Experimental determination of the solubility of the assemblage paragonite, albite, and quartz in supercritical H2O

Gwchimice A Cosmmhimica Acttr Vol. 51, pp. 365-372 0 Pagamon Ltd. 1987.R'in+zdin U.S.A.

Journals

Experimental determinationof the solubility of the assemblage paragonite, albite, and quartz in super&itical Hz0 ALAN B. WOODLAND

and JOHN V. WALTHER

Department of Geological Sciences, Northwestern University, Evanston, IL 60201, U.S.A. (ReceivedJuly 24, 1986; accepted in revisedform November 26, 1986) Abstract-The concentrations of Na, Al, and Si in an aqueous fluid in equilibrium with natural albite, paragonite, and quartz have been measured between 350°C and 500°C and 1 to 2.5 kbar. Si is the dominant sohne. in solution and is near values reported for quartz solubility in pure H20. At 1kbax the inanitions of Na and Al remain fairly constant from 350°C to 425°C but then decrease at 450°C. At 2 kbar, Na increases slightly with increasing temperature while Al remains nearly constant. Concentrations of Si, Na, and Al all increase with increasing pressure at constant temperature. The molality of Al is close to that of Na and is nearly a log unit gmater than calculated molalities assuming Al(GH)j is the dominant Al species. This indicates a Na-Al complex is the dominant Al species in solution as shown by ANDERSON and BURNHAM (1983) at higher temperature and pressure. The complex can be written as NaAl(OH): + nSiOr where n is the number of Si atoms in the complex. The value of n is not well constrained but appears to be less than or equal to 3. The results indicate Al can be readily in pure Hz0 solutions at temperatures and pressures . transnorted _ as low as 350°C and 1 kbar.

INTRODUCMON

BURNHAM (1967,1983)and ANDERSONet al. (198’7) demonstrate that in chloride-free solutions at temperatures and pressures of 6OO’C and 2 kbar and above, THERE CAN BE little doubt that most metamorphic reNa and K form stable complexes with Al which raises actions take place in the presence of a fluid phase. The the molality of AI in solution close to that of the total ability of this fluid to transport material from reactant alkali concentrations. to product sites can control the degree of equilibrium The common occurrence of mica, feldspar, and attained and the nature of the mined-fo~ation proquartz in metapelites suggests the concentrations of cess (WALTHER and WOOD, 1984). When metamorah&is, Al, and Si in metamorphic fluids may be latgely phic reactions ate written, aluminum is ofhen conserved controlled by equilibria between this assemblage and between solid phases. Mass transfer studies often use the fluid phase. In light of the potential importance of a fixed ahuninum reference frame. This tendency stems this assemblage, a series of experiments in the simplified in part Tom the fact that the solubilities of aluminous Nabearing system were undertaken to gain insight into phases in near neutral pH solutions at s&ace conditions (e.g. LIKENS et al., 1977; HEM and ROBERSON, the behavior of the Na-Al complex and its possible role in the transport of Al during rne~rno~~srn. Ap1967) and in many hydro~e~~ en~nmen~ are plication of the phase rule to the assemblage albite extremely low (MEYER and HEMLEY, 1967). Addi+ paragonite + quartz C H&I reveals the fluid comtionaliy, the classic study of CARMICHAEL(1969) SUG position is fixed at constant temperature, pressure and cessfully explained many textural relationships for the activity of HrO. The measured molalities of Na, Al, reaction of kyanite to sillimanite by locally conserving and Si in an Hz0 solution in equilibrium with this aluminum. Yet, careful studies of medium to high solid phases assemblage is, therefore, uniquely detergrade pelitic schists seem to require extensive migration of Al during metamorphism (-INNER, 1961;FOSTER, mined by pressure and temperature. The stability of the Na-Al complex relative to the chemical potentials ~@~~;YA~LEY, 1977; WATKINS, 1983).Indeed the of Na, Al and Si can then be determined from the occurrence of feldspar and mica in high tem~m~re solubility data. A recent survey of metamorphic mica hydrothermal veins suggests an appreciable amount of indicates paragon&. is more common Al can be transported. Because of the ubiquity of Al composition than previously believed (Gurrxrrrr, 1984). Therefore, and its mobility in a variety of metamorphic settings, the relevance of these experiments to natural systems controls on the transport of Al are of considerable imis more direct than one might tacitly assume. portance. If a fluid phase is responsible for the transport of Al, the dominant aluminum-bearing aqueous species EXPERIMENTAL PROCEDURES AND REgULTS must be characterized. Studies of corundum ~lubi~ty in I&O under suThe experiments were performed following the extractionpercritical conditions (RAGNARSDCKMR and WALquench technique described by WALTHERand ORVILLE(1983) using a Momy-type reaction vessel with a volume of -36 THER, 1985;BECKER et al., 1983) indicate the concentrations of Al(OH)! in solution are low (iog cm3. The charge was filled with - 1.5 grams of natual, optically clear quartz, paragonite, and albite together with d& mAKOHhO - -4.0 at 500°C and 2 kbar). However, extilled, deionized, and decarbonixd water. Paragon&e and albite perimental investigations reported by ANDER.S~N and compositions are given in Table 1I When a sample was taken, 365

