- Email: [email protected]
The solubility of calcite and aragonite in sulfate-free seawater and the seeded growth kinetics and composition of the precipitates at 25°C
Chemical Geology, 74 (1989) 309-320 Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands
309
[3]
THE SOLUBILITY OF CALCITE AND ARAGONITE IN SULFATEFREE SEAWATER AND THE SEEDED GROWTH KINETICS AND COMPOSITION OF THE PRECIPITATES AT 25 °C ALFONSO MUCCI, RENI~ CANUEL and SHAOJUN ZHONG Department of Geological Sciences, McGiU University, Montreal, Que. H3A 2A7 (Canada) (Received January 22, 1988; revised and accepted September 6, 1988)
Abstract Mucci, A., Canuel, R. and Zhong, S., 1989. The solubility of calcite and aragonite in sulfate-free seawater and the seeded growth kinetics and composition of the precipitates at 25 ° C. Chem. Geol., 74: 309-320. The influence of dissolved sulfate ions on the precipitation kinetics and composition of calcite and aragonite overgrowths was investigated in artificial seawater solutions at 25 ° C using a constant disequilibrium technique. The solubility of both minerals in sulfate-free seawater was determined to allow the expression of the saturation state of precipitating solutions as the ratio of the CaCO.~ ion concentration product to their stoichiometric solubility constant. SO~- ions inhibit both calcite and aragonite precipitation. The inhibition in seawater is not selective and therefore, whether in the presence or absence of 0.028 m S042- , aragonite precipitates more rapidly on aragonite seeds than calcite does on calcite seeds from supersaturated solutions. SrC03 incorporation in aragonite is independent of the precipitation rate and unaffected by S042- ions in seawater. The distribution coefficient of Sr 2+ in aragonite overgrowths precipitated from seawater is 0.97 _+0.08 while a value of 1.05 ± 0.09 is obtained in sulfate-free seawater. In contrast, M g C Q incorporation in calcite overgrowths is influenced significantly by the presence or absence of SO~ions in seawater. While the distribution coefficient of Mg 2+ in calcite is also independent of the precipitation rate, its value increases from 0.0172 ± 0.0022 in seawater to 0.0230 ± 0.0026 in sulfate-free seawater, corresponding, respectively, to the precipitation of 8.1 ± 1 and 10.5 ± 1 mole% magnesian calcites.
1. I n t r o d u c t i o n
It has often been reported that sulfate ions in solution inhibit both calcite dissolution and precipitation rates (Reddy and Nancollas, 1976; SjSberg, 1978; Busenberg and Plummer, 1985; Walter, 1986). In addition, several investigators (Kitano, 1962; Kinsman and Holland, 1969; Kitano et al., 1975; Walter, 1986) have indicated that the presence of sulfate ions favored the formation of aragonitic CaCOs rather than calcite, while the work of Bischoff and Fyfe
0009-2541/89/$03.50
(1968) demonstrated the inhibition of aragonite-calcite transformation by dissolved S O l - . Seawater contains a relatively high concentration of sulfate ions ( ~ 28 mM at a salinity of 35). The presence of SO~- and Mg 2+ ions is believed to be responsible for the preferential nucleation of aragonite rather than the more stable calcite in highly supersaturated seawater solutions. However, during the diagenesis of organic-rich sediments and following oxic and suboxic diagenesis, dissolved sulfate is used by bacteria as an alternate electron acceptor in the process of organic matter oxidation. Under the
© 1989 Elsevier Science Publishers B.V.
310
appropriate conditions (i.e. closed or restricted sedimentary environment) this metabolic reaction may result in the depletion of SO~- ions in the sediment pore waters. In addition, sulfate reduction is accompanied by the production of alkalinity which often leads to an increase in the pore-water saturation state with respect to CaCO3 and the precipitation of carbonate minerals. Calcite rather than aragonite is most often found as an authigenic phase in these sediments (Bathurst, 1975). Marine waters with low S042- concentrations are also believed (Baker and Kastner, 1981; Kastner, 1984; Baker and Burns, 1985) to be conducive to dolomite formation, but this idea has not yet received widespread recognition as a possible dolomitization mechanism (Morrow and Ricketts, 1986; Hardie, 1987). The purpose of this paper is to provide quantitative data on the growth kinetics of both calcite and aragonite in "normal" and sulfate-free seawater and determine the composition of the calcium carbonate overgrowths precipitated from these solutions at 25 ° C.
