Solution geochemistry of lead and zinc in water containing carbonate, sulphate and chloride ions

Chemical Geology, 29 (1980) 293--311

293

Elsevier Scientific Publishing Company, Amsterdam -- Printed in The Netherlands

SOLUTION GEOCHEMISTRY OF LEAD AND ZINC IN WATER CONTAINING CARBONATE, SULPHATE AND CHLORIDE IONS

A.W. MANN and R.L. DEUTSCHER

CSIRO, Division of Mineralogy, Wembley, W.A. 6014 (Australia) (Received July 9, 1979; revised and accepted January 31, 1980)

ABSTRACT

Mann, A.W. and Deutscher, R.L., 1980. Solution geochemistry of lead and zinc in water containing carbonate, sulphate and chloride ions. Chem. Geol., 29: 293--311. The solution geochemistry of lead and zinc is markedly influenced by the formation of complexes with hydroxyl, carbonate, sulphate and chloride ions, and by formation of insoluble precipitates with these same anions. PbSO4 °, PbCO3 ° and Pb (CO3)22- are the dominant Pb ions in solutions containing sulphate, chloride and carbonate anions. Cerussite (PbCO3), anglesite (PbSO4) and Pb~CI(OH)3 are the phases which limit Pb solubility. ZnSO4° and Zn(OH)~ ° are the dominant complex ions of Zn in solutions containing sulphate, chloride, carbonate and hydroxyl anions. Zn4(SO4)(OH)6 , Zn(OH)2 and ZnCO3 (smithsonite) are the phases which limit Zn solubility. Comparison of the calculated total activities for Cu, Pb and Zn suggests that for solutions with pH < 6.0 the order of mobility is Zn > Cu > Pb and for solutions with 6.0 < pH < 8.3, the order of mobility is Zn > Pb > Cu. In alkaline carbonate solutions, the orders of mobility are Pb < Cu < Zn for 8.3 < pH < 9, Pb < Zn < Cu for 9 < pH < 10 and Zn < Pb < Cu for solutions with pH > 10. These orders of mobility in the acid and neutral pH range are in agreement with observed concentrations of base metals in the vicinity of some weathering base-metal sulphide deposits.

INTRODUCTION

Solubilities of copper, lead and zinc in natural waters in the vicinity of base-metal sulphides are of importance in assessing possible mechanisms for ore formation, as a possible exploration tool for their detection, and for minimizing the environmental consequences of mining them. Skinner and Barton (1973) suggested that by far the greatest number of mineral deposits have been formed by precipitation from solution -- this applies particularly to stratiform base-metal sulphide deposits. Some limitations on the possible composition of ore-forming fluids were presented by Barton (1957), and subsequent fluid-inclusion studies (e.g., Roeder, 1976) and reasoning based on observations of naturally-occurring brines (e.g., White, 1968) suggested that many stratiform base-metal deposits owed their existence to mobilizing fluids containing high concentrations of the approp0009-2541/80/0000--0000/$02.25 ©1980 Elsevier Scientific Publishing Company

294 riate metals. King (1973) discussed the frequently-observed lithological and stratigraphic displacement of Cu-sulphides from those of Pb and Zn. Impressions of the chronological evolution of ideas as to the origin of stratiform deposits and the importance of metal-laden brines, were given by King (1976) and Ridge (1976). Hydrogeochemistry has not been extensively employed in the exploration for base-metal sulphide deposits. Some examples for the case of Cu were quoted in an earlier publication (Mann and Deutscher, 1977). Goleva et al. (1970) described the distribution and migration of Pb in groundwaters in Russia, and Shvartsev et al. (1975) summarised some features of the migration of microcomponents in neutral waters of the supergene zone. Cameron (1977a) described the geochemical dispersion in lake waters and sediments from massive sulphide mineralization in the Agricola Lake area, Northwest Territories. The estimated order of mobility obtained from the related soil survey (Cameron, 1977b), for base metals was Zn > Cu > Pb. From experimental studies on the aqueous dispersion of elements around sulphides contained in glass and plastic tanks, Govett and Whitehead (1974) and Govett et al. (1976) indicated that electrochemical processes are probably responsible for dispersing base metals around sulphide ore bodies. However, the electrochemical experiments, which generated dispersion patterns similar to known stratabound deposits, did not offer any explanation for the observed zonation or segregations of base metals. More extensive and rigorous experimentation, particularly of simpler chemical and mineralogical configurations, was suggested in order to determine the contribution of various possible mechanisms to element dispersion. Experiments conducted on sulphide ore from the Woodlawn deposit in New South Wales (Giblin, 1978) suggested that metals in solution were exchanged with metals in the sulphide in the order Cu > Pb > Fe > Zn. In the vicinity of the Woodlawn deposit, the supergene zone was observed to be Cu enriched and the gossans markedly enriched in Pb carbonates and sulphates; Zn was the only metal expected to persist in water over any distance from the deposit. Potter and Nordstrom (1978), in a detailed study of weathering of sulphide ores in two distinct environments, i.e. in situ and in waste piles in Shasta County, California, observed that Cu and Zn were mobilized as soluble sulphates by oxidation of sulphides but Pb remained and was concentrated in the aerated zone as an insoluble residue. Mine drainage from a sulphide deposit at Captains' Flat, N.S.W., Australia, was noted (Craze, 1977) to contain high concentrations of base metals, with molar concentrations in the order Zn > Cu > Pb. The dumps had metal in concentrations in the order Zn > Pb > Cu, whilst nearby river sediments were noted to have Pb > Zn > Cu. Of the two cases reported by Nunn and Riches (1978) from the Northampton area in Western Australia, the first had surrounding waters with Zn > Pb > Cu (molar concentrations) whilst waters around the second deposit indicated Zn > Cu > Pb. Both dumps contained significant concentrations of Cu, Pb and Zn yet the seepage waters contained appreciably different absolute and