366

A. B. Woodland and J. V. Walther TABLE 1. Electron analyses in wt%.

of paragonite

Paragonite SiO?

.u% MgO

47.56 38.41 0.18 0.34 0.54 6.59 nd 0.05

Fe0 K,O Na20 CaO Cl Total 93.67 nd = not detected

microprobe and albite

increase in the log molality of Al with increasing pressurebut the magnitude of the increase is reduced at higher pressures. DISCUSSION

Albite 68.57 19.38

nd nd 0.18 12.13 0.18 nd 100.44

2.8 grams of extracted solution were flushed t?om the apparatus with 0.01 M HCI solution to make up a total of 18 to 20 grams for analysis. If required, small amounts of distilled, deionized, and decarbonized water were then injected into the vessel to achieve the desired pressure for the next sample to be extracted. The time interval between sample extractions was from 2 to 10 days depending on the temperature and pressure and whether temperature and pressure conditions d&red from the previous extraction. The samples were analyzed for Na, Al, and Si with a direct cut-tent plasma spectrometer (DCP). As the feldspar and paragonite contained minor K this was also analyzed for. None was detected. Analytical reproducibility was within +5 percent for Na and Al and witbin +2 percent for Si. Experimental temperatures and pressures were between 350°C and 500°C and 1 to 2.5 kbar. The maximum temperature for a given pressure was limited by the paragonite breakdown reaction: quartz + paragonite P albite + andalusite + H20.

A primary goal of this study is to derive a set of aqueous species that accounts for the total measured Na, Al, and Si in solution with the constraints on their chemical potential imposed by the mineral phases. We also require this set of species to satisfy the results of previous investigations in simpler systems. To this end, we commence with a discussion of the uncharged silica and alumina species known to exist in the experimental solutions. Since the mineral assemblage includes quartz, the concentration of the species SiOXaqj is buffered by quartz solubility. The close agreement between the quartz saturation curve of WALTHER and HELGESON (1977) and the silica concentrations measured in this study indicates any complexing of Si with Na or Al, if it exists, must be minor in comparison with the species SiO2(,,+ We can, therefore, conclude that the measured Si in solution exists dominantly as the Si02(w9,species buffered by the reaction: Si02cqtz) @

The molality of the species Al( the reaction:

(2) is buffered by

Hz0 + 0.5 NaAISSi30,0(OH)2 * 0.5 NaAlS&Os + AI(O

(1)

The equilibrium temperature and pressure for reaction (1) wascomputed from the “SUPCRT” program of Helgeaon and coworkers (HELGESON et al., 1978). The temperature gradientwithin the chargevolume was - 1“C and the pressure uncertainty was approximately 10 bars. When the molalities of Na or Al increased by 10% or more from the previous sample, equilibrium was considered to be approached from undersaturated conditions. Likewise, the approach was considered to be born supersaturatedconditions if the molalities of Na or Al decreased by 10%or more. When the molality differed by less than 10%between successivedeterminations, the two determinations were considered the same and the direction of approach to equilibrium could not be determined. The experimental results am given in Table 2 and arc shown as a function of temperature at constant pressure in Fig. 1 and as a function of pressureat constant temperatureof 450°C in Fig. 2. The most abundant solute, Si, increases with both temperature and pressure (Figs. la, 2a) and is in reasonably good agreement with the calculated quartz solubility in pure Hz0 of WALTHER and HJXGESON(1977). The concentration of Na in solution is approximately one order of magnitude less than Si. At 1 kbar, the molality of Na remains fai;ly constant with increasing temperature until it decreases at 45O’C (Fia. lb). At 2 kbar. the molality of Na increases slightly with &n&ing temperature. The decrease in concentration at high temperature observed at 1 kbar does not seem to occur at 2 kbar. The log molality of Na increases with increasing pressure at constant temperature, although this effect becomes leas pronounced at higher pressures (Fii. 2b). The behavior of Al is sympathetic to that of Na with molalities generally 0.2 log units lower than Na (Rg. lc). Concentrations of Al remain relatively constant with increasing mmpemmmexceptat 1 kbarwhereAldecma~~above425”C. Like Na, this behavior is not apparent at 2 kbar. There is an

SiOz(as).