2. M a t e r i a l s a n d m e t h o d s
2.1. Solids, solutions and experimental procedures Baker ® "Instra-analyzed flux reagent" grade calcium carbonate was used for all experiments in which calcite was studied, both for the solubility measurements and as a seed material for the precipitation reactions. Procedures for cleaning and size separating this material have been described previously (Mucci, 1986). Aragonite synthesized in the laboratory by the method of Wray and Daniels (1957) as modified by Katz et al. (1972) at a temperature of 70 ° C was used without further treatment for all experiments in which aragonite was studied. Xray diffraction spectrometry (XRD) and scanning electron microscopy examination (SEM) of this material indicate the absence of vaterite and the presence of < 1 wt.% calcite. The cal-
cite and aragonite have surface areas of 0.52 and 3.40 m 2 g- 1, respectively, as determined by the Kr-BET method of de Kanel and Morse (1979). Aged artificial seawater was used for all the experiments. The artificial seawater of salinity 35 was prepared to include all major elements of natural seawater including F - according to the method of Kester et al. (1967), modified to fit the analysis of Millero (1974). The sulfatefree seawater was prepared according to the same procedure but excluding sulfate. The ionic strength of the solution was balanced to equal that of a standard seawater (It=0.697 m) by increasing the amount of NaC1 added to the preparation. The seawater solution was stored in a polypropylene container for two months to reduce the dissolved phosphate concentration ( < 1/IM) prior to use. Solubility measurements were carried out in a closed system at 25°C over long periods of equilibration by a procedure described at length by Morse et al. ( 1980 ) and Mucci ( 1983 ). Equilibration was followed from both undersaturated and supersaturated solutions. Precipitation reactions were conducted in an open system in solutions of close to constant composition over a wide range of precipitation rates. Constancy of composition was maintained during the length of the precipitation by the use of a chemo-stat technique through the simultaneous injection of two titrants in equal amounts by a dual syringe pump. The mixture of the two titrants reproduced the exact composition of the precipitating solution plus an excess in calcium, strontium and carbonate alkalinity to compensate for the calcium carbonate precipitation. The temperature of the precipitating solution was held constant at 25 (+0.05)°C by circulating water through a jacketed 400-ml glass reaction vessel from a constant-temperature bath. The Pco~ of the solution was kept nearly constant at ~ 3000 ppm or 10 -2.5 atm. by bubbling a CO~-N2 gas mixture of known composition. A detailed description of the chemo-stat, its working concept and titrant composition have been presented pre-
311
viously (Mucci and Morse, 1983; Mucci, 1986). The precipitation was carried out on ~ 0.20.7 g of seed material in 300-400 ml of seawater solution until > 10 -4 mol of carbonate was precipitated. The amount of carbonate precipitated was calculated from the weight of the standardized titrant used during the length of the experiment and the carbonate alkalinity of the solution before and after the precipitation reaction (according to equation 1 in Mucci, 1986).
2.2. Calculation o[ the stoichiometric solubility constants The stoichiometric solubility constant of calcite (or aragonite) in an electrolyte solution is defined as: K c* (or a) = [Ca 2+1 [COW-]
(1)
where [i] is the equilibrium total ion concentration. While [Ca 2+ ] was measured directly by titration, [CO~- ] was derived from the carbonate alkalinity, the equilibrium pH of the solution and the carbonic acid second apparent dissociation constant. The carbonate alkalinity was obtained from the titration alkalinity from which the boric acid contribution was subtracted. Analytical procedures for the determination of Ca 2+ and titration alkalinity have been described previously (Mucci, 1986). The equilibrium pH was measured using a combination glass electrode which had been previously calibrated against U.S. National Bureau of Standards (N.B.S.) buffer solutions (6.862, 7.413 and 9.180). The first apparent dissociation constant of boric acid and the second apparent dissociation constant of carbonic acid in sulfate-free seawater were derived from their standard seawater values (salinity 35, 25 °C) as they appear in Millero ( 1979 ) and the ratio of the total ion activity coefficients of the appropriate ions in solution. Thus,
K~(sf) =Ki3tsw) 7(B(OH)~-)sw ~(B(OH)4)si
=2.16-10 -9
(2)
and
Y(co,~ - )~ Y(HCO:r )~w
=8.61"10 -1°
(3)
where Ki~ and K.~ are the first apparent dissociation constant of boric acid and second apparent dissociation constant of carbonic acid, respectively, in seawater (sw) and sulfate-free seawater (sf); and 7 ( f s are the total ion activity coefficients. Total ion activity coefficient estimates were calculated from the ion pairing model of Millero and Schreiber (1982).
