solutions is assumed

Figure 9. Calculated isothermal solubility of molybdenite as 82 a function of pH in MPP-buffered solutions at 1.0 Kb, a^a^ - ag^. =0.1 M. The measured solubility at relevant temperature is also shown for comparison. The circle denotes the solubility data at 400 °C, triangle at 350 °C, and square at 300 °C which is from Wood et al. (1987). Neutral vii

pH for run solutions is assumed except at 300 °C

Figure 10. Calculated molybdenite solubility as a function of 83 pH under unbuffered conditions in terms of fg2 at 300, 350, and 400 °C, 1.0 Kb. Oxygen fugacity is controlled by HM, - 0.1 M, and ^H2S " O'Ol Note the monotonie increase in molybdenite solubility with respect to temperature. The shaded areas represent the limit of error on the solubility calculation

Figure 11. LogfQ2-pH diagram illustrating dependence of 85 molybdenite solubility on fo2 and pH in NaCl solution at 350 °C and vapor saturated pressure. Total sulfur concentration is chosen to be 10"^ ®Na+"^Cl- "O'l Contours show the molybdenum concentration in solutions. Dashed line separates Mo(5+)- and Mo(6+)-predominant field

Figure 12. Fugacities of major gas species in the system 87 S-O-H at 1 Kb. fQ2 and fg2 are controlled by MHP buffer. At temperatures below 500 °C, reduced S species are more abundant than oxidized sulfur species viii

LIST OF TABLES Page

CHAPTER I.

Table 1. Experimental solubility data for M0O2 in NaCl 20 solutions buffered by magnetite-hematite. Concentrations in molarity

Table 2. LogK for some selected reactions used in the 33 calculation of molybdenum speciation

Table 3. Calculated molybdenum speciation scheme 39

Table 4. Equilibrium constants of dissolution of M0O2 (logK) 45

CHAPTER II.

Table 1. Experimental solubility data for molybdenite (M0S2) 64 in NaCl solutions. Concentrations are in molarity unless specified

Table 2. Equilibrium constants (logK) for reactions involving 74 the dissolution of molybdenite (M0S2)

Table 3. Selected equilibrium constants (logK) 76 1

GENERAL INTRODUCTION

Molybdenum is a transition element in Group VI of the Periodic Table.

Its atomic number is 42, and its spherical electron distribution is 4d^5s^.

This electronic configuration makes molybdenum unique in terms of its

chemical properties. It can exist in a variety of oxidation states from -2

to +6, and can coordinate with 4 to 8 neighboring atoms (Mitchell, 1973).

Molybdenum readily reacts with most inorganic and organic ligands, forming

a broad spectrum of monoraeric or polymeric compounds.

Although molybdenum was discovered in 1778, its commercial

significance was not recognized until the 1920s. Since then, molybdenum

has gained wide applications in modern industry despite the fact that

metallurgical uses still constitute the major fraction of molybdenum

consumed. Molybdates are now largely replacing toxic chromâtes as

corrosion inhibitors because of their nontoxic properties. Sodium

molybdate, for example, is found to be more effective than the traditional

chromate corrosion inhibitor when added to solutions in cooling towers and

automobile cooling systems. Other molybdenum compounds, such as zinc and

calcium molybdates, are widely used in places where chromium- and lead-

based inhibitive pigment were employed for equipment and facilities in the

food, beverage, drug, and cosmetic industries (Lander, 1977).

In addition to their role as paints and corrosion inhibitors,

molybdates, molybdenum sulfide, molybdenum oxide, and some

organomolybdenum compounds also serve as important catalysts in the oxidation reactions involving organic substances, in hydrodesulfurization 2

and denltrification of crude oils in the petroleum- processing industry,

and in the gasification and methanation of coal, and in coal liquification

schemes as well (Lander, 1977). Furthermore, the sulfide (M0S2) has been

extensively used in lubricants either as dry lubricant or in colloidal

suspension in oils and greases at high temperatures at which many oils are unstable. This lubricating action arises from the unique structure of M0S2 which consists of layers of MoSg units with molybdenum in trigonal prismatic coordination by sulfide ions (Mitchell, 1973).

The uses of molybdenum in metallurgy are even more diverse. Because of its very high melting temperature (2610 °C), molybdenum is being used in electrodes for furnaces and other electrical equipment in which tungsten has more commonly been used. Molybdenum alloys are not only of improved strength but also of better corrosion resistance. Therefore, they are widely used for high speed tools, high temperature parts in jet engines, space rocket motors, gas turbines, and nuclear reactors (Lander, 1977;

Ross, 1972).

In biochemistry and agriculture, the uses of molybdenum are also expanding. The role of molybdenum in nitrogen fixation and nitrate reduction in plants has received increasing attention in recent years. As a trace element in animal and plant tissues, molybdenum is believed to influence substantially biochemical processes, such as the synthesis of proteins (Mitchell, 1973).

Despite its broad utility, however, there are few places where ores of molybdenum are currently being mined. Molybdenum is widely distributed in the Earth's crust in trace amounts. Its average concentration in different rocks varies considerably, from 0.2 to 27ppra with an average of 2.5 ppm in 3

silicic rocks (Turekin and Wedpohl, 1961). In sedimentary rocks,

molybdenum is preferentially accumulated in organic-rich sediments. Thus

molybdenum is often found to be concentrated abnormally in coal beds,

lignite sandstone units, uraniferous sandstones, vanadiferous shales,

phosphate beds, and black shales (Finch, 1967; Gulbrandsen, 1966;

Valkovic', 1983; Ruzicka and Bell, 1984). In some of the rocks, molybdenum

can be as high as several thousands of ppm (Armands, 1972; Coveney and

Martin, 1983). In spite of its wide distribution, molybdenum rarely forms commercial deposits in sedimentary rocks.

Molybdenum has been found in thirteen minerals (Killeffer and Linz,

1952), among which molybdenite (M0S2) is by far the most important one. It occurs in a variety of geological environments and mainly concentrates in endogenetic ore deposits. Some other molybdenum minerals occurring naturally are wulfenite (PbMoO^), powellite (CaMoO/^), ferrimolybdite

(Fe2(Mo0^)37H20), and ilsemanite (Mo30gnH20), which appear only in oxidized zones of hypogene ores, and are of minor significance.

Molybdenite is currently mined mainly from porphyry molybdenum deposits and as a by-product from porphyry copper deposits. Some other major type of deposits which are of less importance are skarns, quartz veins, greisens, and pegmatites. Almost all the deposits are formed by hydrothermal activities, as evidenced by the intimate relations between hydrothermal alteration and molybdenum mineralization, and by the information generated from fluid inclusion and isotopic studies.

Porphyry molybdenum deposits are the most important source of molybdenum. They occur in stockworks, veinlets, and disseminations in granitic to monzonitic intrusions. Based on the nature of associated 4

rocks, porphyry Mo deposits are subdivided into Climax type deposits and

quartz monzonite deposits (White et al., 1981). Climax type deposits are

associated with high silica, alkali-rich rhyolite, and granite porphyries

that intrude into Precambrian gneisses. The orebodies are stockworks

containing thin quartz veinlets with molybdenite occurring along the walls

(Wallace et al., 1968, 1978; White, 1981). Quartz monzonite deposits are

associated with rocks ranging in composition from biotite granite to quartz

monzonite or granite porphyries. Molybdenite occurs as fracture coatings

and as disseminations in quartz-molybdenite veinlets (Hudson et al., 1979,

1981).

Molybdenum-bearing skarns commonly occur as copper-molybdenum and tungsten-molybdenum skarns (Pokalov, 1977). However, the intrusive rocks associated with molybdenum skarns are usually more felsic in composition than those associated with copper or tungsten skarns. They are developed along the contact between plutons, which are commonly small in volume, and carbonate or calcareous sedimentary rocks. Molybdenite is found as a disseminated phase in the skarn or in quartz veinlets along with tungsten or copper minerals.

Quartz vein type molybdenum deposits are characterized by the occurrence of subparallel fissure-filling quartz veins in which molybdenite is disseminated. The veins vary in width from several centimeters up to several meters. Besides molybdenite, wulframite and scheelite are also present in the veins. The associated igneous rocks are granite or quartz monzonite that were emplaced into sedimentary, volcanic, and Precambrian metamorphic rocks (So et al., 1983a, 1983b; Shelton et al., 1987).

Both greisen and pegmatite deposits are relatively unimportant as 5

sources of molybdenum due to their small volume and hence small tonnage,

although they are usually of high grade and easy to process. Greisens are

typically formed by greisenization of their host granites, which are

leucocratic, and are high in silica, Al, and alkaline elements, low in Mg

and Fe contents. They often occur as thin layers along the contact between

granitic intrusions and low grade metamorphic rocks (Ivanova et al., 1978).

Molybdenum-bearing pegmatites are usually small in size and associated with

granitic rocks. They often occur in the border zone of the intrusives, and

also in the invaded rocks. In most of the occurrences, molybdenite is the

only sulfide that forms large flakes and crystals in the dykes close to the

edge of the main granites (Vokes, 1963).

A common feature of these ore deposits is the extensive hydrothermal

alteration associated with them. This alteration is best seen in

porphyritic systems where a large volume of pervasively altered rocks are developed. In contrast to porphyry deposits, hydrothermal alteration in quartz vein type deposits is highly localized; it occurs mainly in the vicinity of the vein itself. The most common alteration includes K- feldsparthization, albitization, propylitization, greisenization, argillization, and phyllic alteration.

Fluid inclusion studies indicate that molybdenum deposits are formed in a wide temperature range from 250 to 650 °C, with most of them falling between 300 and 450 °C. The ore-forming fluids are apparently saline solutions that contain 3 to 56 eq. wt% NaCl (Roedder, 1971; Hall et al.,

1974; Ivanova et al., 1978; Kamilli, 1978; Bloom, 1981; MacDonald and

Spooner, 1981; White et al., 1981; Shelton et al., 1983, 1986, 1987; So et al., 1983a, 1983b; Lang and Eastoe, 1988). Pressures estimated from fluid 6

inclusion studies are between 100 and 1700 bars, and boiling apparently

took place during mineralization (Ivanova et al., 1978; White et al., 1981;

So et al., 1983a, 1983b; Shelton et al., 1986, 1987). Stable isotopic data

reveal that mixing of magmatic fluid and meteoric water also took place

during the hydrothermal processes (Hall et al., 1974; Shelton et al., 1983,

1987).

Despite the richness of literature about the physico-chemical

conditions under which molybdenum deposits were formed, little work has

been done to study the behavior of molybdenum under these conditions. This

knowledge is essential for us to understand the origin of molybdenum

deposits and to put a bound on future exploration. In this investigation,

the solubilities of molybdenum oxide (M0O2) and sulfide (M0S2) have been

determined in hydrothermal solutions under both sulfur-free and sulfur-rich conditions to elucidate the transport and deposition mechanism of molybdenum in natural systems.

This thesis is divided into two chapters, each of which is autonomous.

Chapter I deals with the solubility of M0O2 in NaCl solutions between 300 and 450 °C with fQ2 buffered by magnetite-hematite. A speciation scheme for aqueous molybdenum species is proposed, and the pertinent thermodynamic data are derived based on solubility measurements. The relative stability of each molybdenum species under geologically relevant conditions is also discussed. Chapter II describes the solubility of molybdenite (M0S2) as a function of temperature, NaCl concentration, fQ2, fg2, and pH in hydrothermal solutions. A dissolution or precipitation mechanism of molybdenite is hypothesized based on the speciation scheme derived in

Chapter I. A comparison is made between solubilities observed from 7 experiments and those by thermodynamic modelling. Together, these two papers provide us with some insights on the behavior of molybdenum at elevated temperatures and pressures. 8

CHAPTER I.

DISSOLUTION OF MoOg AND COMPLEXING OF MOLYBDENUM IN NaCl SOLUTIONS

AT ELEVATED TEMPERATURES AND PRESSURES 9

ABSTRACT

The solubility of M0O2 was measured in NaCl solutions at temperatures

between 300 and 450°C. Experiments were performed under water-saturated

vapor pressure conditions using Barnes vessels with capacities of 0.5 and

1.1 liters. NaCl concentration varied from 0.0316 M to 5.7 M, and oxygen

fugacity was controlled by the Fe30^-Fe203 buffer.

Run solutions, with quench pH mostly in the range of 4 to 5, were

characterized by a blue color, indicating a mixed oxidation state of Mo(5+)

and Mo(6+). Depending on the NaCl concentration, Mo content changes from

several tens of ppm to over 3000 ppm in the solutions. The strong

dependency of M0O2 solubility upon NaCl concentration, especially at

temperatures above 350 °C, suggests a relationship between NaCl and

molybdenum complexing.

Given some constraints on the system (i.e., mild acidity of quench

solutions and coexistence of Mo(5+) and Mo(6+)), a speciation scheme was

devised for aqueous molybdenum complexes based on solubility data.

According to this calculation, HMoO^" is important only at temperatures below 300 °C. With increasing temperature, ion pairs of NaHMoO^° and

Na2MoO^° are the dominant species for Mo(6+), while MoO(OH)2Cl is the dominant Mo(5+) complex. 10

INTRODUCTION

The ubiquitous association of extensive hydrothermal alteration with

molybdenum deposits strongly suggests an intimate relationship between

hydrothermal activity and ore deposition. Although there is a large volume

of literature to date describing the physicochemical conditions under which

molybdenum ores are formed (Hall et al., 1974; Bloom, 1981; White et al.,

1981; So et al., 1983a, 1983b; Shelton et al., 1986, 1987), little is known

about the chemistry and thermodynamics of molybdenum species in

hydrothermal fluids. Considerable controversy exists as to the form that

molybdenum takes during transport and the nature of the fluid capable of

transporting molybdenum. Knowledge of molybdenum speciation and its

dependency on the environment is essential to understanding the behavior of

molybdenum during hydrothermal processes and the thermodynamic modelling of

ore genesis in terms of transport and deposition mechanisms.

Molybdenum commonly exists in an oxidation state of 6+ in aqueous

solutions at room temperature, and can coordinate with six neighboring

atoms. Depending on the concentration of Mo and the pH of solution,

molybenum can readily undergo a range of hydration, protonation, and

polymerization reactions, forming a wide spectrum of complexes (Avesten et

al., 1964; Sasaki and Sillen, 1967; Nabivanets, 1969; Haight and Boston,

1973). It is generally accepted that monomeric complexes predominate only when molybdenum concentration is below lO'^M in solution, although

Podshivalova et al. (1984) have shown that monomers can persist in solutions with Mo concentration up to 0.1 M, based on a theoretical calculation. In neutral to mildly alkaline solutions, MoO^^' is the 11

dominant species. Upon acidification, however, a succession of other

Mo(6+)-bearing species is formed:

M0O42- ..> HM0O4" > H2M0O4O > HM0O3+ > MoOgZ*

(Nabivanets, 1969). Mo(6+) can also react with CI" and F', generating a

series of oxy-chloro and oxy-fluoro complexes in strong acid solutions

(Griffith and Wickins, 1967; Griffith and Lesniak, 1969; Karyakin and

Kryachko, 1967; Shapiro and Volk-Karachevskaya, 1969). The possibility of

forming these complexes in natural systems has been ruled out due to the

extreme geological conditions required. Molybdenum species with lower

oxidation state in aqueous solutions have not been reported in geological

literature, although they have been isolated and described by chemists

(Haight et al., 1973; Averill et al., 1980).