(3)

Because the solid phases and Hz0 can be considered to be pure, their activities are unity at the temperature and pressure of interest using a standard state of the pure phase at any temperature and pressure. For the aqueous species we choose a standard state of a hypothetical one molal solution referenced to infinite dilution at any temperature and pressure. With this standard state, the activity coefficient for the neutral aqueous complex, AI(O can be regarded as unity. The molality of Al(OH)! can be calculated from the mass action equation of reaction 3: AG: = log

-2.303RT

mAI(

where AG! is the standard state molar Gibbs free energy of reaction, R is the gas constant and T is temperature in K. Computed molalities of Al( at 1 and 2 kbar, given by reaction (3), calculated from standard state molar free energy data for albite, paragonite, and water from HELGESONet al. (1978) and the standard state molal free energy of A&OH)! from RAGNAR~DOTIXR and WALTHER(1985), are shown in Table

3.

Comparison with the measured molalities (Table 2) indicates Al concentrations are approximately 1 order of magnitude higher than that predicted assuming AI(O is the only Al species in solution (Fig. lc). Clearly AhOH)! cannot be the only Al species present. Based on the studies of the solubility of the assemblage andalusite + quartz in supercritical HzO, the amount

367

Paragon&e,albite, quartzsolubiities

TABLE2. Experimental Results. Sample

P

T

Dur.

No.

(ban)

(V)

(days)

N-l N-2 N-3 N-4 N-5 N-8 N-7 N-8 N-9 N-IO N-11 N-12 N-13 N-14 N-15 N-18 N-17 N-18 N-19 N-20 N-21 N-22 N-23 N-24 N-25 N-28 N-27 N-28 N-29 N-30 N-31 N-32 N-33 N-34 N-35 N-38 N-37 N-38 N-39 N-40 N-41 N-42 N-43 N-44 N-45 N-48 N-47 N-48 N-49 N-30 N-51 N-52 N-53 sU

=

1220 1065 1040 950 975 1210 1000 1020 1005 1005 1000 1005 990 1020 1100 995 1045 1520 1550 1510 2035 2900 2029 2135 2025 2020 1935 2070 2050 1985 2000 2475 2530 2505 2075 2015 1980 2490 2470 2500 2710 2500 2495 2460 2490 2500 2480 2450 2580 2495 2005 2010 2075

365 355 357 353 345 385 378 378 402 401 399 427 424 425 450 448 458 453 468 454 458 432 453 428 427 427 421 408 403 398 399 448 454 451 480 474 475 480 475 475 500 500 509 500 4M) 400 395 350 355 355 350 353 380

undersaturated, S 3

10 3 4 3 3 4 3 2 5 2 4 5 3 2 8 8 3 10 3 3 4 3 2 5 2 2 3 3 2 2 2 5 3 4 3 5 2 3 2 2 7 4 3 2 8 8 8 8 2 3 4 3 B -

Log msq -1.591 -MS08 -1.908 -1.6OS -1.838 -1.518 -1.587 -1.534 -1.492 -1.509 -1.519 -1.478 -1.463 -1.483 -1.408 -1.411 -1.388 -1.284 -1.278 -1.283 -1.213 -l.lSS -1.218 -1.212 -1.324 -1.313 -1.298 -1.358 -1.357 -1.373 -1.392 -1.234 -1.228 -1.217 -1.188 -1.189 -1.185 -1.182 -1.185 -1.187 -1.076 -1Slf -1.118 -1.110 -1.307 -1.315 -1.350 -1.483 -1.488 -1.473 -1.517 -1.343 -1.548

Log mNa -2.866 -2.934 -2948 -2.946 -2.975 -2.910 -2.975 -2.983 -2.852 -2.988 -2.982 -3.017 -3.022 -2.984 -3.196 -3.240 -3.208 -2.840 -2.832 -2.838 -2.590 -2.583 -2.585 -2.391 -2.414 -2.443 -2.490 -2.506 -2.586 -2.599 -2.802 -2.357 -2.394 -2.438 -2.480 -2.539 -2.563 -2.450 -2.437 -2.452 -2.344 -2.3GO -2.419 -2.435 -2.464 -2.468 -2.475 -2.638 -2.582 -2.553 -2.585 -2.833 -2.859