2.3. Steady-state solution composition and saturation state of the precipitating solution The steady-state ion concentration product, [Ca 2+ ] [CO~- ], of the precipitating solution was calculated using the steady-state pH measured during the experiment, the carbonate alkalinity and the Ca 2+ concentration of an aliquot withdrawn at the end of the precipitation. For most measurements the steady-state saturation of the solution was reached within 15 min. and remained constant for the length of the experiment. The steady-state pH was monitored by two sets of electrodes which were calibrated against N.B.S. buffer solutions. The stoichiometric solubility constants of calcite (or aragonite) determined in this study in sulfatefree seawater and by Mucci ( 1983 ) in "normal" seawater were used to calculate the saturation state of the precipitating solution with respect to calcite (or aragonite), according to: ~2c(or a)
--
[Ca 2+1 [COW-] /~c(or a)
(4)
Uncertainties associated with estimates of the saturation state, which might have been introduced by u s i n g the N.B.S. buffer calibration
312
system in strong electrolyte solutions (e.g., liquid junction potentials, etc.), are minimized since b o t h / ~ c (or a) and the steady-state C a C Q ion concentration product of the precipitating solution were determined using the same electrode and p H scale. The Sr 2+ concentration of the solutions was measured before and after the reaction by flame atomic absorption spectrophotometry (AAS) using aqueous standards in a seawater matrix for calibration. Precision of the AAS analysis is estimated to be better than + 5%.
2.4. Overgrowth composition The Mg content of the calcite overgrowths and the Sr content of the aragonite overgrowths were determined following the acid digestion of a known amount of reacted solid by AAS. The mole fraction of MgC03 and SrCO3 in the overgrowth was calculated from the results of the AAS analyses, the amount of carbonate precipitated and the amount of material dissolved for analysis (Mucci and Morse, 1984 ). No correction was introduced to compensate for the presence of residual solution salts as the overgrowths were rinsed with calcite-equilibrated distilled water after being filtered out of the parent solution (Mucci, 1986). The reacted solid was also examined by X R D spectrometry to determine its composition and mineralogy and identify other carbonate mineral phases which might have precipitated along with the calcite and aragonite. Powder packs were prepared and irradiated using a Siemens ® model D-500 X-ray diffractometer. The Cu-K~ wavelength radiation was used as a source and the diffraction spectra was recorded using a proportional counter detector. 3. R e s u l t s
3.1. Solubility of calcite and aragonite in sulfate-free seawater Analytical results and calculated CaCO3 ion concentration products of sulfate-free seawater
solutions equilibrated with calcite and aragonite are presented in Tables I and II. The C a C Q ion concentration products measured after 188 days of equilibration differ by no > + 7% and their averages are taken to represent the stoichiometric solubility constants of calcite and aragonite in sulfate-free seawater at 25°C and 1-atm. total pressure. This is also taken to be an indication that the solubilities are reversible since measurements were carried out from both undersaturated and supersaturated solutions. Stoichiometric solubility constant values are 4.07+0.12 and 5.94+0.08 for calcite and aragonite, respectively. These values were used in eq. 4 to calculate the saturation state of the precipitating solutions in the overgrowth experiments.
3.2. Influence of dissolved sulfate on the precipitation kinetics of calcite and aragonite in seawater Selected parameters and results from single precipitation reactions are presented in Tables III-V. The precipitation rate data were fitted to the following empirical kinetic equation: R = k ( ~ c ( .... } - - 1 ) n
(5)
or to its logarithmic form, l o g R = l o g k + n l o g ( t 2 c ( .... I - 1 )
(6)