295

relative concentrations of basemetals. These observations suggest that the mobilization of base metals in such cases is not simply related to the composition of the ore, but perhaps to other factors such as the solubilities of the elements in the percolating solutions. Barton and Bethke (1960) presented some t h e r m o d y n a m i c properties of Zn and Cu minerals, and Hem (1972) reported the chemistry and occurrence of Cd and Zn in surface waters and groundwaters, and provided information as to the solubility of Zn in some carbonated solutions. Hahne and Kroontje (1973) assessed the significance of chloride concentration on behaviour of heavy-metal pollutants. In an earlier publication (Mann and Deutscher, 1977), the present authors investigated the solubility of Cu in solutions containing carbonate, sulphate and chloride ions. By using available solubility product constants, mineral species limiting the solubility of Cu and the concentration of Cu ion under a given set of solution parameters were determined. From the concentration of Cu ion, and using the available complex ion association constants for the complex ions likely to be present, a total activity for Cu was calculated for a series of pH values for each solution. In this paper, and in a similar manner, the solubilities of Pb and Zn are calculated for a similar set of widely-varying solution parameters. Comparison of the solubilities of Cu, Pb and Zn for these solution parameters suggests relative hydrogeochemical mobilities for these elements. METHOD

OF CALCULATION

Activity--solu bility relations The activities of Pb and Zn species in solution can be calculated from appropriate t h e r m o d y n a m i c data, but to obtain molalities of individual species (and hence solubilities expressed as the sum of these molalities), activity coefficients need to be known. Since for most species under consideration in this paper, activity coefficients are n o t known and solution parameters are not fully defined, we have chosen to sum activities of individual Pb and Zn species to obtain a total "activity" (= Zapb or Zazn). In comparing this "total activity" with solubility, we are effectively assuming that for each species in which the metal is contained (including complexes), the activity coefficient is unity. In dilute solutions, and when zero-charged complexes are dominant, the error involved in this assumption is likely to be small and for such cases the total "activity" provides an estimate which will be very close to the actual (observed) solubility since it does take account of complexion formation. Application to natural waters Activities of the ions OH-, CI-, SO42-, HCO3- and CO32- are used in all calculations here in preference to concentrations, to circumvent the problem

296 of estimating activity coefficients in hypothetical solutions. To apply e x a c t l y t h e s o l u b i l i t i e s d e r i v e d i n t h i s p a p e r t o a n y n a t u r a l w a t e r s y s t e m , freei o n a c t i v i t i e s o f t h e s p e c i e s O H - , Cl-, SO42-, H C O 3 - a n d CO32- i n t h e n a t u r a l w a t e r m u s t b e u s e d . A t h i g h i o n i c s t r e n g t h s (e.g., :>0.3), t h e a c t i v i t y coeffic i e n t o f a m o n o v a l e n t i o n will be o f t h e o r d e r o f 0.7, w h i l s t t h a t o f a d i v a l e n t i o n will b e c o n s i d e r a b l y less. The procedures and coefficients developed by Garrels and Thompson (1962) for seawater can be applied to m a n y n a t u r a l waters to o b t a i n activity coefficients a n d c o m p l e x - i o n a n d free-ion activities.

Solubility T h e s o l u b i l i t y o f P b was c o n s i d e r e d t o b e l i m i t e d b y t h e f o r m a t i o n o f o n e o r m o r e o f t h e c o m p o u n d s o n t h e l e f t - h a n d side o f t h e r e a c t i o n s s h o w n i n Table I and the solubility of Zn by the compounds shown on the left-hand side o f r e a c t i o n s i n T a b l e II. M i n e r a l n a m e s f o r t h e c o m p o u n d s a n d s o l u b i l i t y p r o d u c t s f o r e a c h o f t h e s e r e a c t i o n s are s h o w n i n t h e c o l u m n s a d j a c e n t t o each entry. TABLE I Compounds limiting Pb solubility Mineral