Studies of the behavior of molybdenum in hydrothermal fluids are

conflicting. Ryabchikov et al. (1981), who studied the partitioning of molybdenum between granitic melt and coexisting fluid, suggested that molybdenum is transported in magmatic fluid in the form of oxy-chloro or oxy-fluoro complexes, since the concentration of Mo in the fluid is dependent on the concentration of CI in the system. Khitarov et al. (1967) had earlier obtained a similar result in their study of partitioning of molybdenum between gas and aqueous phase at temperatures up to 350 °C.

Candela and Holland (1984) reported, however, that the partition coefficient for molybdenum between vapor and liquid is independent of Cl content in the fluid phase. Hence, they concluded that Mo02(OH)2°, instead, is the stable complex in magmatic fluids. This appears to be in 12

agreement with the calculations by Smith (1983) that suggest that molybdate

and molybdic acid are the dominant species in hydrothermal fluids. It is

difficult to justify the factors that cause such a discrepancy, but

temperature and oxygen fugacity effect seem among the most important ones,

Information generated from solubility studies of geological Interest

is scarce. Ivanova et al. (1975) conducted experiments on the solubility

of M0O3 in pure water and in NaOH solutions at temperatures up to 300 °C.

They found that the solubility in pure water increases slightly with

temperature, but drastically with NaOH concentration. They did not make it

clear what species contributed to the high solubility, but claimed the

importance of polymers at temperatures below 200 °C. The low quench pH of

their run solutions evidently indicates the presence of molybdate or

bimolybdate in their experiments if only monomers are consided. Kudrin

(1986) has recently studied the solubility of M0O2 in pure water and in aqueous HCl, NaOH, and KOH solutions in a wide range of oxygen fugacity at temperatures from 250 to 450 °C. He concluded that molybdenum is transported as HMoO^' and H2MoO^° in hydrothermal fluids. Although Kudrin did not give it much attention, his data also apparently indicate the presence of a Mo chloride complex in his low pH experiments. It is interesting that the point at which the HCl dependency of M0O2 solubility begins tends to move toward neutral pH as temperature is raised (e.g. pH-1.5 at 350 °C, and pH-3 at 450 °C). This is in contrast to the extreme conditions reqired to form molybdenum chloride complexes (HC1>4.0 M) at room temperature, suggesting that higher temperature favors the formation of chloride complexes.

Zhidikova and Malinin (1972) measured the solubility of powellite 13

(CaMoO^) in sodium chloride solutions at 50 to 300 °C. They observed an

increased solubility of powellite with NaCl concentration and attributed

this behavior to the change in ionic strength of their solutions. Thus, a

form of MoO^^" was proposed as the main species in their solutions. More

recently, a direct measurement of the solubility of molybdenite (M0S2) was

provided by Wood et al. (1987) who studied the solubility for ten minerals

simultaneously in H20-NaCl-C02 solutions from 200 to 350 °C. They believe

that molybdic acid predominates in their solutions because the solubility

is independent of CI concentration and pH, as well as temperature.

In spite of the works described above, no effort has so far been

directed to clarify the speciation of molybdenum in NaCl solutions, a more

realistic analogue of hydrothermal fluid. The physicochemical conditions

of ore formation, as deduced from fluid inclusion and stable Isotope

studies, are such that the major ore deposition takes place around 300-550

°C and pressures of 200-1800 bars, and that the salinity of the fluids,, expressed as a weight percent equivalent of NaCl, varies from 5 to 60 wt.%

(Hall et al., 1974; White et al., 1981; Bloom, 1981; So et al., 1983a,

1983b; Shelton et al., 1986, 1987). Molybdenum is apparently transported in fluid that is poor in total dissolved sulfur or sulfide (Walder et al.,

1988). In light of this condition, experiments in this study were conducted on the solubility of M0O3 in various NaCl solutions at temperatures between 300 to 450 °C in an attempt to better clarify the Mo speciation in hydrothermal fluids. M0O3 is better than its sulfide counterpart as the source of aqueous molybdenum because the solubility of

M0O3 is much higher than M0S2, making the analysis easier and more reliable. In addition, data generated in this system represent the 14 behavior of molybdenum in a sulfur-free system, which can be used in conjunction with those in asulfur-rich system to infer the behavior of Mo in the entire history of hydrothermal processes. Oxygen fugacity in all the experiments was buffered by the assemblage magnetite plus hematite, which was also used to determine HCl concentration in solutions since the solubilities of Fe oxide are well established (Chou and Eugster, 1977;

Crerar et al., 1978; Boctor et al., 1980). EXPERIMENTAL METHODS

All the solubility measurements were carried out along the saturated

water-vapor pressure curve in two continuous-sampling vessels of the type

described by Barnes (1963, 1971). One chrome-lined vessel with a capacity

of 1.1 liter was used in the lower temperature range 300-350 °C; the other,

with capacity of 0.5 liter, was used at temperatures between 400 and 450

°C. A sampling tube and valve system placed on top of the vessel allowed

direct sampling of run solution at temperature. Pressure was measured by

fluid-filled bourdon-tube gauges with a full scale of 15,000 and 50,000 psi

respectively. An isolator was placed between the gauge and capillary line

to protect the gauge and to minimize the possibility of contaminating the

fluid inside the gauge with brine from the pressure vessel. Run

temperature was monitored and kept constant (± 1 °C) by solid-state controllers which read thermocouples at the top and bottom of a reaction vessel. The measured temperature gradient from top to bottom of a vessel was less than 2.0 °C.

The solid oxides magnetite (FegO^) and hematite (2^20]), M0O3, and

M0O2 were purchased from Alpha Company. NaCl was supplied by Fisher, ACS certified grade. All materials were used without further purification.

Doubly distilled water was used in the preparation of all synthetic fluids.

Depending on the vessel used, 15 to 25 grams each of magnetite and hematite, along with about 6 grams of M0O3 were loaded into the vessel, which was then evacuated and filled with 300 to 600 ml of NaCl solution. A typical run time at selected temperature was 7 to 20 days. Although M0O3 dissolves very rapidly, the equilibration among the buffer phases is 16

relatively slow, as demonstrated by a continuous decrease in Mo

concentration and a gradual increase of Fe content in solution in the first

several days. Therefore, we assumed an apparent equilibration was achieved

when both Mo and Fe concentrations stayed unchanged over a period of time

at constant temperature or when the values could be approached from both

the supersaturated and undersaturated states (Crerar and Barnes, 1976).

After equilibration, sample solution was extracted from the reaction

vessel at temperature and pressure via the tubing and valves system, as

described by Barnes (1971). The sampling system was first rinsed with

doubly distilled water and blown dry with pressurized N2 gas, then followed by flushing with at least three 5 ml. aliquots of the run fluid.

Afterwards, two samples (about 5 ml aliquots each) were withdrawn. One was for pH measurement. The other was filtered using a sintered glass and membrane filter paper with pore size of 0.8 um, diluted with either NH^OH or distilled water and transfered to a Nalgene bottle for later analysis.

Molybdenum was analyzed colormetrically with a Beckman electron photometer using the dithiol method of Clark and Axley (1955), which is sensitive to less than 0.03 part per million Mo. Fe and Na were analyzed on a Perkin-Elmer model 303 atomic absorption spectrometer using an acetylene-air flame. Chloride was analyzed using a Corning chloride electrode.

Solid phases were X-rayed and examined under a microscope to ensure that all original materials were present after the vessel was opened. Both methods confirm that magnetite and hematite persist throughout most experiments. There were several cases, however, in which all the magnetite was oxidized into hematite. These runs were discarded from consideration. 17

700 K 1.Okb

—10

FeS2

-30 -10 —8 -6 -4 -2 LogfS2

Figure 1. Phase relations in the system Mo-Fe-O-S at 700 K, MoOg is the stable phase with respect to M0O3 under conditions where fgn is buffered by the assemblage Fe203 - Fe^O^ 18

Identification of the Mo-bearing mineral was not easy due to its low

abundance in the system. Nevertheless, powder XRD analysis illustrates

that M0O3 was converted to M0O2 during the experiments, except when loss of

magnetite indicated a high oxidation state. The rate of this transition

seemed very high, and the process was completed within several days. To

verify this, a few runs were made with M0O2 as the starting material.

After equilibration, the solubility was found to be in excellent agreement

with that in runs with initial M0O3, and M0O2 was unambiguously identified

in the final products. This confirms that M0O2 was the stable phase under

the experimental conditions, as predicted from thermodynamic calculation

(Figure 1). EXPERIMENTAL RESULTS AND CONSTRAINTS ON SOLUTION MODELLING

The solubility data of molybdenum dioxide and the run conditions are

summarized in Table 1. Figures 2 to 5 are plots of log total molarity of

molybdenum against log total sodium chloride concentration in solutions at

temperatures from 300 to 450 °C. The error of analysis is much smaller

than the symbol size. As can be seen from Table 1, the overall solubility

of M0O2 is very high, ranging from several tens of ppm to over 3000 ppm,

increasing with NaCl concentration in the experiments. These values, however, are still significantly lower than that obtained by Ivanova et al.

(1975) in their study of the solubility of M0O3. At 300 °C, for example, molybdenum concentration in this study is 700 ppm in 3.0 m NaCl solution, while in Ivanova's work, the Mo concentration is about 1400 ppm, even in pure water. This once again confirms the finding that we were dealing with the solubility of M0O2 in spite of the fact that M0O3 was the starting material employed, as described in the previous section.

It is worthwhile to note that the slope of the plots in Figures 2 to 5 increases with temperature, from 0.4 at 300 °C to 0.8 at 400 °C, although it apparently becomes constant at temperatures above 400 °C. This behavior, along with the clear NaCl dependency of M0O2 solubility, cannot be interpreted as an ionic strength (or activity coefficient) effect. It must be the NaCl that enhanced the solubility of M0O2 by complexing with it in some way.

Interpretation of molybdenum speciation based on solubility alone is a very difficult task, given the complexity of the system and the large number of potential stoichiometric molybdenum complexes present. We cannot 20

Table 1. Experimental solubility data for M0O2 in NaCl solutions buffered by magnetite-hematite. Concentrations in molarity

T °C Run # Duration (days) ™NaCl Quench pH %o

15-1 300 15 0.0356 4.50 0.00159 15-4 300 11 0.0371 4.15 0.00087 12-2 300 13 0.117 4.45 0.00280 12-6 300 8 0.130 4.30 0.00106 18-2 300 14 0.731 4.56 0.00430 18-3 300 20 0.751 4.55 0.00379 18-6 300 8 0.746 4.51 0.00264 10-3 300 11 1.620 4.65 0.00500 21-2 300 16 2.920 4.87 0.00730

15-2 350 7 0.092 4.30 0.00075 15-3 350 15 0.042 4.43 0.00088 15-5 350 11 0.0545 4.44 0.00097 12-3 350 7 0.163 4.31 0.00190 12-4 350 13 0.197 4.41 0.00185 12-5 350 18 0.238 4.44 0.00197 18-4 350 7 0.981 4.60 0.00565 18-5 350 14 0.981 4.61 0.00535 18-7 350 6 1.070 4.70 0.00510 10-1 350 7 1.968 4.81 0.01090 10-2 350 13 2.094 4.83 0.01100 21-3 350 8 3.423 5.02 0.01340 21-4 350 17 3.484 5.10 0.01410

14-2 400 19 0.0329 4.69 0.00100 14-5 400 10 0.0272 4.50 0.00032 7-2 400 15 0.100 4.43 0.00150 13-1 400 6 0.312 4.64 0,00540 13-2 400 12 0.318 4.64 0.00460 5-1 400 7 1.126 4.74 0.00900 5-2 400 12 2.060 4.90 0.01470 11-1 400 7 3.610 4.99 0.02840 11-2 400 14 3.660 5.10 0.02970 8-1 400 10 4.340 5.07 0.02130 8-5 400 13 5.252 5.24 0.03070

14-3 450 6 0.0342 4.45 0.00094 14-4 450 14 0.0342 4.31 0.00059 13-3 450 7 0.333 4.83 0.00495 8-2 450 10 5.400 5.07 0.03290 8-3 450 11 5.500 5.14 0.03090 21

-1 T=300 T MH

-2 -

0 +)

07 0 -3 -

R - 0.8988 Y - -2. 380E+00 + <4.G066E-01) * X —4 J I I I I -2 -1 0 Log CNaCl (tot)]

Figure 2. Plot of log[Mo(tot)] vs. log[NaCl(tot)] illustrating dependence of M0O2 solubility on chloride concentration at 300 °C with fQ2 controlled by MH. Solid line represents a least-squares fit to the data 22

-1 T=350 T MH

-2 - +> 0

w 0

07 0 -3

R - 0.9940 Y - -2.221E+D0 + C6. 5669E-01) * X -4 I I -2 -1 Log CNaCl (tot)]

Figure 3. Solubility of M0O2 in NaCl solutions at 350 °C and fQ2 controlled MH. Solid line represents a least-squares fit to the data 23

-1 T=400 °C MH

-2 — •P 0 +) V/

07 0 -3 —

R - 0.9860 •2. 032E+00 + (7. 8977E-01) * X —4 _L -2 -1 0 LogCNaCl (tot)]

Figure 4. Plot of log[Mo(tot)] vs. log[NaCl(tot)] demonstrating dependence of M0O2 solubility on chloride concentration at 400 °C and fQ2 controlled by MH. Solid line represents a least-squares fit to the data. Note the increase in slope of the solid line with respect to temperature in the range from 300 to 400 °C 24

-1 T=450 T

-2 - +> •p0

0 z 1—1 07 0 -3 -

R - 0.9948 -1.984E+00 + (7.6340E-01) * X -4 I I I , -2 Log CNaCl (tot)1

Figure 5. Solubility of M0O2 as a function of NaCl concentration in NaCl- H2O solutions at 450 °C with fQ2 buffered by MH. Solid line represents a least-squares fit to the data 25

assume the presence of a single dominant species although this assumption

has been applied successfully in some other systems (Chou and Eugster,

1977; Frantz and Popp, 1979; Eugster, 1985), because of the apparent

contradiction with observed phenomena from this study. The relevant

phenomena fall into two categories: those related to pH, and those related

to oxidation state. The discussions which follow are based on data in

Table 1.