where R is the precipitation rate normalized to the reactive surface area (mol hr.-~ m - 2 ) ; k is the rate constant; n is the empirical reaction order; and (t2~(or at - 1 ) is the degree of supersaturation with respect to calcite or aragonite. The precipitation rate data for calcite and aragonite in both "normal" and sulfate-free seawater are plotted in Fig. l a and b as a function of the logarithm of their respective supersaturations. The kinetic parameters n (order of reaction ) and log k (rate constant) for each phase in both solutions are presented in Table VI. As was expected from results of previous studies, the precipitation rate of calcite in sulfate-bearing or "normal" seawater is slower
313 TABLE I
Calcite stoichiometric solubility constant in sulfate-free artificial seawater, It = 0,697 m at 25 °C Approach
Equilibration period
[Ca 2+ ]
At (meq kg- ~)
PH T
(mmol k g - ~)
M
(107 mol ~ kg ~)
(days) UND SUP
2 2
10.23 10.19
0.844 1.670
7.800 7.552
3.97 4.83
UND SUP
13 13
10.13 9.87
0.891 1.849
7.773 7.531
3.95 4.96
UND SUP
188 188
10.11 9.74
0.923 1.349
7.749 7.613
3.90 4.22
UND SUP
231 231
10.13 9.72
0.930 1.224
7.753 7.654
3.97 4.15
UND SUP
248 248
10.24 9.82
0.963 1.233
7.730 7.648
3.98 4.17
UND = approach to equilibrium from undersaturation; S U P = approach to equilibrium from supersaturation. TABLE II
Aragonite stoichiometric solubility constant in sulfate-free artificial seawater, I, = 0.697 m at 25 c C Approach
Equilibration period
[Ca `-'+ ] (mmol kg- ~)
A, (meq kg ')
pH Na~s
K: (107 mol e kg e)
(days) UND SUP
2 2
9.83 9.67
1.068 1.866
7.930 7.660
6.28 6.44
UND SUP
13 13
9.68 9.54
1.087 1.654
7.918 7.650
6.16 5.50
UND SUP
188 188
9.77 9.29
1.158 1.671
7.850 7.696
5.84 5.96
UND SUP
231 231
9.78 9.23
1.172 1.533
7,853 7,733
5.95 5.85
U ND SUP
248 248
9.87 9.41
1.206 1.496
7,839 7,750
6.02 6.02
For explanation of UND and SUP see footnote to Table I.
than in sulfate-free seawater at all supersaturations investigated. The rates differ by a factor of 2 (at f2c=3.5) to 3 (at t2c=17). Extrapolation of the least-squares fit lines to lower saturations indicates that the precipitation rate of calcite in both solutions would be identical at a f2c = 1.4. A similar pattern is observed for aragonite in Fig. lb. Above a saturation state, f2a, of 1.7 the precipitation rate of aragonite in sulfate-free seawater becomes faster than in "nor-
mal" seawater. Rates differ by a factor of 4 at ~a = 4. The rates of calcite and aragonite precipitation were also plotted as a function of the supersaturation with respect to calcite in "normal" (Fig. 2a) and sulfate-free (Fig. 2b) seawater to allow comparison of the relative precipitation rates of the two phases in the same solution at identical CaCO:~ ion concentration products. In both supersaturated solutions ara-
314 T A B L E III Kinetic data and composition of aragonite overgrowths precipitated from artificial seawater at 25 : C Experiment No.
[Sr 2+ ] (ppm)
[Ca 2+ ] (mmol kg -1)
pH TM
~l~ a
Amount of carbonate precipitated (104mol)
-log(rate) (mol hr. ' m -2)
SrCO:~ (mole%)
Zl Z2 Z3 Z4 Z5 Z6 Z7 Z8 Z9 ZIO Zll Z12 Z13 Z15
6.7 6.7 6.6 6.7 7.0 6.9 7.0 6.7 --5.1 7.5 8.3 8.1
8.83 8.84 8.84 9.24 9.67 9.87 10.08 9.21 9.66 9.62 9.27 9.75 9.96 9.81
7.896 7.829 7.737 7.714 7.852 7.637 7.718 7.785 7.787 7.873 7.937 7.653 7.672 7.595
3.77 2.87 2.33 -3.38 1.44 -2.84 2,14 3.93 4.95 1.64 1.65 1.19
15.6 15.4 15.6 14.5 11.4 12.0 11.1 14.8 4.26 12.6 14.3 3.18 3.22 0.841
4.010 4.290 4.584 -4.009 4.950 -4.318 4.539 3.703 3.437 4.776 5.066 5.588
0.88 0.75 0.81 0.80 0.73 -0.83 0.68 --0.63 ----
D~")
1.05
0.89 0.98 0.99 0.92 1.08
0.84
1.03
*Calculated from the AAS analysis and the [Sr 2+ ] / [ C a 2. ] ratio of the final solution.
TABLE IV Kinetic data and composition of calcite overgrowths precipitated from sulfate-free artificial seawater at 2 5 C Experiment No.
R1 R2 R3 R4 R6 R7 R8 R9 R12 R13 RI4 RI5 RI6 R17 R18
[Ca ~+] (mmol kg 1)
10.47 10.31 10.75 10.98 11.00 11.10 10.84 10.70 10.93 10.76 10.52 10.52 10.57 10.38 9.56
pH TM
7.934 7.876 7.869 7.887 8.055 7.772 7.990 7.805 7.848 7.704 7.854 7.857 8.032 7.767 7.645
£2,.