Massicot Anglesite Cotunnite Laurionite Cerussite Hydrocerussite Phosgenite

Reaction Pb(OH)~ ~ PbO + H~O PbSO4 PbCl~ PbC1OH PbClo.s(OH)~.s PbCO3 Pb~(OH)~(CO3)~ PbCl(CO3)0. s

log Ks ¢ ¢ ~~ --¢~ ¢ ¢

Pb :+ + 2 O H Pb 2÷ + 2OHPb 2÷ + SO4:Pb 2÷ + 2C1Pb 2÷ + C1- + OHPb ~÷ + 0.5C1- + 1.5OHPb 2÷ + CO3~3Pb 2÷ + 2OH- + 2CO~2Pb 2÷ + C1- + 0.5CO32-

-16.79 15.3 -7.80 -4.67 13.7 -17.0 13.13 -45.46 9.90

TABLE II Compounds limiting Zn solubility Mineral

Reaction

Zincite

ZnO + H20 Zn(OH)2 ZnCO3 Zn(CO3)0.36OHI.~s Zn(SO4)0.2s(OH)~.~ ZnCl0.~OH~. s

Smithsonite Hydrozincite

log Ks ~ ~ ~ ¢ ~ ¢

Zn 2* + 2OHZn 2÷ ÷ 2 O H Zn 2+ + CO32+ Zn 2+ + 0.36CO3:- + 1.28OHZn ~++0.25SO42- +1.5OHZn :+ + 0.5C1- + 1.5OH-

-16.89 -17.15 -10.78 -14.42 -13.9 -13.4

297

Solubility constants for 25°C, and either extrapolated or corrected to zero ionic strength, were obtained from Sill~n and Martell (1964); no attempt was made to evaluate criticallythe experiments or corrections used to obtain these values.

Complex-ion formation Reactions resulting in complex-ion formation with Pb or Zn were chosen from the comprehensive lists compiled by Sill~n and Martell (1964). Selections were made with preference given for values of equilibrium constants at or near 25°C, and derived by extrapolation of experimental data to zero ionic strength. Tables III and IV contain the complex-ion reactions used and the equilibrium constants for complex-ion formations. TABLE III Association constants for complex-ion formation with Pb Reaction P b 2+ P b 2÷ P b 2+ P b 2+ P b 2+ P b ~+ P b 2+ P b ~+ P b 2+

+ + + + + + + + +

OH2OH3OH C O 3 ~2 C O 3 ~SO42C12Cl3C1-

log K ¢~~ ~ ¢ ~ ¢¢¢-

PbOH ÷ Pb(OH)2 ° P b ( O H ) 3PbCO3 ° P b ( C O 3 ) 2 ~PbSO4 ° PbC1 + PbCl~ ° PbC13-

7.82 10.88 13.94 7.5 8.62 2.62 1.6 1.78 1.68

TABLE IV Association constants for complex-ion formation with Zn Reaction Z n ~÷ + O H Z n 2÷ + 2 O H Z n 2÷ + CO3 ~Z n ~+ + S O 4 2 Z n 2÷ + C1Z n 2÷ + 2 C l Z n 2÷ + 3C1Z n 2÷ + 4C1-

log K ¢~~ ~ ~ ~ ~~

ZnOH ÷ Zn(OH)2 ° ZnCO3 ° ZnSO4 ° ZnCl ÷ ZnC12 ° ZnCl3Z n C 1 , ~-

4.36 12.89 5.3 2.3 0.4 , 0.61 0.53 0.2

298

Carbonate equilibria The following three equations apply to carbonate, bicarbonate equilibria in water at 25°C: K = 10 -1"47

(l)

CO2 (aq) + H2) # H + + HCO3-,

K = 10 -6.4

(2)

HCO3-

K = 10 -1°'3

(3)

CO2 (gas)

@- CO2 (aq),

~ H + + CO32-,

The equilibrium constants for eqs. i and 2 are due to Harned and Davis (1943), and for eq. 3 is due to Harned and Scholes (1941). From these equations, an expression relating the carbonate content of water to the prevailing fugacity* of carbon dioxide and pH can be obtained, viz.: log aco32- = log f c o : - 18.17 + 2 pH

(4)

For the bulk of calculations in this paper, the CO2 content of waters was taken as being controlled by the atmosphere, i.e. 0.03% (fco2 = 10-3's bar), and, unless otherwise stated, carbonate values are related to this value of fco2 and to pH, using eq. 4. Commonly a fixed total activity of combined carbonate species has been used in calculations involving carbonate equilibria, for example, in construction of Eh--pH diagrams (Garrels and Christ, 1965). As has been pointed out (Jenne, 1968) this is not a desirable assumption for surficial materials such as soils and streams, and fixed partial pressures, as used here, more closely approximate natural systems. Field analyses of surface and near-surface underground waters in varying geological environments, usually show carbon dioxide partial pressures to be in the range PCO~ = 10-2--10-4 bar. In a recent drainage catchment study in Western Australia (Mann and Deutscher, 1977) a mean PCO2 = 10-2'°1 bar was observed; values similar to this are c o m m o n l y observed for groundwaters in equilibrium with soil. To show the importance of carbon dioxide fugacity (and hence carbonate concentration at any given pH) on the solubility and in complex-ion formation, other sets of calculations were performed with carbon dioxide fugacities different from 10 -3"s bar. These will be discussed in detail later.