Deductions Based on Solution PH

As shown in Table 1, the quench pH decreased with respect to the

almost neutral pH of initial run fluids, indicating that H"*" was generated

during the experiments. Unlike what is observed by Wood and Vlassopoulos

(1989) in the tungsten system, where the lowest quench pH is always

accompanied by the highest metal content, the quench pH in this study increases with molybdenum concentration in solution, although the overall change is only about 1.0 unit. If the solubility of M0O2 is solely due to the reaction

MoOg + H2O + 1/2 O2 — + (2-x)H+

where x=0, 1, or 2, the quench pH then should be either neutral (when x=2) or should have a value equivalent to the magnitude of the log molarity of molybdenum found in the same solution (when x-O or 1). This is, of course, not the case.

By the same token, the possibility of forming an ion pair involving 26

sodium and either molybdate or bimolybdate as the only dissolution

mechanism can also be ruled out because a low quench pH would be generated:

M0O2 + xNaCl + H2O + 1/2 O2 — NajjH2.xMo04° + xHCl

where x-1 or 2.

Also, it is unlikely that the quench pH was affected significantly by

the dissolution of the Fe-0 buffer assemblage because the Fe content in

solutions is much lower than that of Mo. This is true even though as much

as twice the Fe content of may have been consumed by the dissolution of

iron oxides (Chou and Eugster, 1977; Doctor et al., 1980):

FEGO^ + 2H+ + NCL- — FEGOG + FECL^Z'" + HGO

Therefore, a trend of decreasing pH with increasing Mo concentration would

be expected if the dissolution of M0O2 was governed by the reactions described above.

Reactions involving the formation of an oxy-chloro complex are also possible:

MoOg + 2(n-l)H+ + nCl* + 1/2 Og MoO^.^Cl^^^^ + (n-ljHgO

These reactions, however, consume and would shift the quench pH into either the neutral (when n-1) or alkaline region (when n>l), contradicting the experimental results.

Similarly, any reaction involving Mo(5+) species alone cannot explain 27

the behavior of quench pH in these experiments. It seems that at least two

reactions have taken place simultaneously in the solution. One generates

or HCl, while the other either consumes or is independent of H"*".

Deduction Based on Oxidation State

Another distinct phenomenon is the dark blue color associated with

quench solutions. This blue color can persist for as long as two days

standing in the open air at room temperature, then fades to yellow.

Addition of concentrated HCl tends to enhance the color, while a few drops

of H2O2 decolorize the solution immediately. The degree of the color is

apparently a function of molybdenum concentration in solutions; the higher the Mo content, the deeper the color.

This phenomenon has been described by chemists (e.g., Killeffer and

Linz, 1952) as "molybdenum blue". Although the true nature of molybdenum blue is still uncertain, it is generally believed that it results from the mixing of Mo(5+) and Mo(6+) species. This phenomenon occurs in mild acid solutions (at pH near 4) as an intermediate product when Mo(6+) is reduced by a reducing agent or when Mo(5+) is oxidized by an oxidizing agent

(Killeffer and Linz, 1952). Mixing of Mo(5+)-bearing and Mo(6+)-bearing solutions also gives rise to a blue-colored solution as long as the pH is kept near 4.0. The blue solution was easily produced in our lab at room temperature when sodium molybdate was reduced by iron in dilute HCl solution. The blue color was immediately diminished upon the addition of

H2O2. A similar phenomenon was also observed in some tungsten-bearing systems (Wesolowski et al., 1984; Manning and Henderson, 1984; Wood and 28

Vlassopoulos, 1989). Unfortunately, there has been no spectroscopic study which might let us use this observation as a quantitative measure of the

Mo(5+)/Mo(6+) ratio. The existence of mixed oxidation state, however, is itself a valuable constraint on our solution modelling effort. 29

ESTIMATION OF MOLYBDENUM SPECIES

The discussions in the previous section strongly suggest that there

are at least two kinds of molybdenum species in our solutions, one with

oxidation state of 6+, the other with 5+. The presence of more than one

Mo-bearing species in the system severely hinders the effort to estimate

the stoichiometry of molybdenum species. During this study, we made some

preliminary attempts to identify molybdenum species at elevated

temperatures directly by using a "window bomb" designed on principles

outlined by Susak (1981). Our spectroscopic investigation showed us

quickly, however, that there are few distinct absorption features in the UV or Visible range that can be used in a reliable, reproducible fashion.

Clearly, this line of study deserves more attention and may ultimately yield valuable insights on the system we have been studing here.

Considering the great effort necessary, however, we have chosen to set that phase of the problem aside and to focus on deductions that can be made on the basis of thermodynamic manipulations of solubility data. As shown below, this approach allows us to recognize several valuable constraints on speciation.

To begin, we need to postulate a general form for each molybdenum species of different valance state. Regardless of the possibility of forming polymeric species, monomers must be assumed in order to make the calculation possible. This limitation is actually quite reasonable, since depolymerization tends to predominate at high temperatures and high salinity in run solutions will inhibit polymerization (Wesolowski et al.,

1984). Given the bulk compositions in our experiments, the most likely forms for hexavalent molybdenum are Hj^MoO^*'^, Na^H2.^MoO/;^°, and

Mo0^.JJ.C1JJ*'2, where x-0, 1, and 2. Higher values of x, though possible,

are less likely because the resulting complexes are large and easily

destabilized. Difficulty arises when we try writing a general form for

Mo(5+) complexes. One of the profound properties of Mo(5+) in aqueous

solution is its strong tendency to form dimers (Haight and Boston, 1973).

The most well-known monomeric species is MoOClg^", which is stable only

under extreme conditions (HC1>9 M). In view of the existence of solid

MoO(OH)3 (Killeffer and Linz, 1952) and M0OCI3 (Sillen and Martell, 1964),

we postulate that the form for Mo(5+) is MoO(OH)3_yCly®, which should

bracket most of the hydroxy or chloro molybdenum complexes with oxidation

state of 5+.

Having postulated the general forms, the next step is to determine the

stoichiometric coefficients x and y in their formulas. This can be done by

fitting the solubility data with respect to the corresponding ligand using

a multivariant linear regression technique (Crerar and Barnes, 1976; Wood

and Crerar, 1985). This method, however, is sometimes misleading,

especially when the variance is large, and may yield an unexpectedly large

number of species, making thermodynamic modelling almost impossible.

Another way to do this is by prejudicially assuming one dominant

species each for Mo(5+) and Mo(6+). If this method yields results that are

chemically realistic, then we have identified the simplest species mixture that can account for observed solubility behavior. More complicated mixtures are not ruled out by this procedure, but at least we will discover whether they are even necessary. Because there are three different forms for Mo(6+), three pairs of species must be considered in the calculation: 31

case /

Consider first the pair MoO(OH)3_yCly° - Na^H2_xMo02^°. The major

unknown quantities in the system are x, y, and the concentrations of NaCl°,

HCl®, NaOH°, Na+, Gl", H+, OH", Mo(V), Mo(VI), and FeClg®, in which the neutral ferrous chloride is assumed as the dominant Fe-bearing species for

simplicity (Chou and Eugster, 1977; Crerar et al., 1978; Boctor et al.,

1980). The following equations can be written for the system

NaCl° — Na"*" + CI" (1)

NaOH° — Na+ + OH" (2)

HCl® -- H+ + CI" (3)

HgO -- H+ + OH" (4)

FegO^ + 2HC1° — Fe203 + FeClg® + HgO (5)

NaT - [NaCl°] + [Na+] + [NaOH®] + x[Mo(VI)] (6)

GIT - [NaCl®] + [C1-] + [HCl®] + 2[FeCl2°] + y[Mo(V)] (7)

[Na+] + [H+] - [C1-] + [0H-] (8)

FeT - [FeClgO] (9)

MoT - [Mo(V)] + [Mo(VI)] (10) where the brackets denote molal concentrations of the corresponding species. The activity coefficients will be defined as unity for all the species throughout this calculation.

Equations (l)-(5) and (7)-(9) can be solved simultaneously for NaCl°, 32

NaOH®, HCl®, FeClgO, Na+, H+, CI", and OH*. In eq. (7), ytMo(V)] can be

omitted as a first approximation because it is very small compared to CIT

{<11 in most cases), and accurate values can be obtained by iteration later

when y[Mo(V)3 is known. The equilibrium constants for reactions (1) to (5)

are given in table 2.

Assume that the ion pair we are studying is produced by

M0O2 dissolution in the following fashion:

MoOg + 1/2 Og + HgO + xNa+ — NaxHg.xMoO^o + (11)

[Mo(VI)][H+]X

" f02l/2[Na+]X and

MoOg + 1/4 Og + 0.5(3-2y)H20 +yH+ + yCl" — MoO(OH)3.yCly° (13)

[Mo(V)] ^13 —— (14) fo2 / [H i^ici'jy

Combining equations 6,7,and 10, yields

1 [Mo(VI)] - ([Cl-]+[HClO]+2[FeCl°]+y[MoT]-[Na+]-[NaOH°]) (15) X + y and

1 [Mo(V)] ([Na+]+[Na0HO]+x[MoT]-[Cl-]-[HCl°]-2[FeCl2°]) (16) X + y

By substitution of (15) into (12), taking the logarithm and rearranging, we obtain 33

Table 2. log K for some selected reactions used in the calculation of molybdenum speciation

Reaction 300°C 350Oc 400°C 450°C O M M C C NaCl® - Na+ + Cl" -0.94A -1.87* -2.0% C

NaOH® - Na+ + OH" -0.94° -1.87° -2.0° -2.82°

HC1° - H+ + Cl- .1.23d -2.53d -3.2® -5.00®

H2O - H+ + OH" -11.29? -11.91^ -12.28^ -13.48^

FegO^+jHClO - FegOg+FeClgO+HgO 4.868 5.162 5.42% 4.31%

^ Estimated from Wright et al. (1961) and Pearson et al. (1963).

^ From QUist and Marshall (1968).

° Assume the same as NaCl°.

^ From Jackson and Helgeson (1985).

® Calculated from Frantz and Marshall (1984).

^ From Marshall and Franck (1981).

S Calculated from Crerar et al. (1978).

^ Calculated from Doctor et al. (1980). 34

logKii - log(y[MoT]+[Cl-]+[HClO]+2[FeT]-[Na+]-[Na0HO])

+ xlog([H+]/[Na+]) - log(x+y)fo2^/^

At constant temperature and fQ2> * and y are independent of composition

in the experimental range. As long as x and y are not both zero, the

following relation.exists:

aiog(y[MoT]+[Cl-]+[HClO]+2[FeT]-[Na+]-[NaOH®]) X (17) aiog([Na+]/[H+])

In the same fashion, we have

logKi3 - log(x[MoT]+[Na+]+[Na0HO]-[Cl-]-[HCl°]-2[FeT])

- y log [H"*"] [CI"] - log(x+y)fo2^/^ and

31og(x[MoT]+[Na+]+[NaOH°]-[Cl*]-[HC1°1-2[FeT]) y - (18) ôlog([H+][Cl-])

Simultaneous solution of (17) and (18) will yield values for x and y respectively. However, an easy way to do this is by trial and error.

First, let y in (17) vary from 0 to 3. For each y, a plot of log(y[MoT]+[Cl-]+[HClO]+2[FeT]-[Na+]-[NaOHO]) vs. log([Na+]/[H+]) is made for all data gathered at a single temperature, and x is defined as the slope of the plot. Figure 6a shows an example of such kind of plot. At

400 °C, the slope x is found to be 1.9 when y-1, and there obviously is a linear relations between numerator and denominator in eq (17). Because monomeric species are assumed, as mentioned earlier, x must be an integer.

Thus, the slope is rounded to a whole number. The relationship between y 35

T=400 °C

y-1

R - 0.9621 -1.164E+01 + (1.8969E+00) * X -4 4 5 6 Log < [Na+] / [H+] )

Figure 6a. A representative plot of log u vs. log([Na+]/[H+]), where u-y[MoT]+ [Cl ]+ [HCl]+2[FeT] - [Na"*"] -[NaOH] , demonstrating the linearity of equation (17). When y-1, the value of x is given by the slope of the solid line which has a value of 1.9. Thus, an integer of 2 is selected for x 36

T=400 °C

1

x-2

-2

-3

R - 0.9696 5.3298E+00 + (1.0012E+00) * X -4 -g —8 -7 -6 LogCCH+3 CCI-3)

Figure 6b. A representative plot of log v vs. log([H+][Cl"]), where v-x[MoT]+ [Na'^] + [NaOH] - [CI" ]- [HCl]-2[FeT] , demonstrating the linearity of equation (18). When x-2, a value of 1 can be determined for y from the slope of the solid line 37

and the corresponding x in (17) is then plotted on a x-y coordinate, and

for simplicity, is defined as x-f(y). Similarly, for each given value of x

(from 0 to 2) in (18), log(x[MoT]+[Na+]+[NaOH]-[CI"]-[HCl]-2[FeT]) is

plotted against log[H^][Cl"] to find a value for y which is also rounded to

a whole number. The linearity of (18) and the way to assess the value of y

are demonstrated in Figure 6b for 400 °C. Under this specific condition, y

is found to be 1 when x-2. These x and corresponding y values are also plotted on the same x-y coordinate, and define y-f(x). The solution of the system of equations of (17) and (18) thus can be found at the intercept of the two curves x-f(y) and y-f(x).

Figure 7 illustrates the solutions of (17) and (18) at different temperatures. Table 3 lists all the species calculated using equations

(17) and (18), assuming two dominant species in the solutions. The concentrations of Mo(5+) and Mo(6+) at each temperature can be calculated using eqs. (15) and (16) respectively.

The validity of the speciation scheme, however, needs to be tested.

Two criteria are used for this purpose. One involves the prediction of some of the features observed in the experiments, such as quench pH, using the calculated species stoichiometries and their concentrations. The other is the comparison of the calculated average ligation number with that determined from experiments. The average ligation number L is defined as

x[Mo(VI)] + y[Mo(V)] L - (19) [MoT] where x and y are numbers of ligand bound to Mo(6+) and Mo(5+) 38

T-300°C T-350 ®C

CO, 2) x-f(y)

y"fCx) <1, 1) 1 - y-f (x)

x-f Cy)

T-400 °C T-450 °C

x-f(y) x-f Cy) 2 i r- 2 ip- y-f Cx) y-f Cx)

1 — 1 - C2, 1) C2, 1)

Figure 7. Plot of y vs. x, derived from equations (18) and (17) respectively, at 300, 350, 400, and 450 °C. Solution of the system of equations (17) and (18) at each temperature can be found at the intercept of the two curves y=f(x) and x=f(y), which is shown in parentheses 39

Table 3. Calculated molybdenum speciatlon scheme

T °C Mo(V) Mo(VI)

300 MoO(OH)Cl2° HMoO^-

350 MOO

400 MOO(OH)2C1° Na2MoO^°

450 MoO

respectively. L can be determined by plotting (x[Mo(VI)]+y[Mo(V)]) against

[MoT], from which L is given by the slope of the plot.