12.25 9.54 10.01 9.96 13.63 6.60 12.52 7.43 9.02 4.73 -9.62 14.9 5.51 3.43
Amount of carbonate precipitated (104 mol)
-log(rate) (mol hr. -I m -~)
10.5 11.7 9.28 9.70 9.53 8.68 9.83 10.4 9.59 9.51 11.4 11.2 9.05 24.8 22.3
2.964 3.242 3.245 3.100 2.806 3.902 2.957 3.666 3.390 4.023 -3.251 2.664 4.321 4.907
*Calculated from the AAS analysis and the [Mg 2+ ] / [ C a `-,+] ratio of the final solution.
MgCO:~ (mole%) ( × 1 0 ~) AAS
XRD
9.6 10.1 10.2 8.7
-9.0 ---
2.11 2.20 2.31 1.99
11.7 9.7
11.0 9.0
2.80 2.21
--
12.0
--
11.8 11.2 10.3 9.0 10.0
-11.0 9.5 --
2.74 2.52 2.30 1.98 2.18
--
315
TABLE V Kinetic data and composition of aragonite overgrowths precipitated from sulfate-free artificial seawater at 25 °C Experiment No.
[ Sr~+ ] (ppm)
[Ca 2+ ] (mmol kg -1)
PH
R28 R29
8.0 7.7 7.9 8.5 8.5 8.5 8.2 8.7 8.2 8.2 8.0 8.2 8.0 8.1 7.7 7.8
9.79 9.73 9.60 10.87 10.90 10.32 10.13 10.53 10.05 10.39 10.33 10.33 10.16 10.18 9.96 10.24
7.699 7.697 7.683 7.531 7.554 7.665 7.644 7.617 7.629 7.580 7.539 7.660 7.672 7.705 7.562 7.562
R30 R31 R32 R33 R34 R35 R36 R37 R38 R39
R40 R4I R43 R44
-2 O0
i (a)
I
T M
~a
Amount of carbonate precipitated ( 104 mol )
-log(rate) (molhr.-~ rn 2)
SrCQ (mole%)
D~r
3.20 2.76 2.70 -1.82 2.82 2.59 2.25 2.33 1.97 -2.85 3.04 3.44 1.73 1.77
12.4 12.8 12.1 8.25 8.50 10.5 10.3 8.59 10.9 9.39 9.16 10.1 10.3 10.5 10.9 9.47
3.724 ?1.022 4.029 -4.597 3.848 4.012 4.373 4.105 4.513 -3.736 3.638 3.529 4.970 4.824
0.91 0.91 0.93 1.05 1.01 0.89 -0.93 0.82 0.90 1.02 I).85
1.00 1.03 1.02 1.21 1.17 0.97 -1.01 0.91 1.04 1.19 0.96
-
-
,
I
i
-
-0.93 --
,
~
t
-
i
,
o 'E
5
/ j
J
j
g
o
~ o
o!~o
'
o!8o
,'.oo
•
-4 O0
x-
g
r
!
1
E
o 4o
-
-1.08 --
,!2o
log ( Q c - I )
-500
-o'7o
o! . . . .
!. . . .
11o o!1o I o 9 ( f ~ a - I)
o!3o
o!5o
Fig. 1. Precipitation kinetics of calcite (a) and aragonite (b) overgrowths in "normal" and sulfate-free seawater at 25 ~C as a function of their respective supersaturations; . . . . calcite in seawater (from Mucci, 1986); A =calcite in sulfate-free seawater; • = aragonite in seawater; o = aragonite in sulfate-free seawater.
gonite precipitates more rapidly on aragonite seeds than calcite does on calcite seeds.
3.3. Influence of dissolved sulfate ions on the composition of calcite and aragonite overgrowths precipitated from seawater Results of the X R D analysis of calcite and aragonite overgrowths precipitated from sulfate-free seawater confirmed that the mineral-
ogy of the overgrowths is controlled by the nature of the seed material. Only magnesian calcite overgrowths could be detected when calcite seeds were used while the presence of no other phase could be detected on the reacted aragonite seed material. When it could clearly be resolved from the calcite seed material the position of the magnesian calcite overgrowth peak was used to determine its MgCO:, content using the data of Goldsmith et al. (1961) as
316 T A B L E VI Kinetic parameters for calcite and aragonite precipitation rates in seawater and sulfate-free seawater at 25 ~C, expressed according to the empirical rate law described in eq. 6 Mineral
Solution
n
- log k
Source
Calcite
seawater
2.8
6.29
1.9
5.41
Mucci and Morse ( 1983 ); Mucci ( 1986, 1988 ) Burton and Walter (1987
SOs-free seawater
3.0
6.08
this study
seawater
1.6 1.7
4.63 4.39
this study Burton and Waiter (1987
S04-free seawater
2.2
4.43
this study
Aragonite
I (a) ,~;/
//
L o
-6.oo
g
~ / 7 '
oo
-400 /
/
jj
J
f
-6 O0
-]
E
E
-5 oo
J o!oo
'
oL2o
'
log
I
046 ('~C-
060
. I. . . .