Solution parameters The following activities of chloride and sulfate ion were chosen to be representative of the range of values likely to be found in natural and some

* F o r t h e c o n d i t i o n s c o n s i d e r e d in this p a p e r t h e d i f f e r e n c e b e t w e e n t h e fugacity a n d partial pressure o f CO2 is small.

299

synthetic waters, and also to provide useful information as to the effects of complex-ion formation on Pb and Zn solubility: chloride activities: sulphate activities:

0.0, 0.0,

10 -2, 10 -s,

10 -1, 10 -3,

1.0 10 -1

Taking all combinations of these chosen activities for chloride and sulphate, calculations for complex-ion activities and solubility were carried out in the following manner: (a) H y d r o x i d e and carbonate ion activities were calculated at pH intervals of 0.02-pH units in the range pH = 2.0--12.0, the latter using eq. 4. (b) At each pH, for the given chloride, sulphate and carbonate ion activity, an equilibrium Pb 2÷ or Zn ~÷ activity for each of the c o m p o u n d s in Tables I and II was calculated. (c) The minimum of these is the maximum activity of metal ion in a solution of that composition at equilibrium. (d) Equilibrium activities of all complex ions were calculated, using this value of Pb 2÷ or Zn 2÷ in the equations shown in Tables III and IV. (e) Activities for all complex ions were plotted as a function of pH. (f) The sum of activities of all complex-ion and free-ion activities of Pb and Zn (= ~.apb, ~ a z n ) were plotted as a function of pH, as shown in Figs. 2 and 4. Calculations were carried o u t on the CSIRO computing network, using a simple FORTRAN program incorporating an automatic plotting subroutine. A version of this program w i t h o u t plotting facilities has been developed for a smaller PDP-11 ® computer, using BASIC. RESULTS AND DISCUSSION

Compounds limiting lead solubility Fig. 1 shows the c o m p o u n d s observed to be limiting for the solubility of Pb in a variety of solutions with fco~ = 10-3"s. Cerussite and anglesite are the solid phases which limit the solubility of Pb for solutions with low concentrations of chloride and/or sulphate (Fig. la). At chloride activities of approximately 10-1 and higher, hydroxy--chlorides, e.g. laurionite exhibit stability fields in the pH range 4.0--8.0 (Fig. l b and c). In concentrated chloride solutions, cotunnite PbC12, is the equilibrium solid at low sulphate activities (Fig. lc).

Lead solubilities Fig. 2 shows the activities of the complex ions of Pb listed in Table III, plus the hydrated Pb 2÷ ion, as a function of pH, for a solution with fco2 = 10-3"s, ac1- = 1 and a s o 2 - = 10 -1. Under these conditions, ~he zero-charged complexes PbSO4 ° and PbCO3 ° are the dominant ions fQr pH's up to 10.

300 (a)

J

0 _~

Cerussite

--Anglesite

_~ :-~.

~

I

1

aci-

I .....

=

10 -2

1

(b)

Cerussite L

I

a c i - = l o -1 I

I

(c)

I

Anglesite o _@-.~- . . . . . ~" Irlillilrllllll Cotunnite

~,~'Pb2a(6H)3

Cerussite

llb.~L,,L',~,,r/-/////j ~//,~{/~/,~

_~1 l i= I,illll I,11 II I1£\'~\~,\'~\~\1~',i~/~1 2

~

aCl- =

I

I [I I p L a u r i o n i t e

4

5

6

?

,

8

t

9

pH

Fig. 1. C o m p o u n d s limiting the Pb solubility in aqueous solutions w i t h fco~ = l O - L s and various activities of chloride and sulphate.

-4

• PbSOo

,/,~ kO0~-/

. . . . . PbCi20" ..............................................

.~ . .

_.. Pbcr

-8 -

~

PbCl3-

/

-"

/

~ o~.. ; / ,,,, . / " ,/ o .'"

,.."

, 4

/

.

PbCO# -

",T~-7~~

~o~,

--

"",7"-:"-

°~-.7":'" .'" ~'~:, ~'~--'-"".,

//

// ,/ /Z

.j Zo~~

.,o~ -~" -20

./

,ob

\ "'..Z~ "..
~

.--

./

/

"~"~'.~L"

"~.

/

-12

./__

"'~'°..~

/

~o~,~.~

>, >

-16

.. /

/"

1-

~ o

"

- :-~-'~.

?/ 6

pH

,

,

8

I0

12

Fig. 2. Activities o f Pb c o m p l e x e s in a solution w i t h foo2 = 1 0 - L s , acl- = 0 and aso42 = 10 -1

301

The total Pb activity, expressed as a sum of the individual complex-ion activities, and for the same solution parameters, is shown in Fig. 3. Cerussite produces a minimum Pb activity of 10 -s'Ts (= 0.344 mg 1-1) at pH 9.5. Anglesite, PbSO4, has a limiting activity of 2:apb = 10 -s" 16 or 1.4 mg 1-1 in low-chloride solutions with a s o ~ - = 10 -1. This figure increases marginally due to chloride complexing when aCl = 1. The activity of Pb in equilibrium with cotunnite, PbC12, is Zapb = 10 -2"s° or 660 mg 1-1. iO-Z

i

tt

l

I

2050

I

\ %', IO-S

205

10-4

20.5

O.

w

E 10-5

2"05

iO-I~ I ~

i0-7 2

0 1

0 0

0

10 "1

1

10 "1

'205

I

I

4

s

pH

I

[

tl

io

12

Fig. 3. T o t a l activities for Pb m s o m e a q u e o u s s o l u t i o n s w i t h fco~ = 10-3"s.