It can be shown that L can also be determined directly from solubility

measurement. Consider reaction (13) and the reaction

MoOg + 2(x-l)H+ + xCl" + 0.5 0% -- MoO^.xClxX'Z + (x-DHgO, (20)

for which

[Mo(VI)]

f02^/^[H"*"]2*'2[Cl"]* ^ ^

From (10), (14) and (21), we have

MoT - Ki3fo2^/^[H+]y + I'' , (22)

At constant fQ2 and pH

3log MoT [Cl'jaMoT

aiog[Cl"] MoT Ô[C1-]

[Cl-](yKi3fo2^/^[H+]y[Cl-]y-l + xK20fo2^/^[H+]2x-2[Cl-]X-l)

MoT

y[Mo(V)] + x[Mo(VI)]

MoT which is exactly the average ligation number L in (19). Therefore a plot of the logarithm of total molybdenum concentration in solution against the logarithm of CI" would yield information on the average ligation number of aqueous molybdenum species: 41

ôlog MoT L (23) Ôlog[Cl-]

It should be realized, however, that it is the free CI' which actually participates in the reactions. Thus CI", rather than total NaCl as proposed by Wood et al. (1985, 1987), must be involved in making the plot to find average ligation number L. Substitution of Na"^ for CI" in either term of eq. (22) will not invalidate (19) and (23), although in this case L cannot be strictly termed an average "ligation" number.

Attempts were also made to calculate speciation scheme for the other two pairs Mo0(0H)3.yClyO-H3jMo04*-2 and Mo0(0H)3.yClyO-Mo04.j^Clj^^"2.

However, in case 2, the results violate the average ligation number obtained from eq. (23); and the results in case 3 cannot predict the quench pH. 42

GEOLOGICAL APPLICATIONS

As discussed in the previous section, both Mo(V) and Mo(VI) can exist

stably in electrolyte solutions at high temperatures and pressures.

However, their relative stability and significance is apparently a function

of temperature, oxygen fugacity, pH, and total chloride concentration. At

temperatures below 300 °C, HMoO^" is the most important hexavalent

molybdenum species (see table 3) under experimental conditions. With

increasing temperature, a sequence of ion pairs are formed with NaHMoO^°

and Na2MoO^° stable at moderate temperature and above 400 °C respectively.

These forms are similar to those for tungsten in NaCl solutions reported by

Wood and Vlassopoulos (1989). While ion-pairs are the dominant molybdenum

species for Mo(VI) in the solutions, the chloro-molybdenum complexes are

the most important species for Mo(V), although further work needs to be

done to clarify their true identity.

The possibility of forming the chloro-molybdenum complex under geologically relevant conditions was demonstrated by Kudrin (1986), who studied the solubility of M0O2 in dilute HCl solutions at 350 and 450 °C at ambient oxygen fugacity. Regardless of the reliability of his way of achieving control over oxygen fugacity in his experiments, the dependence of M0O2 solubility on HCl concentration is evident. Fig. 8 shows the relationship between log(MoT/fQ2^/^) and log[HCl] at 450 °C and 0.5kb constructed from Kudrin's data, assuming an oxidation state of 5+ for molybdenum species because the oxygen fugacity was below the magnetite- hematite buffer for the majority of his experiments. As can be seen, the solubility curve is concave upward. At lower HCl concentration (HCl < 43

450 X 0. 5kb

(M O (*. S" _J

# o I ?

-1 X _L -4 -3 —2 —1 Log CHCn

Figure 8. Solubility of M0O2, normalized to fQ2 in HCI-H2O solutions at 450 C and 0.5 Kb. The shape of the solubility curve suggests a stepwise complexing of chloro molybdenum species. An oxidation state of 5+ is assumed for aqueous Mo species (see text for discussions). The value of n designates the ligation number, and a maximum value of 3 can be derived from the solubility curve. Constructed from the solubility data of Kudrin (1986) 44

O.OIM), the solubility curve is flat, suggesting the presence of H^MoO^o in

the solution (Kudrin, 1986) if Mo(6+) is the major species, or MoOCOH)^^ if

Mo(5+) is predominant. Increasing HCl concentration, however, makes the

solubility increase dramatically. The shape of the curve seems to suggest

a stepwise complexing of Mo with chloride, and a maximum ligation number of

three can be derived from the plot.

It is important to note that the pH (about 3) as well as the HCl

concentration (around 0,1 M) required to form a chloro-molybdenum complex

at elevated temperatures and pressures can easily be encountered in

geochemical systems. Holland (1972), for example, reported a very low

quench pH (1 to 2) in his solutions in equilibrium with granite magma. My

own unpublished data verify that up to 0.1 M HCl can be generated in the

fluid phase in equilibrium with granitic melt in the system granite-NaCl-

H2O at 780 °C and Ikb through the reaction

2 NaClmgit + HgOgiuia " 2 HClgi^ia + ^220^21%

(Burham, 1979; Eugster, 1985, 1986). An even higher value of HCl can be expected in a real magraatic-hydrothermal processes. The high CI content in hydrothermal fluids would also undoubtedly enhance the formation of chloro molybdenum complexes.

Figures 9a-9d are logfQ2-pH diagrams depicting the stability field of aqueous molybdenum species from 300 to 450 °C at 0.5 kb and " 3^^+ =

0.1 M. In constructing these diagrams, only those species which are derived from experiments at pertinent temperatures are considered here.

The equilibrium constants of reactions involving the formation of these 45

Table 4. Equilibrium constants of dissolution of M0O2 (logK)

Reaction 300Oc 350°C 400^0 450OC Réf.

M0O2+ 1/2 O2+ HgO 8.48 7.59 6.86 6.28 - HgMoO^o

M0O2+ 1/2 O2+ H2O 7.38 4.82 2.93 0.51 a - H+ 7.87±0.31 "b

M0O2+ 1/2 O2+ H2O+ Na+ - NaHMoO^® + H+ 6.52+0.20

M0O2+ 1/2 O2+ HgOf 2Na+ - Na^MoO^o + 2H+ •0.55+0.15 -2.54+0.23

M0O2+ 1/4 O2+ 2H++ 2C1" - MoO(OH)Cl2° + I/2H2O 16.02+0.47

M0O2+ 1/4 O2+ H++ Cl" + I/2H2O - Mo0(0H)2C1O 10.60+0.15 10.89+0.14 11.32+0.13

^From Kudrin (1986).

^This study. 46

species are summarized in Table 4.

The solid lines in Figure 9 represent the boundary lines between an

ion pair (either NaHMoO^° or Na2MoO^°) and other molybdenum species, while

the dashed lines are boundaries between HMoO^" and other Mo species.

Generally, Mo(6+) is stable at high fQ2 and high pH conditions, and Mo(5+)

occurs at low pH and low f.Q2' 300 °C or below, the dominant Mo(6+)

species is HMoO^", and the dominant Mo(5+) is MoO(OH)Cl2° under

geologically reasonable fQ2 and pH conditions. With increasing

temperature, ion pairs of NaHMoO^® or Na2Mo02° become more stable with

respect to HMoO^", which predominates only when 8^^+ is low (<0.1 M near

neutral pH). At the same time, the stability fields of H^MoO^o and

MOO(OH)2C1° both expand at the expense of ion pairs or HMoO^".

Extrapolation of the experimental results to higher temperature seems to

suggest that H2MoO^° and MoO(OH)2Cl° are the dominant species for Mo(6+)

and Mo(5+) respectively near magmatic conditions-(Figure 10). This is in accordance with the conclusion drawn by Candela and Holland (1984) that

H2MOO^° predominates in magmatic fluids.

The relative stability of ion pairs versus HMoO^* or H2MoO^° as a function of loga^^^ and pH is illustrated in Figure 11, where the solid lines are boundaries between an ion pair- and a HMoO^"-predominant field, and the dashed lines between an ion pair- and a H2MoO^°- predominant field.

It can be seen that the stability field of H2MoO^° (dashed lines) increases very rapidly with temperature, while that of HMoO^" (solid lines) decreases gradually with temperature. The curve with hatch marks confines a field above which neither H^MoO^o nor HMoO^" is possible at 350 °C. In view of the high salinity of ore fluids associated with most molybdenum deposits 47

a T=300®C b 1=350*0 10 10 r HgMoO \ M0O4*" 0 0 - \ \ S MaHMoO| -10 -10 ' O \ HM0O4" G)

• 1 0 -20 \ M

• 1 0 -30 w MoO(OH)ClA MoO(OH)2C1 V

-40 -40 L 8 10

T=400°C d T=450°C

HgMoO? HgMoO^ MagMoO: NagMoO^

HMoO \ HM0O4

MoO(OH)2Cr \ \ MoO (OH) 2cry

8 10

Figure 9. Relative stability of aqueous molybdenum species as a function of log fQ2 and pH at 300, 350, 400, and 450 °C. Only those species which are determined from experiments are considered here. At 350 °C and above, solid lines represent boundary between an ion pair and other molybdenum species, and dashed lines between HMoO^" and other molybdenum species. P=0.5 Kb, ^Na+ " ^Cl- O.IM 48

T=600°C

10 1 1 ' X ' ' 1 " 1 '

\H2M004 Na2Mo04

or -

\ \ \ \ HM0O4 \ \ \ \ \ \ \ \ \ \ OJ -10 \ s \ s o \ ^\ u- \ s \ U) \ MoOfOHjgCl" \ 0 s -20 \ s \ \ \ \ V \ \ -30 \ \ >

—40 1 1 1 i 1 1 L 1 1 A 1 6 8 10 12 pH

Figure 10. Relative stability field of aqueous molybdenum species as a function of log fo2 and pH extrapolated to 600 °C. All other conditions and notations are the same as in Figure 9 49

.1 V \ I T T \ \ \ ion pair HMOO," \ \ \ \ A \ ion pair HgMoO? - \

\ V -1 + \ o \ \ \\\ 350*0 ^ \ \ \ \ \ \ \ \ \\\ \\W\ 07 0 —2 — \ \ \ \ \ \ \ \ 400®C -3 - \ \ \ \ 450^0 \ \ \ 350®C \ 400®C\ \450®C —4 ^ I 1 _l ULl I 8 10 12

pH

Figure 11, Activity-activity diagrams depicting the relative stability among aqueous Mo(6+) species as a function of log[Na+], pH, and temperature. Solid lines represent boundary between an ion P3.ir (NaHMoO^® or Na^MoO^^) and bimolybdate, and dashed lines between an ion pair and molybdic acid 50

(Hall et al., 1974; White et al., 1981; So et al., 1983a, 1983b; Shelton et

al., 1987), it is reasonable to believe that at this temperature, ion pairs

are the most Important form of molybdenum species in hydrothermal fluids at

relatively high oxygen fugacity.

The influence of temperature on the speciation of Mo above 450 °C was not determined in this study. As mentioned earlier, it seems that both

H2MOO^° and MoO(OH)2Cl are stable well above the upper temperature of major hydrothermal events, based on the extrapolation of my results. The

importance of H^MoO^o has been addressed by several authors (Candela and

Holland, 1984; Tingle and Fenn, 1982; Wood et al., 1987). However, the possibility of Mo(5+) as MoO(OH)2Cl° in hydrothermal fluids has not been reported by geochemists. The reason for this is complicated. It may be due to fundamental differences in experimental method. On the other hand, high temperature properties of the saline fluid itself may have contributed to the contradiction, in addition to the influence of temperature and oxygen fugacity. Works done by Sourirajan and Kennedy (1962) and Bodnar and Sterner (1987) show a heterogeneity of the fluid phase in che system

NaCl-H20 at high temperatures. At temperatures above 650 °C, at which most partitioning studies were carried out, there exists a wide unmixing region with two phases coexisting. One is an H^O-rich fluid and the other is

NaCl-rich liquid (Fig. 11). Variation of composition within this miscibility gap will yield no difference in fluid and liquid compositions but will change the relative volumes of the fluid and liquid phases. As a consequence, a partitioning study is virtually performed as if at constant composition. Any change in the partition coefficient is simply due to the changes in the proportion of H20-rich fluid and NaCl-rich liquid which 51

^5 MMBurvmtnU

_ Sourtrajan and Kennedy (1962) 1000 P: 1000 Ban

? 900 2 3 CD 2

^r 800

700

600 100 Wt. % NaCI

Figure 12. Phase relations in the system NaCl-H20 at high temperatures and 1.0 Kb. There exists a large unmixing region above 630 °C, within which two phases are coexisting, one rich in H2O, and the other rich in NaCI. Adopted from Bodnar and Sterner (1987) 52 remix together upon quench. Because many partitioning experiments were performed under conditions very close to the critical point of the system

NaCl-H20, the partitioning coefficient generated is thus meaningless because it is nothing more than a quantity that depends on the size of the system. The large uncertainty associated with partition coefficient indicates that this is an inadequate way to infer information about speciation. 53-54

CONCLUSIONS

Solubility of M0O2 in NaCl solutions represents an ideal condition

under which the behavior of molybdenum can be deduced. This investigation

demonstrates that a significant amount of molybdenum can be carried in

hydrothermal fluids in an environment that is lacking or poor in reduced

sulfur. The apparent NaCl concentration dependence of M0O2 solubility,

particularly at high temperatures, implies that a relatively high salinity

fluid favors the transport of molybdenum. The influences of temperature,

oxygen fugacity, and pH on the solubility of M0O2 lie largely in their role

of affecting Mo speciation.

In magnetite-hematite buffered solutions, molybdenum species with

oxidation states of both 5+ and 6+ can coexist stably. With increasing

temperature, both Mo(5+) and Mo(6+) species undergo a stepwise complexing.

At lower temperatures (<300°C), the charged molybdenum species , HMoO^',

predominates. At high temperatures, however, ion pairs of NaHMoO^° and

Na2MoO^° become increasingly important, and HMoO^" is significant only when

a^^^ is low. In contrast to Mo(6+), hydroxyl-chloro Mo complexes are the

principal Mo(5+) species with MoO(OH)Gl2° stable at lower temperatures and

MOO(OH)2C1° stable at high temperatures.

This study, then, has provided some insights on the mode of Mo

transport in hydrothermal fluids. These are of petrological and economic

interest, particularly when the molybdenum is subsequently removed from solution and deposited as molybdenite (M0S2). Experiments and model calculations relevant to this second stage are the subject of a second paper in this series. 55

CHAPTER II.