Log

mM

-3.181 -3.204 -3.225 -3.233 -3.225 -3.118 -3.198 -3.223 -3.220 -3.210 -3.180 -3.254 -3.218 -3.220 -3.378 -3.398 -3.384 -3.032 -3.054 -3.049 -2.935 -2.853 -2.903 -2.892 -2.791 -2.819 -2.834 -2.847 -2.851 -2.888 -2.902 -2.783 -2.753 -2.777 -2.801 -2.850 -2.873 -2.804 -2.799 -2.794 -2.893 -2.749 -2.773 -2.782 -2.789 -2.749 -2.806 -2.907 -2.914 -2.877 -2.915 -2.965 -2.991

&w-ch’

U

s

U S U S _ S S U _ U U ii S S S U

S U U S

S S S S _

supersaturated, - = same.

of Al that complexes with Si is negligible and cannot account for the elevated Al concentrations observed (WALTHER,unpublished data). This suggests the dominant Al species must contain Na and possibly Si as well. coned BURNHAM (1983) and ANDERSON et al. (1987) have also argued for alkali + Al complexes to account for corundum solubilities increasing on almost a mole-to-mole basis with the addition of NaOH or KOH. Note that in these solutions, no Si is present so a stable complex of alkali + Al exists without Si.

THE Na + Al COMPLEX The presence of a Na + Al zk Si complex, which is indicated by the elevated concentration of Al, is also suggested by the s~~~eti~ behavior of Na and Al shown in Figs. lb and lc. Because the molality of Al in solution is consistently only slightly lower than the molality of Na, it is reasonable to assume a sir&e complex throughout the pressures and temperatures investigated with an Al to Na ratio of 1:1. As concluded

368

A. B. Woodland and J. V. Walther

r-

-25

TernPer.lYr.

00

T*mpe.rat”re

oc

_~T.-_-.I_

Temperature

__ _.._?

‘C

FIG. 1. Log molality for Si (a), Na (b), and Al (c) plotted as a function of temperature at 1 and 2 kbar. In (a) the solid curve is quartz solubility computed from WALTHER and HELGESON ( 1977).