Effect of increase of PCO, As pointed o u t earlier, the fugacity of CO2 in the atmosphere (fco2 = 10-3"s) is not always appropriate for waters in contact with soil, particularly groundwaters. Activities at fco2 = 10-2 and thus more appropriate to the latter situation have also been calculated. In the neutral and alkaline pH range where cerussite is the solid limiting Pb solubility, the Pb solubility is virtually unchanged for the increase in f c o , to 10 -2 bar. Only in the chloride solutions on the acid side of neutral pH, when sulphate activity is low, are the solids which limit Pb solubility and Pb solubilities altered. Table V compares the Pb activities in some different solutions at fco~ -- 10-3"s and f c o , = 10 -2 bar. f c o , clearly has only a small effect on the Pb solubility as compared to chloride and sulphate. This is in contrast to the case recorded previously for Cu (Mann and Deutscher, 1977), where dicarbonate complexes, particularly, have a strong influence on Cu solubility in alkaline solution. It is o f note that for Pb the chloro-carbonate, phosgenite, replaces laurionite, and the stability field for cerussite is displaced to lower pH's for the case where fco~ = 10-2 bar.

302 TABLE V E f f e c t o f PCO2 o n Pb solubility for a s o l u t i o n with ac] = 1 and aso4~- = 10 pH

PCO= = 10-3"5

3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5 10.0

PCO 2 = 10-2

limiting solid*

solubility (mg 1-')

limiting solid*

solubility (rag 1-')

Cot Cot Cot Cot Lau Lau Lau Lau Lau Pb2CIOH 3 Pb=CIOH3 Pb2C1OH 3 Cer Cer Cer

660 660 660 660 616 195 61.9 19.8 6.46 2.05 0.576 0.365 0.371 0.343 0.365

Cot Cot Cot Cot Pho Pho Pho Pho Cer Cer Cer Cer Cer Cer Cer

660 660 660 660 489 155 49.2 15.9 2.89 0.608 0.364 0.334 0.340 0.432 1.367

*Cer = cerussite; Cot = c o t u n n i t e ; Lau = laurionite; P h o = phosgenite.

Compounds limiting zinc solubility Fig. 4 shows the compounds observed to be limiting the Zn solubility in a variety of solutions with fco2 = 10-3"5.

(a) Zn(OH)2

cO_-~FZn4(SO4)(OH)61 ~/I -~/1"/ I -I~ "~, _~

I

.,¢,~OH

J (b)

/i Z n4.(S04) (OH)6 --

(~

I

aci- =10-2

//"

Zn(OH)2

aci-. 10-1

, t

1

I

I

I

I

(c) -'~Zn4(SO4)(OH)6~ ~O F'//// Zn2Cl(OHY3Y/////'J

z n,o.,,

acl-. 1

Fig. 4. C o m p o u n d s limiting the Z n solubility in a q u e o u s s o l u t i o n s w i t h f c o : = 10- ~"s and various activities o f c h l o r i d e and sulphate.

303 Zinc hydroxide and Z n 4 ( 8 0 4 ) ( O H ) 6 , are the solid phases which limit Zn solubility in most solutions. The former is stable in neutral and alkaline solutions, the latter in acid solution. Zinc hydroxy--chloride, Zn2CI(OH)3, limits solubility of Zn in acid chloride solutions with low as02- (Fig. 2b and c).

Zinc solubility Fig. 5 shows the activities of the complex ions of Zn listed in Table IV, plus the hydrated Zn 2÷ ion, as a function of pH, for a solution with fco~ = 10-3"s, aCl- = I and aso4 ~- = 10 -1. Under these conditions, the zero-charged complexes ZnSO4 ° and Zn (OH)2 ° are the dominant ions. The total Zn activity, expressed as a sum of the individual complex-ion activities, and for the same solution parameters, is shown in Fig. 6. This curve, showing Zn to be extremely soluble in acid solution and relatively insoluble for pH > 8, is representative of all solutions studied at this f c o • The activity of Zn in a solution in equilibrium with solid Zn(OH)2 is 10 -4"~4 or 3.78 mg 1-1. Both of the solids, Z n 4 ( S O 4 ) ( O H ) 6 and Zn2CI(OH)3 are exceedingly soluble. In equilibrium with Zn4(SO4)(OH)6 and Zn2CI(OH)3, Zn has an activity, Zazn, > 1 for pH < 5.

ZnS040

........ z;,EF"z;~c,?"~;~*":-~'~.~"... -2 ~Zn

T

,,~,% ,%.