SOLUBILITY OF MOLYBDENITE (M0S2) IN HYDROTHERMAL SOLUTIONS 56

ABSTRACT

The hydrothermal solubility of molybdenite was measured in NaCl

solutions at temperatures from 300 to 400 °C under water vapor saturated

pressures. All of the solubility experiments were performed by

precipitating molybdenite (M0S2) from solutions buffered either by Fe^O^-

Fe203-FeS2 (HHP) or Fe30^-FeS2-FeS (MPP), to which M0O3 was added as the

source of molybdenum. NaCl concentrations were in the range from 0.1 to

4,0 M.

The total concentration of molybdenum in solution varies from below

0.1 ppm to over 10 ppm, increasing with the NaCl concentration in run fluids. These values are 2 to 3 orders of magnitude lower than measured under sulfur-free conditions. At all temperatures within the experimental range, molybdenite is more soluble in MHP-buffered solutions than it is in

MPP-buffered solutions, suggesting an effect of fQ2 and fg2 on the solubility. A retrograde behavior of molybdenite solubility with respect to temperature from 350 to 400 °C was observed under both buffer conditions.

Solubility of molybdenite under geologically relevant conditions was also modelled using the speciation scheme devised for the sulfur-free system, and is compared with the measured values. The results indicate that Mo(6+), present as Na2MoO^°, is more important at high temperatures.

At lower temperatures, MoO(OH)2Cl° is responsible for the transport of molybdenum in hydrothermal solutions. 57

INTRODUCTION

Molybdenite is the most important Mo-bearing phase that occurs

naturally. It is found in a variety of geological environments as

disseminations in porphyries, skarns, W-Mo quartz veins, and even in

uranium sandstones. The formation temperature of the relevant ore

deposits, as deduced from fluid inclusion studies, spreads over almost the

entire range of hydrothermal activity, from over 600 °C in porphyry

molybdenum deposits (Hall et al., 1974; Bloom, 1981; White et al., 1981) to

less than 200 °C in U-Mo associations (Tugarinov et al., 1973). This

feature of molybdenite deposits raises a fundamental question about the

mobility of molybdenum in ore-forming fluids under various geochemical conditions. Solution of this problem, however, requires knowledge of the solubility of molybdenite as a function of temperature, oxygen and sulfur fugacities, and fluid composition. This information, coupled with molybdenum speciation, will allow a comprehensive understanding of the mass transport of molybdenum and ore deposition.

Despite the importance of this information, surprisingly few studies of molybdenite solubility have been attempted. The only data available from a systematic study are provided by Wood et al. (1987), who studied the solubility of molybdenite along with nine other minerals in the system

NaCl-C02-H20 up to 350 °C. They found that molybdenum content in their solutions rarely exceeded 1 ppm, indicating a sparingly soluble nature of molybdenite. Westrich (1974), on the other hand, reported an increased solubility of molybdenite in KCl-HCl solutions in their preliminary study.

Aside from these works, there is a complete lack of experimental 58

constraints on the role of oxygen and sulfur fugacity on the molybdenite

solubility. Obviously, these two factors are critical to the transport of

molybdenum because molybdenite is stable only under conditions of

relatively low oxygen fugacity and high sulfur fugacity (Hsu, 1977).

Calculations by both Kudrin (1986) and Smith (1983) also demonstrate that

molybdenite solubility is a function of reduced sulfur concentration in

addition to its strong temperature dependence. These deductions are

further supported by Walder et al. (1988) who conducted fluid inclusion

studies to investigate the molybdenum and sulfur contents of the fluids

associated with molybdenum deposits in New Mexico. They observed a negative correlation between molybdenum and sulfur, particularly reduced sulfur, in the inclusion fluids. This phenomenon suggests that thiomolybdate, contrary to the proposal of Arutyunyan (1966), is unimportant to the transport of molybdenum in hydrothermal solutions.

To understand the origin of molybdenum deposits better, we need to know the behavior of molybdenum during hydrothermal processes, from the onset of hydrothermal activity to the end of ore formation. This requires a complete knowledge of speciation and mass transport of molybdenum under geologically relevant conditions. Since the solubility of molybdenite in the sulfur-rich system is apparently very low, there must be a significant difference in conditions under which molybdenum is transported and under which molybdenum is precipitated. In the first paper of this series (Cao,

1989; hereafter "Paper I"), the transport mechanism of molybdenum in the sulfur-free system was discussed, representing the behavior of molybdenum in an idealized pre-ore stage. It is important, however, to know not only how molybdenum is transported but how molybdenum is deposited also. In 59

T=250 "C

—30 Hemat i te

-34

Pyr i t -42

—46 Pyrrhot i te

-50 -20 -25 22 19 -16 -13 -10 -7 Log fS2

Figure 1. Phase relations in the system Fe-O-S, illustrating the buffering capacity of the assemblages Fe304-Fe203-FeS2 (HHP) and FeoÛA- FeS2-FeS (MPP). MHP buffer has higher values of f^o and foo than MPP buffer does 60

this regard, this second paper is devoted to characterizing the deposition

mechanism of molybdenite, and to generating information about the transport

of molybdenum in sulfur-rich systems as well.

The measurements in this study were performed by precipitating

molybdenite (M0S2) from solution. This method has shown to have a great

advantage over direct solubility measurements of molybdenite in terms of

time required to attain equilibration. The great stability of molybdenite

makes its rate of dissolution reaction so small that no detectlble

concentration (<0.02 ppm) of molybdenum was found in run solutions even

after one month of standing at 300 °C in our reconnaissance runs. When we

replaced molybdenite with molybdenum oxide, however, equilibration was

generally achieved within two weeks at 350 °C. Molybdenite could then be

precipitated from the Mo-rich fluid by maintaining a known value of fg2.

high enough to exceed the probable K^p for molybdenite. Fugacities of

oxygen and sulfur in the experiments were buffered by the mineral

assemblages magnetite-hematite-pyrite and magnetite-pyrite-pyrrhotite

respectively. Their relative positions in fQ2 and fg2 space are shown in

Figure 1. NaCl was also added to all experimental solutions in order to simulate natural hydrothermal fluids, in which NaCl is the dominant constituent (Roedder, 1971; Skinner and Barton, 1973; Hall et al., 1974;

White et al., 1981). 61

EXPERIMENTAL PROCEDURES

The experimental methods in this study were similar to those in the

M0O2 solubility study, described in Paper I. Because the same vessel was

used for two different buffer assemblages, caution was exercised to avoid

performing experiments using the two buffers alternately. Instead, the

experiments were arranged so that runs involving magnetite-pyrite-

pyrrhotite started only after all the experiments involving magnetite- hematite-pyrite had been completed. In this way, the buffering effect was guaranteed.

All the solids, except pyrite, were synthetic minerals which were purchased and used without further purification. Synthetic magnetite, hematite, M0O3, M0P2, and molybdenite were purchased from Alpha Company.

Ferrous sulfide (FeS) and NaCl were supplied by Fisher, ACS certified grade. The pyrite (FeS2) was natural crystalline material from Elba,

Italy. Doubly distilled water was used in the preparation of synthetic fluids. In operation, 15 to 25 grams of each mineral of a buffer assemblage and 3 to 6 grams of MoOg or M0O2 were loaded into a Barnes type vessel (Barnes, 1963, 1971) which was then sealed and placed vertically into a two-zone furnace. The vessel was then connected to a vacuum pump to evacuate the air inside, and the vacuum was used to draw run solution into the vessel through a sampling tube. Afterwards, the vessel was ready for heating.

Run duration at a selected temperature was between 10 to 20 days for most of the experiments, although longer or shorter run times were also used. During a run, the furnace was tipped at least twice a day to 62

accelerate equilibration, which was approached from both the high and low

temperature side of run conditions. Run solution was extracted from the

reaction vessel at the run temperature and pressure through a sampling tube

which was rinsed with doubly distilled water and then with run fluid prior

to sampling. For each run, two samples of 5 and 10 ml aliquots were

withdrawn. One was for pH measurement. The other was Immediately filtered

and diluted with a different solution depending on the purpose of analysis.

For molybdenum analysis, the sample was diluted with dilute HgCl2 solution;

for Fe and Na, the sample was diluted with dilute HNO3; and for sulfide

(S^'), the sample was diluted with an antioxidation buffer solution made by

dissolving 70 grams of ascorbic acid and lOOgrams of NaOH in 1 liter of

distilled water.

Molybdenum was analyzed with a Beckman electron photometer using the

dithiol method of Clark and Axley (1955). Iron and sodium were analysed on

a Perkin-Elmer model 303 atomic absorption spectromter using an acetylene-

air flame. Sulfide was analysed using a Cole-Parmer sulfide electrode.

Solid phases at the end of each run were examined with both powder X-ray diffraction and optical microscopy. Both methods indicated that all the minerals of a buffer assemblage (either magnetite-hematite-pyrite or magnetite-pyrite-pyrrhotite) remained. Molybdenite was identified under the microscope and also from its XRD pattern when initial M0O3 was abundant. 63

EXPERIMENTAL RESULTS

Most of the solubility measurements were performed at 350 °C and 400

°C in NaCl solutions ranging from 0.1 to 4.0 m. The results are summarized

in Table 1. Also listed is total sulfur concentration analyzed as sulfide

(S^") for some of the runs. Because of the apparent loss of sulfur during

sampling (as evidenced by the evolution of H2S gas), the measured sulfur

content is not a true total sulfur concentration in run solutions, and

cannot be used in calculations of mass balance and pH, even though it is

very important in assessing the speciation of molybdenum under these

conditions. Nevertheless, these data are still useful because they are

probably proportional to the true values and provide a relative

concentration in each run for rough comparison.

As can be seen from Table 1, the solubility of molybdenite in both

magnetite-hematite-pyrite (MHP) and magnetite-pyrite-pyrrhotite (MPP)

buffered solutions is low compared to values reported in Westrich's (1974)

abstract describing preliminary studies (500-2000 ppm Mo). Moreover, in

MPP-buffered solutions the solubility is even lower than that obtained by

Wood et al. (1987) at the same temperatures. Figures 2 and 3 illustrate the relationship between molybdenite solubility and fluid composition. It is obvious that the solubility in both MHP- and MPP-buffered solutions increases with NaCl concentration, but decreases with increasing temperature from 350 to 400 °C. This retrograde behavior of solubility of molybdenite with respect to temperature under experimental conditions seems directly related to the total sulfur content in solutions. As shown in

Figure 4, there is a negative correlation between molybdenum and sulfur 64

Table 1. Experimental solubility data for molybdenite (M0S2) in NaCl solutions. Concentrations are in molarity unless specified

Run # T°C Duration m^aci Quench m^^ S (ppm) (days) pH

MHP buffer

31-3 300 6 1.0 5.5 2.60*10-4 0.1

32-1 350 9 0.1 4.23 5.25*10-5 28.0 32-2 350 13 0.1 3.70 4.17*10-5 19.8 34-1 350 9 0.32 4.20 7.24*10-5 6.0 31-1 350 10 1.0 5.20 1.45*10-4 0.1 31-2 350 19 1.0 5.15 1.55*10-4 0.9 35-1 350 11 2.5 4.20 1.99*10-4 15.0

23-1 400 11 0.1 3.43 1.05*10-5 13.6 23-2 400 17 0.1 3.90 6.31*10-6 19-1 400 7 1.0 4.65 3.55*10-5 19-2 400 14 1.0 4.90 1.58*10-5 20-1 400 7 2.5 5.33 8.32*10-5 20-2 400 13 2.5 4.90 3.98*10-5 25-2 400 13 4.0 4.60 6.92*10-5 4.6

MPP buffer

30-2 350 11 0.1 5.15 2.00*10-6 29,7 36-1 350 19 0.5 4.50 3.16*10-6 25.5 27-1 350 17 1.0 5.40 1.10*10-5 10,4 27-4 350 21 1.0 5.92 8.90*10-6 22.7 33-1 350 8 2.5 3,60 6.31*10-6 16.9 33-3 350 12 2.5 4,20 5.01*10-6 29-2 350 12 4.0 4,90 1.51*10-5 4.3

30-1 400 14 0.1 5.55 3.63*10-7 67.5 30-3 400 10 0.1 5.00 5.50*10-7 90.8 36-2 400 8 0.5 3.92 1.26*10-6 27-2 400 8 1.0 5.30 5.01*10-6 24.8 27-3 400 17 1.0 5.25 3.16*10-6 29.0 33-2 400 8 2.5 4.23 2.00*10-6 18.1 29-1 400 10 4.0 4.05 4.27*10-6 43.5 65

Mt.—Hem—Py —3

-4 —

0

vy O

o? 0 -5

-6 -2 -1 0 Log CNaCl (tot)]

Figure 2. Plot of logMo(tot) vs. logNaCl(tot) demonstrating dependence of molybdenite solubility on NaCl concentration and temperature in MHP-buffered solutions 66

Mt—Py—Po

+) 0 -p w 0

m 0

—1 0

LogCNaCl (tot)]

Figure 3. Solubility of molybdenite as a function of NaCl concentration and temperature in MPP-buffered solutions. In both Figures 2 and 3, the solubility increases with decreasing temperature under the experimental conditions 67

-3

-4 —

0 E -5 0 z 01 0

-6 —

-7 20 40 60 80 100 S(ppm)

Figure 4. Plot of logMo vs. total sulfur concentration in some run solutions buffered either by MHP or by MPP. There is negative correlation between Mo and S concentrations in the solutions 68

contents in the fluids. Although the data points are quite scattered, the

general trend indicates that the higher the sulfur content in solution, the

lower the molybdenite solubility is. The other possibility may be due to

the pH difference at run temperature, and will be discussed later. This

feature once again demonstrates the ineffectiveness of thiomolybdate

complexes proposed by Arutyunyan (1966) for transporting molybdenum under

geologically relevant conditions.

The influence of oxygen and sulfur fugacities on molybdenite

solubility is depicted in Figure 5, where a direct comparison of solubility

in the sulfur-free system (buffered by magnetite-hematite (MH)) and the

sulfur-rich system (MHP, MPP) is given. In this diagram, the solid lines

represent solubility curves at 400 °C and the dashed lines at 350 °C. Both

solid and dashed lines are obtained by least-squares fitting to the

experimental data. It can be seen that the solubility of molybdenite

decreases by 1 to 3 orders of magnitude from MH-buffered solutions to MHP-

buffered solutions simply due to the introduction of pyrite. This result

strongly suggests that sulfur, particularly reduced sulfur, is the key

factor causing the precipitation of molybdenite. At all temperatures in the experimental range, however, molybdenite is more soluble in MHP- buffered solutions than it is in MPP-buffered solutions, in spite of the higher sulfur fugacity imposed by the MHP buffer assemblage. This is probably due to the higher oxygen fugacity and consequently lower H2S activity associated with the MHP buffer. It should be noted that the unusually high molybdenum concentration in MHP-buffered solutions at 350 °C may partly be the result of sluggishness of equilibration among the buffer minerals, particularly pyrite. The facts that the sulfur content is 69

T=400 °C

-1

-2 -

-3 — +> 0 +> —4

D7 O -5

-G

-7 -2 -1 0

Log CNaCl (tot)1

Figure 5. Comparison of molybdenite solubility in sulfur-free and in sulfur-bearing systems. HM denotes the buffer assemblage hematite-magnetite. Solid lines are solubility curves at 400 °C, and dashed lines at 350 °C. At constant temperature, the solubility in three different systems decreases in the following order: HM > MHP > MPP 70 unreasonably low in those solutions, and that the measured solubilities are about 10 times higher than the predicted ones are consistent with this interpretation. 71

GEOLOGICAL APPLICATIONS

As discussed earlier, most molybdenite deposits are formed at

temperatures between 300 and 450 °C, although some higher or lower

formation temperatures have also been documented (e.g., Hall et al., 1974;

White et al,, 1981; Bloom, 1981; Shelton et al., 1986; Lang and Eastoe,

1988). The fugacities of oxygen and sulfur seem to be between values that

are expected in fluids in equilibrium with the buffer assemblage

magnetite-hematite-pyrite and those in equilibrium with magnetite-pyrite-

pyrrhotite. Therefore, the results of this investigation have a direct

application in assessing the origin of molybdenite deposits.