0!80
!,o
o6o' log
I}
(~'c-
o!6o
,!oo
I)
Fig. 2. Precipitation kinetics of calcite and aragonite overgrowths in "normal" (a) and sulfate-free (b) seawater at 25 C as a function of the supersaturation with respect to calcite. Symbols are as in Fig. 1.
presented by Milliman et al. (1971). These results are presented in Table IV. Since the calcite and aragonite overgrowths were precipitated from solutions of nearly constant composition despite the ongoing precipitation, the partitioning of a trace component, X 2+ (e.g., X 2 + = M g 2 + , Sr2+), between the CaCO:~ solid and the solution can be defined by the H e n d e r s o n - K r a c e k (Henderson and Kracek, 1927) or homogeneous distribution coefficient: c
c
Mx2./Mca2+ D ~ + -- Mx~'+M / ca~+L L
(7)
where Mx2+ and Mca2+ denote the molar concentrations of X 2+ and Ca 2+, respectively, in the C a C Q overgrowth (c) and the parent solution (L). The distribution coefficients measured in this
study are independent of the precipitation rate. The averaged distribution coefficient of Mg 2+ in calcite overgrowths precipitated from sulfate-free seawater is 0.0230_+0.0026, corresponding to the precipitation of a 10.5_+ 1 mole% magnesian calcite. This value is significantly higher than the one obtained in "normal" seawater (0.0172 _+0.0022 or 8.1 + 1 mole% MgCO~) and cannot be explained by the small variation (i.e. < -+ 1% ) in the ratio of the activity of Mg ~+ to Ca 2+ ions between the two solutions. Processes occurring at or near the surface of the growing crystal are probably responsible for the discrepancy. This hypothesis will be expanded upon further in Section 4.3. On the other hand, the averaged distribution coefficients of Sr "'+ in aragonite overgrowths precipitated from seawater (0.97_+0.08) and sulfate-free seawater ( 1.05 _+0.09 ) are statisti-
317
cally identical. This is consistent with the fact that the ratio of the ion activity of Sr 2+ to Ca s + in the parent solutions are the same ( < +_ 1% ). The values of D ~ . determined in this study are in excellent agreement with the results of Kinsman and Holland ( 1969 ) and Kitano et al. (1971). Kinsman and Holland (1969) measured the distribution coefficient of Sr s + in aragonite precipitated from seawater between 16 ° and 96~C. A value of 1.13 + 0.04 at 25°C can be interpolated from the linear relationship they observed between the distribution coefficient and temperature. Kitano et al. (1971) determined a distribution coefficient of 1.1-2_ 0.1 following the precipitation of aragonite from a Ca (HCO:~) 2-MgC12-SrCls solution at 20 ° C. 4. D i s c u s s i o n
4. I. Solubility of calcite and aragonite in sulfate-free seawater Calcite and aragonite do not precipitate as pure C a C Q from seawater and yet, the calcite and aragonite solubility behavior in seawater closely follows that predicted from thermodynamics of pure calcium carbonate (Plummer and Sundquist, 1982; Mucci, 1983). The same conclusion can be shown to apply to their solubility behavior in sulfate-free seawater by comparing the C a C Q ion activity coefficient product in sulfate-free seawater estimated from model calculations with the value derived from solubility measurements. The C a C Q ion activity coefficient product can be calculated from the ratio of the thermodynamic to the stoichiometric solubility constants according to: 0 t ~r 7~c~,~*~)'(co~ ~= (Ksp/K~p)
02
(8)
where K °" is the thermodynamic solubility constant of calcite or aragonite at 25 °C and 1-atm. total pressure (pK(,! = 8.48 _+0.02, p K ° = 8.34 ± 0.03; from P l u m m e r and Busenberg, 1982; Mucci, 1983; Sass et al., 1983); and 0 is a correction factor used to convert the total ion ac-
tivity coefficient product from kg 2 sw m o l - " to kg 2 H 2 0 mo1-2 (0=0.964). CaCO3 ion activity coefficient product values calculated from eq. 8 and the calcite and aragonite solubility constants (0.0076 ± 0.0006 and 0.0072 +_0.0006, respectively) compare thvorably with a value of 0.0065 _+ 10% estimated using the ion pairing model of Millero and Schreiber (1982).