... ~-"~.-":'..~

~o_y,~._-

-6

~5

O0~H *

~

""~" ~ ' ~~o~ ~ * .----

~

~--"

Zn(OH) 0

",0,, ",,

. ~

~

ZnCC~ °

,~. ".,,., "~.

~

. ~

"':":.',"'.~'~ "',',, ~o ..~,:~."" -... .

"',C>.~._ I • "-,,',,",,~.. I z°cJ---~,.., "." "-.1 z°ci~

-IC

"'<",'"',. I

-14

I

4

I

6

I

pH

8

IO

12

Fig. 5. Activities of Z n complexes in a solution with fc02 = 10-3"5, acl- = 0 a n d as042_ = 10 -1.

304 10-2

j

I

650

10-3

65

~

10-4

,5

6'5 7

W

E

iO-S

iO-e

10-7

fco=

= 10"a'5

aci-

= 1

.65

065

a s o z - : lo -~

I 4

I e

oH

I e

[ I0

12

Fig. 6. Total Zn activity as a function of pH for an aqueous solution with acv = 1, aso42- = 10 -l and fco 2 = 10 -3"s.

Effect o f increase in PCO: Zn solubility calculations f o r f c o : = 10-2 bar indicate t h a t the m a j o r e f f e c t o f increasing CO2 pressure is t o decrease the solubility o f Zn in the neutral and alkaline region. S m i t h s o n i t e , ZnCO3, replaces Zn (OH)2 as the solid limiting the solubility o f Zn. In Table VI, the solubilities o f Zn f o r fco~ = 10 -a's and fco~ = 10-2 in a solution with ac1- = 10 -1 and a s o / - = 10 -1 axe compared. A t fco2 = 10-2, s m i t h s o n i t e p r o d u c e s a Zn solubility, 2~azn, o f 10 -s'27 or 0 . 3 5 0 mg 1-1. F o r p H > 6, the Zn solubility is decreased b y at least an o r d e r o f m a g n i t u d e b y the increase o f fco2 f r o m 10 -3"s t o 10 -2 bar.

Comparison of Cu, Pb and Zn solubilities Many h y d r o g e o c h e m i c a l processes relating to ore deposits o c c u r in or are c o n t r o l l e d by, g r o u n d w a t e r , i.e. sub-surface w a t e r in i n t i m a t e c o n t a c t with soft, w e a t h e r e d r o c k a n d / o r u n w e a t h e r e d material. F o r this reason, and because in m a n y cases the c o n t a c t t i m e f o r soil and g r o u n d w a t e r is sufficiently long t o enable " e q u i l i b r i u m " to be a p p r o a c h e d , c o m p a r i s o n o f Cu, Pb and Zn activities as calculated here is m o s t meaningful f o r the situation w:~ere fco~ = 10-2 bar. Fig. 7 shows the activities f o r Cu, Pb a n d Zn as a f u n c t i o n o f pH, for a s o l u t i o n in which aCl- = 1, aso42- = 10 -1 and fco~ = 10-2. T h e activities for Cu were calculated b y the m e t h o d described in Mann and D e u t s c h e r (1977).

305 T A B L E VI Effect o f P c o 2 on Zn solubility for a solution with ac1- = 10 -1, aso4 2- = 10 -1 pH

PCO~ = 10-3's

3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5 10.0

PCO~ = I0-2

limiting solid*

solubility (mg 1-1)

limiting solid*

Gos Gos Gos Gos Gos Gos Gos Gos Zn(OH)2 Zn(OH)2 Zn(OH) 2 Zn(OH)2 Zn(OH) 2 Zn(OH)2 Zn(OH)~

> > > > 983,000 175,000 31,100 5,530 987 102 13.6 4.76 3.88 3.79 3.78

Gos Gos Gos Smi Smi Smi Smi Smi Smi Smi Smi Smi Smi Smi

solubility (mg 1-1 )

> > > > 365,000 36,500 3,650 366

Gos

36.7 4.0 0.716 0.387 0.354 0.350 0.350

>

= greater than 10 ~ m g I-I. *Gos = goslarite; Smi ffismithsonite.

-3

l >~-4 >

fc02 =10"2 aci- ffi1 aso J

-6

• lO -1

I

Pb(Cu(Zn .7

2

I 4

Cu(Pb(Zn

I s

Pb(Ou
]tZn(Pb(Cu

I _u

pn

8

io

12

Fig. 7. T o t a l activities o f Cu, Pb and Zn as a f u n c t i o n o f p H in an aqueous s o l u t i o n with acl- = 1, a s o 2 - = 10 -t a n d f c o 2 -- 10 -2.

306

Clearly, the solubilities of these three elements do not have the same dependence on pH in this solution (or in any other solution studied in this work) and there are four different orders of relative mobilities of Cu, Pb and Zn, depending on pH. At low pH and when 8.5 > pH < 10, Pb is least mobile (on a mole for mole basis) but in the pH range of 6.2--8.5 (a c o m m o n pH range for naturally-occurring groundwaters) Cu is the least soluble element. For any given pH towards the centre of this range, the differences in solubilities are very marked. For example at pH = 7, from Fig. 7, the activity of Cu, EaCu = 10 -s'36 (0.280 mg 1-1); of Pb, Z a p b = 10 -4.27 (11.1 mg l-L); and of Zn, Eazn = 10 -3.07 (56.3 mg 1-1).