by ANDERSONand BURNHAM (1983), the alkali-Al complex is probably uncharged, although this has yet to be confirmed. Thus we will write the Na + Al + Si complex as NaAl(OH)i k nSiOz,, where n refers to the number of silica atoms attached to the complex. ANDERSON and BURNHAM(1983) have suggested the complex, at least at the higher temperatures and pressures of their studies, may have a hydrated feldspar stoichiometry and n would therefore be 3. This conclusion is based in part on solubility measurements of pegmatite where Si concentrations exceed those for quartz solubility, indicating an additional Si complex other than SiOzC,,. The value of n can be computed by subtracting the molality of the species Si02taqj at quartz saturation, determined from the mass action equation for reaction 2, from the measured molahty of SiOz and then dividing by the molality of NaAl(OH)i . This is valid as long as the species SiOICaqj and NaAl(OH)q k nSiOltPsj are the only significant Si species in solution. Values of n computed in this manner are shown in Fig. 3 for those sample solutions given in Table 4. For the most part, n is between 0 and 3. The 2% ana@tical uncertainty in the measurement of Si, however, gives a considerable uncertainty to n of roughly Z!Z 1. The value of n must therefore be viewed as approximate. An alternative hypothesis is that the amount of silica in excess of quartz saturation is due to the presence of Na + Si complex containing no Al.

~~~

This possibility will be discussed below when we examine the distribution of Na in solution. As not to predicate a value of n in our analysis, we will consider n as equal to zero. The choice of n = 0 or n = 3 will not affect our speciation calculations. With the assumption that ah Al in solution is present only as Al(OH)‘: and NaAl(OH$, the difference between the total measured molality of Al and the calculated molality of Al(OH)! from Eqn. 3 is equal to the molality of NaAl(OH)!j. Having determined the molality of NaAl(OH)j, the difference between the molality of NaAl(OH)i and the measured molahty of total Na will yield the sum of the molalities of Na+, NaOH, and any Na + Si complex ifpresent. We require the dissociation constant for NaOH, KNaon, in order to determine the concentration of the species NaOH in solution. Because KNaH has not been experimentally determined at high temperatures and pressures, we have calculated log KNaoH in a manner similar to that presented by ANDERSONef al. (1987). The first step in this procedure is to calculate the activity of AlzOs, aAzo3,in our assemblage relative to corundum (i.e. a standard state of unit activity of corundum at any pressure and temperature). The equilibrium constant for the reaction: NaAl$iJOlo(OH)n P NaA1Si308 + A1203+ Hz0

:I*Ly

;,-@q

1.0

2.0

7.5 P.e.s”rB,

Kbar

2.5

3.0

(5)

1.0

1.6

Prewe.

2.0

2.6

1.0

Iear

FIG. 2. Leg molality for Si (a), Na (b), and Al (c) plotted as a function of pressure at 450°C.

1.8 Prss*ure,

2.0 Kbllr

2.5

369

Paragon&, albite, quartz sohrbihties

and the equilibrium action for water

TABI,E 3, Calculated activities of Al&It and NaOH and the dbocistion coxwtant of NaOH at 1 and 2 kbar. See text. T

Los &AI.& Log ~IWX

KHp = an+aoH-.

LOgKNSOH

400 425 450 475 500*

lkbar -3.74 -3.86 -4.12 2 kbar -3.32 -0.940 -3.36 -0.758 -3.53 -0.590 -3.58 -0.42% -3.65 -0.273

-0.808 -0.621 -0.446

-2.58 -2.76 -2.88 -1.97 -2.03 -1.97 -2.07 -2.15

KNaOtl=

* The average molality of NA(OH),D at 475 ’ C wan used to compute the values at 5OO’C.

the au* since paragonite, albite, and Hz0 are considered to be pure phases and therefore have unit activities in our chosen standard state. The values of uMfi computed for reaction (5) from the data references given previously are presented in Table 3 for the temperatures and pressures investigated in this study. Next we compute the activity of the aqueous species NaOH, &or+. Anderson and coworkers have used the solubility of corundum in NaOH and KOH to obtain the equilibrium constant for the reaction:

equals

NaOH + 0.5 Al203 + 1.5 Hz0 # Nap

(8)

Again we will assume unit molal activity coefficients for the aqueous species and use the values of KH,O computed from the data of I-IELGESONand KIRKIiAM (1974)andH~~~~so~e2al.(l981).Theresultsofthese calculations are also given in Table 4. Because we now know the distribution of species in our solutions, we can compute KNaOHfrom:

(W)

400 425 450

constant for the dissociation re-

(6)