4.2. Influence of dissolved sulfate ions on the relative rates of calcite and aragonite precipitation from seawater Based on the results described in Section 3.2 it appears that a decrease of the sulfate content of seawater following sulfate reduction in anoxic sediments would not result in a change in the mineralogy of authigenic C a C Q . These results apparently contrast with conclusions reached by previous investigators (Kitano, 1962; Kitano et al., 1975; Walter, 1986) who noted that dissolved SO~ ions favor aragonite precipitation by selectively inhibiting calcite precipitation, thus leading to a reversal in the relative precipitation rate of the two minerals following addition of SO4- to the precipitating solution. However, the high [Mg ~+ ] / [ C a ~+ ] or [Mg 2+ ] in seawater is responsible for inhibiting calcite precipitation even in the absence of SO4-, effectively shadowing its influence on the relative precipitation rate of the two minerals. Calcite and aragonite inhibition by SO4- is likely due to a combination of' its adsorption on the surface and incorporation in the mineral structure. Sulfate adsorption at Ca ~+ sites and partial binding of CaSO ° ion pairs on the surface of calcite and aragonite require that the CaS04 bond be broken and SO~- be desorbed from the surface before growth can occur. Furthermore, several workers (Kitano et al., 1975; Takano et al., 1980; Busenberg and Plummer, 1985; Takano, 1985) have shown that SO4 ions are more easily incorporated in calcite than aragonite. In tact, Busenberg and Plummer
318 (1985) observed that the incorporation of SO~-- in synthetic calcites increased with the precipitation rate and sulfate concentration, and resulted in an increase in the solubility and unit cell dimensions of the sulfate-bearing calcites. Processes accompanying sulfate reduction in marine sediments, such as the release and accumulation of dissolved phosphate (Walter, 1986) and organic matter decomposition products (Berner et al. 1978) and an increase in the Pco~ in the interstitial waters (Reddy et al., 1981; Given and Wilkinson, 1985) may influence the mineralogy of marine carbonate cements. In fact, Walter (1986) observed that phosphate ions are selective in inhibiting aragonite precipitation from seawater. Temperature (Burton and Walter, 1987) and substrate mineralogy also certainly play important roles
reaction and may behave differently on the surface of the mineral. The incorporation of sulfate in calcite may help to relieve or accentuate lattice strain induced by the incorporation of Mg 2. in magnesian calcites. Mg 2+ ions (0.65 ~,) are smaller than Ca 2+ ions (0.99 A) while S O l - ions ( ~ 1.19 ,~3) are significantly larger than CO~ions (~0.63 A~). Whether or not this mechanism could buffer Mg 2+ incorporation in calcite precipitated from seawater is uncertain but warrants further investigation. Takano (1985) claimed that S O l - ions in inorganically precipitated calcite occupy lattice sites replacing carbonates and that high-Mg calcites can incorporate more sulfate than if Mg-free.
4.3. Influence of sulfate ions on the composition of calcite overgrowths precipitated from seawater
Based on a comparison of calculated and estimated C a C Q total ion activity coefficients it appears that the solubility or calcite and aragonite in sulfate-free seawater closely follows the predicted solubility of pure CaCO:~. Quantitative seeded precipitation kinetic data of calcite and aragonite in "normal" and sulfate-free seawater indicate that dissolved SO42- ions inhibit the precipitation of both minerals in seawater. The inhibition is not selective and could not be responsible for a reversal in the mineralogy of authigenic CaCO:~ precipitated following a decrease in S O l - concentration in sediment pore waters as a result of microbiologically-mediated sulfate reduction. SrCO., incorporation in aragonite overgrowths is unaffected by the removal of SO] ions from seawater. The distribution coefficient of Sr 2+ in aragonite precipitated from seawater is independent of the precipitation rate and approximately equal to 1. In contrast, MgCO3 incorporation in calcite is influenced significantly by the presence or absence of SO~- ions. The distribution coefficient of Mg 2+ in calcite precipitated from sulfate-free seawater is 34% larger than in "normal" seawater.