Transport of Molybdenum

All the information generated from studies in MH-, MHP-, and MPP-

buffered solutions reveals that molybdenite solubility is a strong function

of temperature, oxygen fugacity, HgS activity, NaCl concentration, and

apparently pH (although it was neither controlled nor defined in this

study). Because of the lack of information on the true total sulfur

concentration, which makes the calculation of pH at run conditions

uncertain, estimation of molybdenum speciation in the sulfur-rich system is

impossible. However, the similarity between molybdenum solubilities with

respect to NaCl concentration in both sulfur-free and sulfur-rich systems

(Fig. 5) suggests that the forms of molybdenum derived from sulfur-free system (in Paper I) are also applicable in sulfur-rich systems. This assumption allows us to model the behavior of Mo in the presence of sulfur 72

thermodynamically beyond the experimental limit.

It has been shown in the Paper I that molybdenum can exist stably in

both +5 and +6 valence states under geologically relevant conditions.

Mo(5+) is mainly present as MoO(OH)Cl2° and MoO(OH)2Cl°; but Mo(6+), on the

other hand, experiences a stepwise complexing with HMoO^', NaHMoO^®, and

Na2MoO^°, stable at 300, 350, and above 400 °C respectively. Given this

speciation scheme, a dissolution or precipitation of molybdenite can be

hypothesized:

MoSg +1/4 Og +3/2H2O +2H++2G1--- MoO(OH)Cl2°+2H2S (T-SGQOC) (1)

[Mo(5+)][H2S]2 Ki

M0S2+I/4 O2+5/2 H2O+H++CI" — MoO(OH)2G1°+2 H2S (T>300°C) (2)

[Mo(5+)][H2S]2 Ko

M0S2 + 1/2 O2 + 3 HgO — HMoO^- + H+ + 2 HgS (T-300°C) (3)

[Mo(6+)][H2S]2[H+] Kc

MoSg +1/2 O2 + 3H2O +Na+ . NaHMoO^o + H+ + 2H2S (T-SSOOC) (4)

[MO(6+)][H+][H2S]2 K/, f02^/2[Na+]

M0S2 +1/2 O2 +3H2O +2Na+ — Na2Mo04O + 2H++2H2S (T>400°C) (5)

[Mo(6+)][H+]2[H2S]2 K= f02^/2[Na+]2 73

where the brackets denote the molarity, and in dilute solution a^-fi].

The equilibrium constants (Kj) for the reactions above can be

calculated by combining thermodynamic data in Table 4 of Paper I and those

for M0S2, H2O, and ^2^(aq.)' example, the calculation of logK for

reaction (2) at 350°C is demonstrated here.

From Table 4 (Paper I), we have,

M0O2+I/4 Og + l/ZHgO +H+ +01" - MOO(OH)2C1°

logK - 10.6

AGj.o - - 2.303RTlogK

- -126.44 KJ/mol.K

AGj.O- ACOf MoO(0H)2Cl- q2-1/2 AG°f AG°f J^q02

- AG' - AG°J^q02

AG' - ACOf Mo0(0H)2Cl' AG°f Q2-1/2 AG°f ^20" AG°f ^1-

where AGj.° is the free energy of reaction, and AG°£ the free energy of formation of substance i. The standard state in all calculations is 298 K and 1 bar pressure of pure substance or 1 mole of hypothetical solution for aqueous species.

AG' - AGj.o + AG°f MQ02 - - 684.19 KJ/mole K

For reaction (2)

AGj.° - AG'+ 2 AG°f H2S " ^ AG°jj20 ' AG°f,MoS2

- -684.19 - 181.46 + 2(269.91) + 298.78 74

Table 2. Equilibrium constants (logK) for reactions involving the dissolution of molybdenite (M0S2)

Reaction SOQOC 350Oc 400°C 450°C

M0S2+ 1/2 O2+ 3H2O - HMoO^-f H++ 2H2S -2.36+0.31

M0S2+ 1/2 O2+ 3H2O+ Na+ - NaHMoO^o + H+ + 2H2S -1.81±0.20

M0S2+ 1/2 O2+ 3H2O+ 2Na+ - NagMoO^o + 2H+ + 2H2S -7.00±0.15 -7.27+0.23

M0S2+ 1/4 O2+ 3/2H2O+ 2H+ + 2C1- - MoO(OH)Cl2 5.80+0.47

MoS2+ 1/4 O2+ 5/2H2O+ H+ + Cl- - MoO(OH)2C1 2.27±0.15 4.44±0.14 6.59±0.13 75

- -27.07 KJ/mole.K

AGj.0 - -2.303RTlogK2

which leads to

logKg - 2.27

where the free energy of formation of H2O is calculated from Helgeson et

al. (1978), M0O2 and M0S2 from Smith (1983), and H2S from Naumov et al.

(1974). The equilibrium constants for reactions (1) to (5) at different

temperatures are summarized in Table 2.

Having, the equilibrium constant Ks, we are now in a position to

calculate the solubility of molybdenite as a function of temperature and

fluid composition. In buffered conditions, the fugacities of oxygen and

sulfur are defined by reactions

ZFegO^ + 1/2 Og — SFegOg

FeS2 + 4Fe202 SFegO^^ + S2

for MHP buffer, and

1/2 FegO^ +3/2 FeS2 — 3FeS + O2

FeS2 — FeS + 1/2 S2

for MPP buffer. The activity of H2S is thus determined by 76

Table 3. Selected equilibrium constants (logK)

Reaction 300°C 350°C 400Oc Ref.

- H+ + HS" -8.18 -8 ,81 a - H,S -1.50 -1,,44 _b H+ + -5.70 - 6,.42 _a 2 Or so. •f 2H-' 47.20 39,,80 _a Sari : •= S02 27.55 25,,25 23.35 _c I/2S2 + H2 - H2S 3.82 3,,60 3.40 _c I/2S2 + 3/2 0 so. 31.70 28,,65 26.14 __c H 2 + 1/2 ©2 H2O 19.78 17.,93 16.40 c

^Murray and Cubiciotti (1983). ^Drununond (1981). CRobie et al. (1978). 77

250 300 350 400 450 500 T *C

Figure 6. Calculated molybdenite solubility as a function of temperature with fo2 controlled by HM at 0.5 Kb. a^^* - a^i, - 0.1 M, and pH - 5.5. See text for calculation model 78

-1

-2

O

Z 07 0 _J -5

—6

-7 250 300 350 400 450 500 T °C

Figure 7. Calculated solubility of molybdenite as a function of temperature with fQ2 and fg2 controlled simultaneously either by MHP or by MPP buffer at 1.0 Kb. pH = 5.5, and a^^^, = a^^. =0.1 M. Note the retrograde behavior of the solubility in the lower temperature range 79

HgO + 1/2 S2(g) — HgScaq) + 1/2 Ogfg)

The equilibrium constants for some selected reactions used in the

calculation are summarized in Table 3.

Figure 6 shows the calculated solubility of molybdenite from 300 to

450 °C under conditions where pH-5.5, [CI"]-[Na"*"]-0.1 m, and fQ2 is buffered by magnetite-hematite. The solubility curve increases monotonically with temperature and more rapidly at temperatures above 375

°C. This is in contrast to the solubility under conditions where fg2 is buffered by mineral assemblages (Figure 7). In Figure 7, both solubility curves do not increase or decrease monotonically, but exhibit a retrograde behavior in the low temperature range. Moreover, the MHP curve lies above the MPP curve.

The calculated effect of pH on molybdenite solubility is illustrated in Figures 8, 9 and 10. Figures 8 and 9 are molybdenite solubility in solutions buffered by MHP and MPP buffer respectively, and figure 10 is the solubility under unbuffered conditions in terms of fg2, with 3^23- 0,01 m at all temperatures. Again, [Na*]-[C1"]-0.1 m and pressure of 0.75 Kb are assumed in constructing these diagrams. It can be seen that at constant temperature, each solubility curve shows a V-shaped pattern. The portion with negative slope of the solubility curve is contributed by Mo(5+), and the portion with positive slope by Mo(6+). Therefore, the solubility tends to increase with either increasing acidity or basicity of a solution, consistent with the experimental results of Westrich (1974). Although the minima of solubility curves in each diagram all lie near the neutral pH at relevant temperature, there is an apparent trend that the minima move 80

towards the acid region with decreasing temperature, suggesting the

increasing importance of Mo(6+) species in transporting Mo at lower

temperatures.

A remarkable distinction between buffered (Figures 8 and 9) and

unbuffered (Figure 10) solubility curves is that, in Figure 10, high

temperature curves lie above low temperature ones almost in the entire pH

range. In buffered systems (Figures 8 and 9), however, the solubility

curves of different temperature cross over each other, such that in some pH

ranges, the solubility of molybdenite increases or decreases with

temperature monotonically; while in some other pH ranges, a maxmum or

minimum occurs. Recognition of this feature is essential to understanding

the retrograde behavior of molybdenite solubility obtained from the

experiments, and from the calculation as well, as described earlier.

Therefore, the influence of pH is very complicated when fg2 buffer is

employed. Its role at constant temperature is very clear. However, if comparison of solubility at different temperatures is attempted, the pH of the run fluid at relevant temperature must be known. Otherwise, ambiguities may arise because at some pH values where two solubility curves meet, the influence of temperature is masked by pH, and may lead to an improper conclusion that.the solubility is independent of temperature. All of these calculations simply suggest that if the pH varies somewhat during the run, it may cause the molybdenite solubility to change in a complicated fashion that may explain the difference in calculated solubility curves in buffered and constant 8^23 conditions.

In Figures 8 and 9, the measured solubilities of molybdenite are also plotted for comparison with calculated values. Because the pH at run 81

MHP

400 °C 300 "C 350 °C

+) 0 +)

Ol 0

pH

Figure 8. Calculated solubility of molybdenite as a function of pH in MHP-buffered solutions at 300, 350, and 400 °C, and at 1.0 Kb. ®Na+ " &ci_ - 0.1 M. The measured solubility at relevant temperatures is also plotted for comparison. Triangle represents the solubility data at 350 °C and circle at 400 °C. Arrow indicates the direction that the measured solubility approaches the calculated value. Neutral pH for the run solutions is assumed 82

MPP

-1 400 "C

300 X -2

350 rs -3 +> O +) -4 V-/ o z -5 07 0 -6

-7

—8

-9 J L 4 5 6 7 8 9 10 11 pH

Figure 9. Calculated isothermal solubility of molybdenite as a function of pH in MPP-buffered solutions at 1.0 Kb, aj^^^ = a^^, « 0.1 M. The measured solubility at relevant temperatures is also shown for comparison. The little circle denotes the solubility data at 400 °C, triangle at 350 °C, and square at 300 °C which is from Wood et al. (1987). Neutral pH for run solutions is assumed except at 300 °C 83

50 X

Figure 10. Calculated molybdenite solubility as a function of pH under unbuffered conditions in terms of fg2 at 300, 350, and 400 °C, 1.0 Kb. Oxygen fugacity is controlled by HM, 3^%+ = aci- =0-1 M, and ay2s ~ 0.01 M. Note the monotonie increase in molybdenite solubility with respect to temperature. The shaded areas represent the limit of error on the solubility calculation 84

condition in this study was unknown, we assume a neutral or slightly lower

pH for the run fluids. This assumption is reasonable, given the facts that

neither strong acid nor strong base is present in the systems, and that

associated species predominate at high temperatures even in strong

electrolyte solutions. As can be seen, the calculated and the measured

solubilities agree surprisingly well, with a deviation within one order of

magnitude except at 350 °C in MPP-buffered solutions. This indicates that

the speciation scheme derived in Paper I and the dissolution or

precipitation mechanisms proposed earlier are probably valid.

The solubility of molybdenite as a function of fQ2 and pH is contoured

in Figure 11 at 350 °C where the stability fields of sulfur species are

illustrated. Total sulfur was chosen to be lO'^-S m because the concentration of reduced sulfur in most hydrothermal solutions is between

0.001 and 0.1 m (Ohmoto, 1972), and ajj^^.-aQj^.-0.1 m. A dashed line separating Mo(5+)- and Mo(6+)-predominant fields is also shown in Figure

11. To the left of the dashed line. Mo(5+) predominates, and to the right

Mo(6+) predominates. We can see that the solubility increases with Ïq2 at constant pH. However, in sulfide-predominant fields, the solubility is much less sensitive to the change in fQ2 than in the sulfate stable fields.

Lower solubility is found at lower logfQ2 and near neutral pH conditions at this particular temperature.

Based on the studies of fluid inclusion and alteration assemblages,

Crerar and Barnes (1976) determined the composition of an average ore fluid for porphyry Cu deposits which roughly fall into the range of pH=-3 to 6 and logfQ2" -28 to -30 at 350 °C. If this composition can be applied to molybdenum deposits because of the similarity between porphyry Cu and 85

T=350°C

-20

o(\J q- -30 - O? 0

Mo(V) \Mo(VI)

-40 8 10 12 pH

Figure 11. Logfo2-pH diagram illustrating dependence of molybdenite solubility on fo2 and pH in NaCl solution at 350 °C and vapor saturated pressure. Total sulfur concentration is chosen to be 10" • M, a^a^, - aQ2_ - 0.1 M. Contours show the molybdenum concentration in solutions. Dashed line separates Mo(5+)- and Mo(6+)- predominant field 86

porphyry Mo deposits, we will find that the solubility of molybdenite under

this condition is from 0.01 to 10 ppm (Figure 11). This concentration,

however, is too low to account for the observed amount of molybdenum in the

deposits. In order for a significant amount of molybdenum to be

transported in such solutions, the activity of H2S must be lowered. From

reaction (1) to (5), one can see that one log unit decrease of aj^2s will

result in an increase of molybdenite solubility by 2 orders of magnitude

with all other variables being fixed. Otherwise, deposition of molybdenite

will be the major event under this condition.