and its potassium analog based upon the work of PASCAL (1984). They find no pressure dependence of the equilibrium constant nor any change with temperature between 500°C and 6OO’C at 2 kbar. They report a log & of 0.908. As an approximation, we will assume this value of log X@.> is valid to temperatures as low as 400°C and pressures as low as 1 kbar. Therefore, using the mass action expression for reaction 6, the aNti in our experimental solutions can be calculated since the equilibrium constant, the a,,,+,, and the molality of NaAl(OH)i am known. Again we assume the a&o and the molal activity coefficient for NaAl(OH$ are unity. The activities of NaOH have been calculated using the average molality of NaAl(OH)j at each temperature and pr~ure (Table 3). If we equate the activity of the uncharged NaOH complex with its molality by muming the molal activity coefficient is unity, we can obtain the molality of Na+ plus any Na + Si complex in solution by subtraction of mNWo& + mN.oH from the measured tObl molality of Na in our solutions. For the moment, we will assume the extent of Na + Si complexing is small in comparison to Na+. This allows us to assign this difference solely to Na”. These values are given in Table 4. With these assumptions, the pH of our solutions can now be computed with the aid of the charge balance equation: mN.++ r&r+= tin(7)

mNa+mOHmN,OH

Values of log KNpoH are presented in Table 3 and Pig. 4 at 1 and 2 kbar and temperatures between 400°C and 5OO’C. For comparison, the dissociation constant of KOH, obtained from the electrical conductance measurements of FRANCK( 1956) are also plotted. Note that although the dissociation behavior of the two species as a Em&ion oftemperature and pressure is similar, our computed values for &,ou are nearly an order of magnitude lower than those of KKoH. In the context of electrostatic theory, the greater association of Na’ and OH- is consistent with the smaller size, and thus the greater charge density of Na+ as compared tith K+. Apparently the greater association in the sodium system is also observed at lower temperatures (cf: BARIW~~ERNST, 1963, Pig. 9). The values of log (mN~+/m~+~based on the dist& bution of species from our experimental charges are plotted as a function of temperature at 1 and 2 kbar in Fig. 5. For comparison, the equilibrium boundary between albite and pamgonite in the presence of quartz computed from the Helgeson data set for the reaction: II+ + 0.5 NaAlSirOs 8 0.5 N~l~Si~O~~OH)~ + 3 SiOr + Na+

(10)

as well as other phase boundaries in the system NazO - AlrOr - SiOz - Hz0 at quartz saturation are shown in Fig. 5. The experimental points would lie on the albite-paragonite phase boundary if there were perfect agreement between the Helgeson data set and our experimental determinations modified for the speciation assumptions. Agreement is reasonably good considering activity coefficients am taken as unity and the

1

n RG. 3. Values of n cakulati for the samples given in Table 4. VIis the number of Si atoms in the IbAI(OH)$' f n5i02

complex. See text for method of calculation.

A. B. Woodland and J. V. Walther

370

4. Calculated pN and the distribution 1 and 2 kbar. See text.

of Na and Al species in the sample solutionsat

TABLE

1 kbar N-9 N-10 N-11 N-12 N-13 N-14 N-15 N-16 N-17

400 400 400 425 425 425 450 450 450

-4.43 -4.43 -4.43 -4.38 -4.38 -4.38 -4.31 -4.31 -4.31

-3.25 -3.24 -3.21 -3.29 -3.25 -3.25 -3.43 -3.45 -3.44

N-28 N-29 N-30 N-31 N-25 N-26 N-27 N-21 N-22 N-23 N-36 N-37

400 400 400 400 425 425 425 450 450 450 475 475

-4.57 -4.57 -4.57 -4.57 -4.40 -4.40 -4.40 -4.25 -4.25 -4.25 -4.16 -4.16

-2.86 -2.86 -2.90 -2.91 -2.80 -2.83 -2.85 -2.96 -2.87 -2.92 -2.87 -2.90

2

-3.76 -4.22 -4.20 -4.12 -4.23 -4.07 -4.40 -4.53 -4.43

7.49 7.26 7.28 7.45 7.40 7.48 7.53 7.47 7 52

4.32 3.86 3.89 4.02 3.91 4.07 3.89 3.75 3.86

-2.82 -2.96 -2.94 -2.94 -2.72 -2.74 -2.81 -2.89 -2.91 “2.91 -2.87 -2.89

-3.68 -3.96 -3.93 -3.92 -3.44 -3.49 -3.61 -3.74 -3.78 -3.77 -3.66 -3.71

7.03 6.89 6.90 6.91 7.17 7.14 7.08 7.06 7.04 7.04 7.17 7.14

4.21 3.93 3.96 3.97 4.45 4.40 4.27 4.17 4.13 4.13 4.30 4.25

kbar