Application of results Table VII lists some orders of mobility for Cu, Pb and Zn obtained from the literature for in situ and laboratory experiments on some naturallyoccurring sulphide deposits. All of these suggest Zn to be the most mobile of the base metals, but there is little apparent consensus as to whether Pb or Cu is the least mobile. Detailed solution parameters, apart from pH, e.g. aCl- , a s o 2 - and fco~, are not available for the solutions from which the data in Table VII were obtained and it is not possible to correlate exactly these solutions with data from our work. However, using pH as a general indicator, solutions with pH < 6 have Pb least mobile and those with 6.0 < pH < 8.3 have Cu least mobile, as suggested by our calculations and as depicted, for example, in Fig. 5. Govett et al. (1976) described the deficiency of Cu (around the weathering sulphides of their experiments) as puzzling. A simple leaching experiment with crushed sulphide in 0.1 M NaC1 after 49 hr. resulted in concentrations of Zn, Pb and Cu of

T A B L E VII Order o f m o b i l i t y for base metals f r o m s o m e r e c e n t literature Location of ore material

A u t h o r (s)

Mobility order

Solution parameters

Agricola Lake (N.W.T., Canada) Woodlawn (N.S.W., Australia) N e w Brunswick (N.W.T., Canada) Shasta C o u n t y (Calif., U.S.A.) Captain's Flat (N.S.W., Australia) Northampton (W.A., Australia)

C a m e r o n (1977a, b)

Zn > Cu > Pb

p H = 3.0--5.0

Giblin ( 1 9 7 8 )

Z n > Pb ~ Cu

G o v e t t et al. (1976)

Zn > Pb > Cu

pH = 6.8--8.0 [SO4 :- ] = 230 mg 1-~ pH = 7.0

P o t t e r and N o r d s t r o m (1978) Craze (1977)

Zn, Cu > Pb Zn > Cu > Pb

pH = 2.4 [SO4 ~- ] = 12,000 mg 1-~ pH = 2.9

N u n n a n d Riches (1978)

Zn > Pb > Cu Zn > Cu > Pb

p H = 6.6 pH = 3.3

307 1 0 -3"49, 1 0 - s ' l ° and 10 -s's° M, respectively, in a solution which had a resulting pH of 7.0. These concentrations are close to the values suggested by our calculations, Zazn = 10 -3"2, Zapb = 10 -4.4 and Zacu = 10 -s'3 for a solution with fc02 = 10-2, ac1- = 10-1 and aso42- = 10-1; at pH 7. It is evident that further extensive experiments and comparisons of this type need to be done, with all solution parameters, particularly concentrations of chloride, sulphate and PCO2 being carefully monitored. However, the suggestion is strong that precipitation of secondary oxidized minerals from saturated solutions in the vicinity of weathering sulphides, may play an important part in controlling the availability of base-metal ions to, and their distribution in, groundwaters. This is not to say that the weathering of sulphides or the rate of release of metals from the sulphides is necessarily so controlled, but rather that secondary oxidized minerals may be important intermediates between release of base-metal ions from sulphides and their subsequent migration and dispersion via groundwater systems. These phases might be expected to be fine grained, well dispersed and difficult to detect, for example by X-ray diffraction. Additionally, the prospect of their preferential nucleation on, or adsorption to, surface active sites is very real. James and McNaughton (1977) have pointed out that heavy-metal uptake on inorganic minerals is strongly dependent on pH, that this dependence is similar to that for the formation of insoluble hydrolysis products and often occurs at similar, but slightly lower pH values than hydrolysis. Ratios of base-metal ions in solution so established, might be expected to persist in solution for some distance from the site of initial precipitation, until groundwater parameters altered sufficiently to preferentially remove one or other of the ions. Spain et al. (1964) demonstrated that the zoning of metal-sulphide precipitates in a sulphide-containing agar gel was dependent on diffusion. The distribution or order of precipitation away from the metal-ion source was based on solubility product equilibria, those precipitates with the least solubility being formed nearest the source. We have carried out similar experiments with gels containing bicarbonate, sulphate and/or chloride which show similar zoning effects. Furthermore, it is likely that the distribution of Cu, Pb and Zn around base-metal deposits is in part controlled by this type of mechanism. From the results for individual Cu, Pb and Zn solubilities, it is now po~ sible for any given solution, to define the minerals responsible for limiting the solubilities of Cu, Pb and Zn, and to define an "equilibrium" secondarymineral assemblage for these elements for that solution. Secondary-mineral assemblages for pH's in the range 3 < pH < 8 all with f c 0 2 = 10-2 bar, are shown in Table VIII. The carbonate assemblage malachite--cerussite--smithsonite is dominant for pH > 8.0 in any solution and in chloride- and sulphate-free solutions for pH > 5.0. For solutions with pH < 7 there is a wide range of possible secondary-mineral assemblages, depending on the exact conditions in solution.