The discussion above reveals that molybdenite is virtually insoluble

in subcritical saline fluids in the presence of even 10*^of reduced

sulfur. Its solubility decreases very rapidly at temperatures below 300

°C. Thus, in the temperature range studied, a typical hydrothermal fluid

précipités molybdenite, and significant transport of molybdenum in such

solution is seemingly impossible, unless fQ2 is abnormally high and aj^2S

low, or in an acid or alkaline solution.

Ore Formation

If most of the molybdenum in a molybdenum deposit is derived from the

associated igneous rocks (Candela and Holland, 1985), then the removal of

molybdenum from the granitic magma will be accomplished by the exsolution

of fluid during crystallization of the melt. At least 100 ppm or more of

molybdenum per liter of water can be expected to fractionate into the magmatic fluid, according to Candela and Holland (1984). The molybdenum extracted in the fluid can exist stably as Na2MoO^°, or probably as H2MoO^° 87

4

0

-4 SO

80 U) 0

-12

-16

P=1.0 kb

20 300 400 500 600 700

T °C

Figure 12. Fugacities of major gas species in the system S-O-H at 1.0 Kb. Oxygen and sulfur fugacities are controlled by MHP buffer. At temperatures below 500 ®C, reduced S species are more abundant than oxidized sulfur species 88

as suggested by Candela and Holland (1984), and most possibly, as

MoO(OH)2C1°. At the same time, sulfur can also be partitioned into the

fluid in significant abundance (Burnhara, 1979; Holloway, 1981), and is

carried along with molybdenum to the top of the magma chamber. Because

oxidized sulfur species predominate at high temperatures, no deposition

will occur. For example, in the MHP buffered system, reduced sulfur

species will not exceed oxidized sulfur species until temperature is below

500 °C (Figure 12). Therefore, major deposition of molybdenite is likely

to occur at a temperature around 500 °C in a porphyry system. This is in

accordance with fluid inclusion studies (White et al., 1981; Bloom, 1981),

While a simple temperature drop and a consequent increase in 3^23 may

account for the deposition mechanism for some high temperature molybdenum

deposits (T>500°C), the situation for those deposits formed at moderate or low temperatures becomes rather complicated. Since most molybdenum carried in magmatic fluid will precipitate out above 450 °C, there would seem to be not enough molybdenum left to form lower temperature deposits. With decreasing temperature, the pH of the fluid will change from neutral or mildly basic to mildly acidic (Brimhall and Crerar, 1987), and molybdenum of higher oxidation state will be replaced by Mo(5+) (Figures 8, 9, and

10). Furthermore, reduced sulfur species will predominate at lower temperatures (Ohmoto and Rye, 1979), making the transport of molybdenum even less favorable as discussed earlier.

A possible mechanism for molybdenum transport under this condition is boiling. Fracturing of the cupola or country rocks leads to a release of pressure and results in boiling (Burnhara, 1979). The differential partitioning of H2S and H2 between liquid and vapor phases will deplete H2S 89

in the liquid phase and increase oxygen fugacity (Brimhall and Crerar,

1987). As a result, the solubility of molybdenite will be greatly

enhanced, and redissolution and remobilization of molybdenum can occur.

This mechanism allows molybdenum to be transported to the deposition site

where H2S is already enriched, forming vein type deposits. Evidence of

boiling is found in many molybdenum deposits (So et al., 1983a, 1983b;

Shelton et al., 1986, 1987), suggesting that this is the most dynamic way

to transport molybdenum at moderate temperatures. Extreme conditions of

low pH, low aH2g, and high fo2» however, are not impossible in geological

environment, and may contribute to the formation of lower temperature

deposits (T<300°C). As shown in Figure 8, more than 100 ppm molybdenum can

be carried in a solution buffered by magnetite-hematite-pyrite if the pH is

below 4.0 at 300 °C.

It should be noted that dilution or mixing with meteoric water at

temperatures below 350 °C may not necessarily be a cause of precipitation,

although this will bring a solution towards neutral pH. Referring to

Figures 8, 9, and 10, the minimum of each solubility curve all lies to the

acid side of its corresponding neutral pH. Therefore, increasing pH from

the mildly acid region will actually increase molybdenite solubility, particularly at temperatures below 300 °C. Furthermore, dilution will decrease the activity of H2S if it is not effectively buffered. As described earlier, one log unit decrease in ay2s would result in an increase of molybdenite solubility by two orders of magnitude. Stable isotope study has shown that mixing of ore-forming fluids with meteoric water is a common phenomenon in molybdenum deposits with a lower formation temperature (So et al., 1983a, 1983b; Shelton et ai., 1986, 1987). From 90

this point of view, the role of dilution or mixing in affecting molybdenum

transport should be reconsidered.

All the discussions above are based on a simplified system, i.e.,

NaCl-H20 fluid. In reality, however, the situation is much more complicated. Because F-bearing minerals and CO^-rich fluid inclusions are

often found in close association with some molybdenum deposits, arguments have been raised on the possibility of fluoro Mo species or carbonate complexing of molybdenum in the transport of molybdenum in hydrothermal solutions (Garten et al., 1981; MacDonald and Spooner, 1982). In addition, the importance of silicomolybdate complex has also been proposed by

Khitarov et al. (1965). Although there is no experimental evidence supporting F" or C02^" complexing of molybdenum at elevated temperatures and pressures, their role should not be underestimated. They may influence the solubility of molybdenite in one way or another. Further work needs to be done to clarify this point. 91

CONCLUSIONS

This investigation demonstrates that the solubility of molybdenite is

a function of temperature, oxygen fugacity, H2S activity, pH, and NaCl

concentration. Temperature tends to enhance the solubility under

conditions of unbuffered fg2; but ambiguity arises in the case that fg2 is

buffered in the temperature range 350 to 450 °C where reduced sulfur

species predominates. H2S has a profound influence on the transport of molybdenum. One log unit increase in H2S activity would decrease molybdenite solubility by 2 orders of magnitude. The solubility increases invariably with increasing fQ2 and NaCl concentration of the hydrothermal system. At any single temperature, deviation of pH from the solubility minimum would result in an increase of molybdenite solubility.

Under magmatic conditions, or in the early stage of hydrothermal events, oxidized sulfur species predominate. These conditions resemble the sulfur-free system; significant amounts of molybdenum (>10"%) can be transported in hydrothermal fluids. Molybdenite precipitation can take place simply due to the cooling of hydrothermal fluids at about 500 °C when reduced sulfur species become increasingly abundant, if there is sufficient sulfur present in the system. In an average ore solution, the solubility of molybdenite is so low that significant mass transport seems unlikely.

However, the solubility can be substantially enhanced if boiling takes place. Dilution or mixing of hydrothermal fluid with meteoric water may not neccessarily cause precipitation of molybdenum; instead, it may account for the transport of molybdenum at low temperatures (<300°C).

During hydrothermal processes, both Mo(5+) and Mo(6+) species play a 92 part In transporting molybdenum. When fQ2 is large and pH is relatively high, molybdenum species with higher oxidation state are more favorable, and Mo is likely to be transported as the ion pair Na2MoO^°, or as H2MoO^°.

At lower lower fQ2 and pH, however, Mo(5+) species are more important, and molybdenum is transported as MoO(OH)2Cl°. It is the difference in stability field of the two molybdenum species that allows transport of molybdenum in hydrothermal solutions of virtually all possible compositions. 93

GENERAL SUMMARY

This investigation shows that under geologically relevant conditions,

molybdenum can exist stably in two different valence states, Mo(5+) and

Mo(6+). Species involving Mo(5+) are of the forms of MoO(OH)Cl2 and

MoO(OH)2C1, which are stable below and above 300 °C respectively. Mo(5+), however, mainly forms bimolybdate and ion pairs, depending on temperature and fluid composition. At low temperatures (<300 °C), or when 3^^+ is low,

HMoO^" is the dominant species. With increasing temperature, ion pairs of

NaHMoO/^° and Ha^MoO^o become increasingly important with Na2MoO^° stable above 400 °C. Molybdic acid predominates only in near magmatic or very high fQ2 conditions. In a tipycal hydrothermal solutions. Mo is most likely transported as Mo(5+) at high temperatures. At low temperatures, however, Mo(6+) may be the most impotant species responsible for the transport of molybdenum.

The mobility of molybdenum in hydrothermal solutions is a strong function of temperature, fQ2, 3^23, pH, and NaCl concentration. Under sulfur-free or low ajj2g conditions, a significant amount of molybdenum can be transported in the solutions, increasing with temperature and NaCl concentration. In sulfur-rich systems, particularly when aj^2s high, the solubility of molybdenite decreases dramatically with respect to the sulfur-free conditions. High fQ2 tends to enhance the solubility, especially in the sulfate stable field. At constant temperature and fQ2, the solubility is lowest in near neutral to mildly acidic solutions.

Changes in pH either way from the minimum will significantly increase the solubility of molybdenite. Precipitation of molybdenite can take place due 94 to the decrease in temperature and oxygen fugacity, or increase in the activity of Shift of pH from mildly alkaline to mildly acidic solutions will also caused precipitation. 95

REFERENCES

Armands, G. 1972. Geochemical studies of Uranium, Molybdenum and Vanadium in the Swedish Alum shale. Stockholm Gontr. Geol. 27:1-148.

Arutyunyan, L. A. 1966. Stability of water-soluble forms of molybdenum in sulfur-bearing solutions at high temperatures. Geokhymiya 4:479-483.

Averill, B. A., H. C. Sllvls, R. H. Tlekelmann, and W. H. Orrae-Johnson. 1980, Chemical physical studies of polynuclear molybdenum-iron- sulfur and mononuclear molybdenum (III) complexes, p. 217-228. In W. E. Newton and S. Otsuka (eds) Molyb. Chem. Biol. Significance. Plenum press, New York.

Avesten, J., E. W, Anecker, and J. S. Johnson. 1964. Hydrolysis of molybdenum (VI). Ultracentrifugatlon, acidity measurement, and Raman spectrum of polymolybdate. Inorg. Chem. 3:735-746.

Barnes, H. L. 1963. Ore solution chemistry. I. Experimental determination of mineral solubilities. Econ. Geol. 58:1054-1060.

Barnes, H. L. 1971. Investigations in hydrothermal sulfide systems, p. 317-335. in G. G. Ulmer (ed.) Research techniques for high pressure and high temperature. Springer-Verlag, New york.

Bloom, M. S. 1981. Chemistry of inclusion fluids in stockwork molybdenum deposits. Econ. Geol. 76:1906-1920.

Boctor, N. Z., R. K. Popp and J. D. Frantz. 1980. Mineral-solution equilibrium - IV. Solubilities and the thermodynamic properties of FeCl2° in the system Fe203-H2-H20-HCl. Geochim. Cosmochim. Acta

Bodnar, R. J. and S. M. Sterner. 1987. Synthetic fluid inclusions, p. 423-457. In G. C. Ulmer and H.L. Barnes (eds.) Hydrothermal experimental techniques. John Wiley & Sons, New york.

Brimhall, G. H. and D. A. Crerar. 1987. Ore fluids: magmatic to supergene. p. 235-321. In I. S. E. Carmichael and H.P. Eugster. (eds.) Thermodynamic modeling of geological materials: minerals, fluids and melts. Rev. Mineral, v. 17. Mineral. Soc. Am., Washington, D.C.

Burnham, C. W. 1979. Magmas and hydrothermal fluids, p. 71-136. In H. L. Barnes (ed.) Geochemistry of hydrothermal ore deposits. 2nd ed. Wiley-Interscience, New York.

Candela, P. A. and H. D. Holland. 1984. The partitioning of copper and molybdenum between silicate melts and aqueous fluids. Geochim. 96

Cosmochim. Acta 48:373-380.

Garten, R. B., J. R. Shannon, A. D. Ward, and E. P. Geraghty. 1981. Controls on molybdenite deposition in Henderson Mine, Empire, Colorado. Geol. Soc. Amer. Abstracts with Prog. 13:423.

Candela, P. A. and H. D. Holland. 1985. A mass transfer model for Cu and Mo in magmatic hydrothermal systems: The origin of porphyry-type ore deposits. Econ. Geol. 8:1-19.

Chou, I-M and H. P. Eugster. 1977. Solubility of magnetite in supercritical chloride solutions. Am. J. Sci. 277:1296-1314.

Clark, L. J. and J. H. Axley. 1955. Molybdenum determination in soils and rocks with dithiol. Anal. Chem. 27:2000-2003.

Coveney, R. M., Jr. and S. P. Martin. 1983. Molybdenum and other heavy metals of the Mecca Quarry and Logan Quarry shales. Econ. Geol. 78:132-149.

Crerar, D. A. and H. L. Barnes. 1976. Ore solution chemistry - V. Solubility of chalcopyrite and chalcocite assemblages in hydrothermal solutions at 200-350°C. Econ. Geol. 71:722-794.

Crerar, D. A., N. J. Susak, M. fiorcslk, and S. Schwartz. 1978. Solubility of buffer assemblage pryrlte + pyrrhotite + magnetite in NaCl solutions from 200 to 350 °C. Geochim. Cosmochim. Acta 42:1427-1437.

Drummond, S. E. 1981. Boiling and mixing of hydrothermal fluids: chemical effects on mineral precipitation. Ph.D. Dissertation. Pennsylvania State University, 396p.

Eugster, H. P. 1985. Granites and hydrothermal ore deposits: A geochemlcal framework. Mineral. Mag. 49:7-23.

Eugster, H. P., 1986, Minerals in hot water. Am. Mineral. 71:655-673.

Finch, W. I. 1967. Geology of uranium deposits in sandstone in the United States. U.S. Geol. Surv. Prof. Paper 538.

Frantz, J. D. and W. L. Marshall. 1984. Electrical conductances and ionization constants of salts, acids, and bases in supercritical aqueous fluids: I. Hydrochloric acid from 100° to 700°C and at pressures to 4000 bars. Am. J. Sci. 284:651-667.

Frantz, J. D. and P. K. Popp. 1979. Mineral-solution equilibrium - I. An experimental study of complexlng and thermodynamic properties of aqueous MgCl2 in the system Mg0-Si02-H20-HCl. Geochim. Cosmochim. Acta 43:1223-1239. 97

Griffith, W. P. and T. D. Wickins. 1967. Raman studies on species in aqueous solutions. Part II. Oxy-species of metals of groups VIA, VA, and IVA. J. Chem. Soc. A. 675-679.

Griffith, W. P. and P. J. B. Lesniak. 1969. Raman studies on species in aqueous solutions. Part III. vanadates, molybdates, and tungstates. Chem. Soc. London Jour. A: 1066-1071.

Gulbrandsen, R. A. 1966. Chemical composition of phosphorites of the Phosphoria Formation. Geochim. Cosmochim. Acta 30:769-778.