inherent uncertainties in the experimental and thermodynamic data. Based on this agreement, it appears the assumption of a Na:Al ratio of 1: 1 for the complex is justified. As seen in Fig. 5, the values of log (m&mH+) are generally somewhat less than the activity ratio calculated from the He&son data set. This systematic small discrepancy may be explained in a number of ways. For instance, the calculated stability field of albite relative to the other solid phases, using the Helgeson data set, may not be great enough (i.e. the computed Gibbs free energy of albite is too high) as has been suggested from analysis of some geothermal systems (RAGNARSDOTTIR et uf., 1984; BIRD, pers.commun.). Alternatively, the disagreement may be attributed to our assumption of unit activity coefficients for the aqueous species in the experimental solutions. However, the dilute nature of our samples would seem to preclude activity coefficient effects of sufficient magnitude. Essentially there is a deficiency of Na in our samples

FIG. 4. Log KNIoHand log KKonplotted as a function of temperature at 1 and 2 kbar. The values of log Kw~ are calculated in the manner discussed in the text. Lq iKKonis from FRANCK ( 1956).

-3.18 -3.41 -3.39 -3.43 -3.49 -3.41 -3.65 -3.71 -3.66

relative to that computed from the Helgeson data set if the dominant aluminum species has a Na:AI ratio of 1: 1. If a Na + Si complex were significant, the discrepancy with the experimental data would be even greater; shil%ng values of log (mNa+/mH+)even lower. The only way to include a significant amount of Na + Si complex@ while retaining the activity ratios shown in Fig, 5 is to Iower the ratio of Na to Al in the dominant aluminum complex. This, however, would make the solubility of corundum in NaOH solution measured by ANDERSON and BURNHAM (1967)difficuh to explain. We therefore conclude that Na-Si complexing, if it exists, is minor in our experimental SO&ions. From the above calculations, the stability of NaAi(OH)z relative to other species in solution can be obtained from the reaction:

FIG. 5. Stability relations in the system Nat0 - A1203SiOz- Hz0 at quartz saturation at I and 2 kbar as a function of tempera~meand log (a&+/&+).The symbols are the values of log (mN.+/mH+) determined from the solubility experiments of the assembIageparagcmite,albite, and quartz.

371

Paragon&e,albite, quartz solubilities

H+ + NaAl(OH)! sztNa+ + Al(OH)I:-i-HzO.

( 11)

Log K&,), based on the experimental measurements and the assumptions atated above, is plotted as a function of temperature at 1 and 2 kbar in Fig. 6. At 2 kbar, kill1 increases from 490” to 425% and then remains fairly constant. The data at I kbar do not permit any trend to be discerned. Note that all values of log qI1, from 400’ to 475’C at 1 and 2 kbar are in the restricted range of one log unit; between 2.2 and 3.2 with no tendency for the stability of the NaAl(OH)i complex to decrease relative to other speciesin solution as temperature is decmased. An alkali-Al complex may, therefore, also be important in hy~othe~~ solutions at temperatures or pressures even lower than those in this study. CONCLUDING REMARKS

In the presence of Na, the molality of Al is greatly elevated above that predicted for AlfOH)‘:upturn;

enhancing the mobility of Al in supercritical fluids (Fig. lc). Therefore, the common practice of conserving Al between minerals in metamorphic reactions is not always justified if our experiments in pure Hz0 can be extrapolated to natural systems. Preliminary investigations in the analogous K-bearing system indicate K also forms a similar stable complex with Al to ternperatures and pressures as low as 1 kbar and 350°C

which sign~~nd~ increases the applicability of this study to natural systems (WOODLAND, unpublish~ data). The study of ANDERSONef al. (1987) suggests the stability of the alkali-Al complex may be somewhat reduced in 0.1 to 1.O molar alkali chloride solutions. The presence of Cl does not preclude significant transport of Al, however. The phlogopite segregations in the contact between dobmite and quartz veins described by WALTHER( 1983) indicate Al transport can occur in solutions with 5 equivalent weight percent

NaCl. The occurrence of feidspar and mica in hydrothermal veins suggestsAl transport can be appreciable in solutions with greater than 40 equivalent weight percent NaCl (BwOM, 1981; ROEDDER, 1984). Whether the N~(OH~~ is the dominant Al complex in these environments

awaits further investigation.

Acknowledgements-We wish to thank G. M. Anderson and

J. J. Hemley for their pemeptivereviewsand Cberil Cheverton for her contribution to the preparation of the manuscript. Financial support wasprovided in part by NSF EAR-8519237. Editorial handling: D. M. Shaw

REFERENCXS ANDERSONG,

M. and BURNHAMC. W. (1967) Reactions of quartz and corundum with aqueous chloride and hydroxide solutions at high temperatures and pressures.Amer. J. Sci. 265, 12-27.

ANDERSONG.