308 TABLE VIII Mineral assemblages* in equilibrium with some solutions saturated with Cu, Pb and Zn at PCO 2 = 10 -5 pH

No chloride (acl = 0)

High chloride (acl = 1)

no sulphate (aso, 2- = 0)

sulphate (aso, ~- = 10 -1 )

no sulphate (aso~- = 0)

sulphate (aso42- = 10 -~)

3.0

> >

:> Ang

Ata Cot

> Ang

4.0

Mal Cer

Ant Ang

Ata Cot

Ata Ang

5.0

Mal Cer Smi

Bro Ang Smi

Ata Pho >

Ata Ang >

6.0

Mal Cer Smi

Bro Ang Smi

Ata Pho Smi

Ata Ang Smi

7.0

Mal Cer Smi

Mal Cer Smi

Ata Cer Smi

Ata Cer Smi

8.0

Mal Cer Smi

Mal Cer Smi

Mal Cer Smi

Mal Cer Smi

> = solid phase extremely soluble ( z a M > 1). *Ang = anglesite; Ant = antlerite; Ata -- atacamite; Bro -- brochantite; Cer = cerussite; Cot = cotunnite; Mal = malachite; Pho = phosgenite; Smi = smithsonite.

Assemblages with atacamite or brochantite, with anglesite or phosgenite, and s m i t h s o n i t e a r e all p o s s i b l e in t h e p H r a n g e o f 5 . 0 - - 7 . 0 . C o t u n n i t e is p r e c i p i t a t e d f r o m s o l u t i o n s w i t h h i g h c h l o r i d e a n d l i t t l e s u l p h a t e a t p H < 5.0. T h e assemblages are reasonably specific to a certain range of solution parameters, a n d w h e r e r e c o g n i t i o n o f t h e s e a s s e m b l a g e s in g o s s a n s a n d / o r s o i l s is p o s s i b l e , they could prove to be a useful diagnostic tool for evaluating the conditions under which the gossan or soil anomaly was formed. CONCLUSIONS (1) Activity calculations indicate that cerussite, anglesite, laurionite and cotunnite are the phases which limit the solubility of Pb at various conditions o f p H , f c o 2 , ac1- a n d a s o ~-. ( 2 ) A n i n c r e a s e in fCO~ t o 1 0 - : b a r d o e s n o t s i g n i f i c a n t l y a f f e c t t h e s o l u b i l -

309 ity of Pb; phosgenite replaces laurionite as the phase which limits Pb solubility in the pH range of 5--6.5. (3) Zinc hydroxide and Zn4(SO~)(OH)6 limit the solubility of Zn in sulphate solutions with fco~ = 10-2 bar. (4) At fco2 = 10-2 bar, smithsonite, ZnCO3, replaces Zn(OH)2 as the least soluble Zn phase, and the solubility of Zn decreases significantly at all pH's from the values at f e o ~ = 10-3"s bar. (5) Sulphate solutions with pH < 6.0 have an order of solubility Zn > Cu > Pb and solutions with 6.0 < pH < 8.3 have an order of solubility Zn > Pb > Cu. These orders of solubility are in agreement with some observed orders of mobility for base metals near some base-metal sulphide deposits. In this treatment of solution geochemistry of Pb and Zn, and in the assessment of relative mobilities of Cu, Pb and Zn, many complicating and possibly important factors have been ignored. Thornber and Wildman (1979), point out that Cu, Pb and Zn (as well as Co and Ni) can coprecipitate with goethite. Adsorption of base metals onto clays and hydrous oxides of Fe and Mn (Jenne, 1968; James and McNaughton, 1977) is a potentially important factor which is likely to influence solubilities particularly in neutral and alkaline solutions. However, in the immediate vicinity of many deposits we would suggest that a zone exists in which the adsorption capacity of clays and hydrous oxides has been exceeded and within which direct precipitation of secondary minerals is possible. We have not attempted in these calculations, to take account of anions other than carbonate, sulphate and chloride, e.g. phosphate, which are capable of producing relatively insoluble precipitates with many metals. Likewise, interference from and compound formation with, other cations has not been included. Complexing of metals, particularly of Cu, with humic acids, can also be important (Mantoura et al., 1978). The results of the calculations performed here should therefore be used guardedly, and only in systems which appear to be simple, inorganic, well "equilibrated", and where the adsorption capacities of oxides and clays are likely to be exceeded. Activity coefficients need to be considered if total activities are to be compared with observed solubilities. With these important provisos and qualifications, calculated solubilities of Cu, Pb and Zn may contribute to a better understanding of the processes involved in the formation of base-metal deposits and particularly to their subsequent dissolution. ACKNOWLEDGEMENTS The authors would like to thank W.E. Ewers and E.H. Nickel for their helpful advice and comment. Thanks are also due to Mrs. C. Harris for preparation of the manuscript and C.R. Steel for diagrams.

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