Haight, G. R. and D. R. Boston. 1973. Molybdenum species in aqueous solutions, a brief summary, p. 48-51. la P. C. H. Mitchell (ed.) Int. Conf. on the Chemistry and Uses of molybdenum. Climax Molybdenum Co. Ltd., London.

Hall, W. E., I. Friedman, and J. T. Nash. 1974. Fluid inclusion and light stable isotope studies of the Climax molybdenum deposits, Colorado. Econ. Geol. 69:884-901.

Helgeson, H. C. 1969. Thermodynamics of hydrothermal systems at elevated temperatures and pressures. Am. J. Sci. 267:729-804.

Helgeson, H. C., J. M. Delany, H. W. Nesbitt, and D. K. Bird. 1978. Summary and critique of the thermodynamic properties of rock- forming minerals. Am. J. Sci. 278A:l-229.

Holland, H. D. 1972. Granites, solutions, and base metal deposits. Econ. Geol. 67:281-301.

Holloway, J. R. 1981. Volatile interactions in magmas, p. 273-293. In R. C. Newton, A. Navrosky, and B. J. Wood (eds.) Adv. Physical Geochemistry. Vol. 1. Thermodynamics of minerals and melts. Springer-Verlag, New York.

Hsu, L. C. 1977. Effects of oxygen and sulfur fugacities on the scheelite-tungstenite and powellite-molybdenite stability relations. Econ. Geol. 72:664-670.

Hudson, T., J. G. Smith, and R. L. Elliott. 1979. Petrology, composition, and age of intrusive rocks associated with Quartz Hill molybdenum deposits, Southeastern Alaska. Canadian J. Earth Sci. 16:1805-1822.

Hudson, T., J. G. Arth, and K. G. Muth. 1981. Geochemistry of intrusive rocks associated with molybdenite deposits, Ketchikan Quadrangle, Southeastern Alaska. Econ. Geol. 76:1225-1232.

Ivanova, G. F., N. I. Lavkina, L. A. Nesterova, A. P. Zhudikova, and I. L. Khodakovskiy. 1975. Equilibrium in the M0O3-H2 system at 25-300OC. Geochem. Int. 12:163-174. 98

Ivanova, G. F., Z. M. Motorina, and V. B. Naumov. 1978. Formation features of the mineral associations of the Yugodzyr' molybdenum- tungsten deposit (Mongolia). Int. Geol. Review 20:855-863.

Jackson, K. J. and H. C. Helgeson. 1985. Chemical and thermodynamic constraints on the hydrothermal transport and deposition of tin: I. Calculation of the solubility of casslterlte at high pressures and temperatures. Geochim. Cosmochlm. Acta 49:1-22.

Kamilll, R. J. 1978. The genesis of stockwork molybdenite deposits: implications from fluid inclusion studies at the Henderson Mine. Geol. Soc. Am. Abst. with Programs 10(7):431.

Karyakln, Y. V. and E. N. Kryachko. 1967. Fluoro-complexes of Mo(VI) in dilute aqueous solutions. Russ. J. Inorg. Chem. 12:1355-1257.

Khltarov, N. I., L. A. Arutyunyan, and E. B. Lebedev. 1967. Experimental study of separation of molybdenum from granite melts at water vapor pressures up to 3000 atmospheres. Geochem. Int. 8:730-737.

Khltarov, N. I., L. A. Arutyunyan, and B. N. Ryzhenko. 1965. The effect of hydrogen sulfide on the migration of molybdenum In the form of a slllcomolybdate complex under elevated temperarures. Geochem. Int. 3:192-195.

Killeffer, D. H. and A. Llnz. 1952. Molybdenum compounds: their chemistry and technology. Interscience Publisher, New York.

Kudrln, A, V. 1986. The solubility of tugarlnovlte M0O2 in aqueous solutions at elevated temperatures. Geochem. Int. 23:126-138.

Lander, H. N. 1977. Energy related uses of molybdenum, p. 773-806. In W. R. Chappell and K. K. Petersen (eds.) Molybdenum in the environment. Vol. 2. M. Dekker, New York.

Lang, J. R. and C. R. Eastoe. 1988. Relationship between a Porphyry Cu-Mo deposit, base and precious metal veins, and Laramide Intrusions, Mineral Park, Arizona. Econ. Geol. 83:551-567.

MacDonald, A. J. and E. T. C. Spooner. 1981. Hydrothermal fluid flow and M0S2 transport in Boss Mountain brecela pipe, B.C.. Geol. Assoc. Canada Ann. Meeting 6:A90.

MacDonald, A. J. and E. T. C. Spooner. 1982. Carbonate complexing of molybdenum in hydrothermal solutions: Evidence from Boss Mountain, British Columbia. Geol. Soc. Amer. Abstrcts with Prog. 14:552.

Manning, D. A. C. and P. Henderson. 1984. The behavior of tungsten in granitic melt-vapor system. Contr. Mineral. Petrol. 86:286-293. 99

Marshall, W. L. and E. U. Franck. 1981. Ion product of water substance, 0-1000°C, 1-10,000 bars: New International formulation and its background. J. Phys. Chem. Réf. Data 10:295-304.

Mitchell, P. C. H. 1973. The chemistry and uses of molybdenum: Introductory lecture, p. 1-5. Jn P. C. H. Mitchell (ed.) Int. Conf. on the Chemistry and Uses of Molybdenum. Climax Molybdenum Co. Ltd., London.

Murray, R. C. Jr. and D. Cubiciotti. 1983. Thermodynamics of aqueous sulfur species to 300°C and potential-pH diagrams. J. Electro. Chem. Soc. 130:866-869.

Nabivanets, B. I. 1969. Acid and basic dissociation constants of hydrated molybdenum trioxide. Zhur. Neorg. Khim. 14:653.

Naumov, G. B., B. N. Ryzhenko, and I. L. Khodakovsky. 1974. Handbook of thermodynamic data. Âtomizdat, Moscow.

Ohmoto, H. 1972. Systematics of sulfur and carbon isotopes in hydrothermal ore deposits. Econ. Geol. 67:551-578.

Ohmoto, H. and R. 0. Rye. 1979. Isotopes of sulfur and carbon, p. 500-567. In H. L. Barnes (ed.) Geochemistry of hydrothermal ore deposits. 2nd ed. Wiley-Interscience, New York.

Pearson, D., C. S. Copeland, and S. W. Benson. 1963. The electrical conductance of aqueous sodium chloride in the range 300 to 383°C. J. Am. Chem. Soc. 85:1044-1047.

Podshivalova, A. K., A. S. Chernyak, and I. K. Karpov. 1984. Some transformations between molybdenum (VI) species in acidic solutions. Russ. Jour. Inorg. Chem. 29:1458-1461.

Pokalov, V. T. 1977. Deposits of molybdenum, p. 125-179. In V. I. Smirnov (ed.) Ore deposits of the USSR. V. 3. Pittman, London.

Quist, A. S. and W. L. Marshall. 1968. Electrical conductance of aqueous sodium chloride solutions from 0 to 800°C and at pressures to 4000 bars. J. Phys. Chem. 72:684-703.

Robie, R. A., B. S. Hemingway, and J. R. Fisher. 1978. Thermodynamic properties of minerals and related substances at 298.15K and 1 bar (10^ Pascals) pressure and at high temperatures. U.S. Geol. Surv. Bull. No. 1452.

Roedder, E. 1971. Fluid inclusion studies on the porphyry-type ore deposits at Bingham, Utah, Butte, Montana, and Climax, Colorado. Econ. Geol. 66:98-120. 100

Ross, R. B. 1972. Metallic materials. Specification Handbook, E & F.N. SPON LTD., London.

Ruzlcka, V. and R. T. Bell. 1984. Sandstone uranium, p. 36. In 0. R. EcKstrand (ed.) Canadian mineral deposit type. Geol. Surv. Canada Econ. Report 36, Ottawa.

Ryabchikov, I. D., V. I. Rekharsky, and A. V. Kudrin. 1981. Mobilization of molybdenum by magmatic fluids in the process of crystallization of granitic melts. Geokhimiya 8:1243-1245.

Sasaki, Y. and L. G. Sillen. 1967. Equilibrium studies of polyanions, 16. Equilibria of molybdates in 3M Na(ClO^) medium at 25°C. Arkiv. For. Kemi. 3:253-277.

Shapiro, K. Y. and I. V. Volk-Karachevskaya. 1969. Solubility of molybdic acid in hydrochloric, sulphuric, and nitric acids. Russ. J. Inorg. Chem. 14:571.

Shelton, K. L. 1983. Composition and origin of ore-forming fluids in a carbonate-hosted porphyry copper and skarn deposit: A fluid inclusion and stable isotope study of Mine Gaspe', Quebec. Econ. Geol. 78:387-411.

Shelton, K. L., C. S. So, D. M. Rye, and M. E. Park. 1986. Geologic, sulfur isotope, and fluid inclusion studies of the Sannae W-Mo Mine, Republic of Korea: Composition of sulfur isotope systemetics in Korea W deposits. Econ. Geol. 81:430-446.

Shelton, K. L., R. P. Taylor, and C. S. So. 1987. Stable isotope studies of the Dae Hwa tungsten-molybdenum mine. Republic of Korea: Evidence of progressive meteoric water interaction in a tungsten- bearing hydrothermal system. Econ. Geol. 82:471-481.

Sillen, L. and A. E. Kartell. 1964. Stability constants of metal-ion complexes. Chem. Soc. (London) Spec. Paper 17:754.

Skinner, B. J. and P. B. Barton. 1973. Genesis of mineral deposits. Ann. Rev. Earth Planet Sci. 1:183-211.

Smith, R. W. 1983. Aqueous chemistry of molybdenum at elevated temperatures and pressures with applications to porphyry molybdenum deposits. Ph.D. Dissertation, New Mexico Inst, of Mining and Technology.

So, C. S., D. M. Rye, and K. L. Shelton. 1983a. Carbon, hydrogen, and sulfur isotope and fluid inclusion study of the Weolag tungsten- molybdenum deposit. Republic of Korea: Fluid histories of metamorphic and ore-forming events. Econ. Geol. 78:1551-1573.

So, C. S., K. L. Shelton, D. E. Seidemann, and B. J. Skinner. 1983b. 101

The Dae Hwa tungconsten-raolybdenxira mine, Republic of Korea; A geochemical study. Econ. Geol. 78:920-930.

Sourirajan, S. and G. C. Kennedy. 1962. The system H20-NaCl at elevated temperatures and pressures. Am. J. Sci. 260:115-141.

Susak, N. J. 1981. Spectra, thermodynamics, and molecular chemistry of some divalent transition metal chloro-complexes in hydrothermal solution to 300°G. Unpublished Ph.D. thesis, Princeton Univ., 265pp.

Tingle, T. N. and P. M. Fenn. 1982. The concentration and transport of molybdenum in magmatic systems, experimental evidence. Geol. Soc. Am. Abstracts with Programs 14:240.

Tugarinov, A. I., I. L. Khodakovskiy, and A. P. Zhidikova. 1973. Physicochemical conditions for molybdenite production in hydrothermal uranium-molybdenum deposits. Geochem. Int. 10:731- 739.

Turekian, K. K. and H. K. Wedpohl. 1961. Distribution of the elements in some major unites of the Earth's Crust. Geol. Soc. Am. Bull. 72:175-192.

Valkovic', V. 1983. Trace elements in coal. Vol. 1, 210pp. CRC Press, Boca Raton, Fla.

Vokes, F. 1963. Molybdenum in Canada. Canada Geol. Surv. Econ. Geol. Report 20.

Walder, I., D. I. Norman, and R. Smith. 1988. Calculations of molybdenite solubility, and its use in explaining the molybdenite mineralizations in the Drammen Granite, Norway. Proceedings of the third International Conference of the Mineral Exploration Research Institute, Abstracts 1988:24.

Wallace, S. R., N. K. Muncaster, D. C. Jonson, W. B. MacKenzie, A. A. Bookstrom, and V. E. Surface. 1968. Multiple intrusion and mineralization at Climax, Colorado, p. 605-640. In J. D. Ridge (ed.) Ore deposits of the United State, 1933-1967 (Graton-Sale Vol.). Am. Inst. Mining Metall. Petroleum Engineers, New York.

Wallace, S. R., W. B. MacKenzie, R. G. Blair, and N. K. Muncaster. 1978. Geology of the Urad and Henderson molybdenum deposits, Clear Creek County, Colorado, with a section on a comparison of these deposits with those at Climax, Colorado. Econ. Geol. 73:325-368.

Wesolowski, D., S. E. Drummond, R. E. Mesmer, and H. Ohmoto. 1984. Hydrolysis equilibriua of Tungsten (VI) in aqueous sodium chloride solutions to 300°C. Inorg. Chem. 23:1120-1132. 102

Westrich, H. R. 1974. The solubility of molybdenite in pH buffered KCl-HCl fluids. EOS 55:1200.

White, W. H., A. A. Bookstrom, R. J. Kamilll, M. W. Ganster, R. P. Smith, D. E. Ranta, and R. C. Steinlnger. 1981. Character and origin of Climax-type molybdenum deposits. Econ. Geol. 75th Anniv. Vol. 1981:270-316.

Wood, S. A. and D. A. Crerar. 1985. A numerical method for obtaining multiple linear regression parameters with physically realistic signs and magnitudes: Applications to the determination of equilibrium constants from solubility data. Geochim. Cosmochim. Acta 49:165-172.

Wood, S. A. and D. Vlassopoulos. 1989. Experimental determination of the hydrothermal solubility and speciation of tungsten at 500 °C and 1 Kb. Geochim. Cosmochim. Acta 53:303-312.

Wood, S. A., D. C. Crerar, and M. P. Borcsik. 1987. Solubility of the assemblage pyrite-pyrrhotite-magnetite-sphalerite-galena-stibnite- bismuthinite-argenite-molybdenite in H20-NaCl-C02 solutions from 200° to 3500c. Econ. Geol. 82:1864-1887.

Wright, J. M., W. T. Lindsay, and T. R. Druga. 1961. The behavior of electrolyte solutions at elevated temperatures as derived from conductance measurements. USAEC Report WAPD-TM-204.

Zhidikova, A. P. and S. D. Malinin. 1972. Solubility of powellite CaMoO^ in aqueous NaCl at 50-300°C. Geokhimiya 1:28-34. 103

ACKNOWLEDGMENTS

I gratefully acknowledge the encouragement and support of Drs. S.

Richardson and C. Richardson throughout this research project. I would also like to thank Dr. Nordlle, Dr. Windom, and Dr. Franzen for their critical review of the thesis. I would especially like to thank Dr.

McCarley for his constructive discussion on molybdenum speciation. This research was funded partly by National Science Foundation grant EAR 81-

15852 to C. Richardson.