M. and BURNHAMC. W. (1983) Feldspar solubiity and the transport of aluminum under metamorphic conditions, Amer. J. Sci. 2384 283-297. ANDERSONG. M., PASCALM. L. and RAO J. (1987) Aluminum sueciation in metamornhic fluids. In Chemical Transport i; Metmomatic Proceises, (ed. H. C. HELGESON) D. Reidel Pub. Co. (in press). BARN%H. L. and ERNSTW. G. (1963)Ideality and ionization in hy~othe~~ fluids: the system MgCI- Hz0 - NaOH. Amer. J. Sci. 261, 129-150. BEC!KERK. H., CEMIC L. and LANGER K. E. 0. E. (1983) Solubiity of corundum in supercritical water. Geochim. Cosmochim. Acta 47, 1573-1578.

BLOOM M. S. ( I98 1) Chemistry of fluid inclusions: st~kwork

molybdenum deposits from Questa, New Mexico and Hudson Bay Mountain and Endako, British Columbia. Econ. Geol. 76, 1906-1920. CARMICHAELD. M. (1969) On the mechanism of prograde metamorphic r&ons in quartz-bearing p&tic rocks. Corttrib.mineral. Petrol 20, 244-267, C?HRWER G.

A. (1961)The origin of &lima&e in Glen Clova, Angus. J. Perrol.2, 3 12-323. FOSTERC. T. JR.( 1977)Mass transfer in sillimanite-bearing pelitic schists near Rangeiey, Maine. Amer. Mineral. 62, 127-746.

FRANCK E.

2.3 -

y 3

2.7.. _

0

2.5-*

FIG. 6. The logarithm of the equilibrium constant for reaction (11) at 1 and 2 kbar computed from the values given in Table 4.

U. (1956) Hcchverdichteter Wassetimpf III. Ionendissoxiation von HCL, KOH, und Hz0 in uberkrit&hem Wasser. Z. physik. Chem. Neue Folge 8, 192-206. GWDO~ C. V. (1984) Micas in metamorphic rocks. In Micas (ed. S. W. BAILEY), Vol. 13, pp. 357-468. Mmeralogical society of America. HELGESON H. C. and KIRKESAMD. M. (1974) Theoretical prediction of the thermodynamic behavior of aqueous eiectrolytesat high pressuresand temperatures: I. Summary of the thermodynamic/electmstaticpropertiesof the solvent. Amer. J. Sci. 274, 1089-l 198. HEI.GESON H. C., K~KHAMD. H. and FI.ownnsG.

C. (1981) Theoretical prediction of the thermodynamic behavior of aqueous electrolytes at l&h pressures and temperatures: IV. Calculation of activity coefficlenta osmotic coefficients and apparent and relative partial molal properties to 600% and 5&u. Amer. 1. Sci.281, 1249-1516; HELGESONH. C.. DEUNY J. M.. NE.$BIT~ H. W. and BIRD D. K. ( 1978)Summary and critique of the thermodynamic properties of rock-forming minerals. Amer. J. Sci. 278A, 229 p. HEMJ. D. and ROBERSON C. E. (1967) Form and stability of aluminum hydroxide complexes in dilute solutions. U.S. Gee!. Surv. WaterSupply Paper 1827-A, 55 p.

312

A. B. Woodland and J. V. Walther

LIKENSG. E., B~RMANNF. H., PIERCE R. S., EATONJ. S. and JOHNSONN. M. (1977) Biogeochemistryof 4 Forested Ecosystem. Springer-Verlag. 135 p. MEYERC. and HEMLEYJ. J. (1967) Wall rock alteration. In Geochemistry of Hydrothermal Ore Deposits (ed. H. L. BARNES),pp. 166-235. Holt, Rinehart and Winston. PA~CALM. L. (1984) Nature et propribtes des esp&zes en solution dans le systeme KzO - NazO - SiOr - Hz0 - HCl: contribution exp&imentale. These de Doctorat d’etat, Universite Pierre et Marie Curie. RAGNARSM~TTIR K. V. and WALTHERJ. V. (1985) Experimental determination of corundum solubihties in pure water between 400-700°C and l-3 kbar. Geochim. Cosmochim. Act4 49,2 109-2 115. RAGNARSDOTTIR K. V., WALTHERJ. V. and ARNORSSON S. ( 1984) Description and interpretation of the composition of fluid and alteration mineralogy in the geothermal system at Svartsengi, Iceland. Geochim. Cosmochim. Acta 48, 1535-1553. ROEDDERE. ( 1984) Fluid Inclusions. Mineralogical Society of America, Washington DC., 644 p.

WALTHERJ. V. (1983) Description and interpretation of metasomatic phase relations at high pressures and temperatures: 2. Metasomatic reactions between quartz and dolomite at Campolungo, Switzerland. Amer. J. Sci. 283A, 459-485. WALTHERJ. V. and HELGESONh. C. (1977) Calculation of the thermodynamic properties of aqueous silica and the solubility of quartz and its polymorphs at high pressures and temperatures. Amer. J. Sci. 277, 13 15- 135 1. WALTHERJ. V. and ORVILLEP. M. (1983) The extractionquench technique for determination of the thermodynamic properties of solute complexes: Application to quartz solubility in fluid mixtures. Amer. Minernl. 68, 73 l-74 1. WALTHERJ. V. and WOOD B. J. (1984) Rate and mechanism of prograde metamorphism. Contrib. Mineral. Petrol. 88, 246-259. WATKINSK. P. (1983) Petrogenesis of Dah&ian albite porphyroblast schists. J. Geol. Sot. London 140,601-618. YARDLEYB. W. D. (1977) The nature and the significance of the mechanism of shlimanite growth in the Connemara schists, Ireland. Contrib. Mineral. Petrol. 65, 53-58.