One of the more interesting findings of this study is the strong influence which dissolved SO]- ions appear to exert on the amount of Mg 2+ incorporated in magnesian calcite overgrowths. In the absence of SO42- ions the distribution coefficient of Mg 2+ in calcite precipitated from seawater increases by 34%. This result has significant implications with respect to the genesis of inorganic high-Mg calcites and dolomite in marine sedimentary environments. However, as the activity ratio of Mg 2+ and Ca 2+ in the two solutions does not vary by more than + 1% we can only speculate on the processes responsible for the increased incorporation of Mg 2+ in calcite in the absence of sulfate. Variation of the distribution coefficient can only be explained partially in terms of an increase ( ~ 4 . 4 % ) in the concentration of MgCOI~ ion pairs relative to CaCO ° between the two solutions. These species may be important intermediate "reactants" in the precipitation
5. Conclusions
319
Acknowledgements The authors wish to express their gratitude to Marco Ciarlo for his technical assistance during most stages of this study. Financial support for this study was provided by the Natural Sciences and Engineering Research Council of Canada through grants Nos. U0432 and E1446. This project also benefited from the award of an equipment grant by the Faculty of Graduate Studies and Research, McGill University. Finally, one of us (S.Z.) wishes to acknowledge the financial support of the Government of the People's Republic of China. References Baker, P.A. and Burns, S.J., 1985. Occurrence and formation of dolomite in organic-rich continental margin sediments. Am. Assoc. Pet. Geol. Bull., 69: 1917-1930. Baker, P.A. and Kastner, M., 1981. Constraints on the formation of dolomite. Science, 213: 214-216. Bathurst, R.G.C., 1975. Carbonate Sediments and Their Diagenesis. Elsevier, Amsterdam, 2nd ed., 658 pp. Berner, R.A., Westrich, J.T., Graber, R., Smith, J. and Martens, C.S., 1978. Inhibition of aragonite precipitation from supersaturated seawater. Am. J. Sci., 278: 816837. Bischoff, J.L. and Fyfe, W.S., 1968. Catalysis, inhibition, and the calcite-aragonite problem, I. The aragonitecalcite transformation. Am. J. Sci., 266: 65-79. Burton, E.A. and Walter, L.M., 1987. Relative precipitation rates of aragonite and Mg calcite from seawater: Temperature or carbonate ion control? Geology, 15:111114. Busenberg, E. and Plummer, L.N., 1985. Kinetic and thermodynamic factors controlling the distribution of SO~- and Na + in calcites and aragonites. Geochim. Cosmochim. Acta, 49: 713-725. de Kanel, J. and Morse, J.W., 1979. A simple technique for surface area determination. J. Phys. E, Sci. Instrum., 12: 272-273. Given, R.K. and Wilkinson, B.H., 1985. Kinetic control of morphology, composition, and mineralogy of abiotic sedimentary carbonates. J. Sediment. Petrol., 55: 109119. Goldsmith, J.R., Graf, D.L. and Heard, H.C., 1961. Lattice constants of the calcium-magnesium carbonates. Am. Mineral., 46: 453-457. Hardie, L.A., 1987. Dolomitization: A critical view of some current views. J. Sediment. Petrol., 57: 166-183.
Henderson, L.M. and Kracek, F.C., 1927. The fractional precipitation of barium and radium chromates. J. Am. Chem. Soc., 49: 739-749. Kastner, M., 1984. Control of dolomite formation. Nature (London), 310: 410-411. Katz, A., Sass, E., Starinsky, A. and Holland, H.D., 1972. Strontium behavior in the aragonite-calcite transformation: an experimental study at 40-98 °C. Geochim. Cosmochim. Acta, 36: 481-496. Kester, D.R., Duedall, J.W., Connors, D.N. and Pytkowicz, R.M., 1967. Preparation of artificial seawater. Limnol. Oceanogr., 12: 176-179. Kinsman, D.J.J. and Holland, H.D., 1969. The co-precipitation of cations with CaCO:,, IV. The co-precipitation of Sr 2+ with aragonite between 16 ° and 96 ° C. Geochim. Cosmochim. Acta, 33: 1-17. Kitano, Y., 1962. The behavior of various inorganic ions in the separation of calcium carbonate from a bicarbonate solution. Bull. Chem. Soc. Jpn., 35: 1973-1980. Kitano, Y., Kanamori, N. and Oomori, T., 1971. Measurements of distribution coefficients of strontium and barium between carbonate precipitate and solution - Abnormally high values of distribution coefficients measured at early stages of carbonate formation. Geochem. J., 4:183 206. Kitano, Y., Okumura, M. and Idogaki, M., 1975. Incorporation of sodium, chloride and sulfate with calcium carbonate. Geochem. J., 9:75 84. Millero, F.J., 1974. The physical chemistry of seawater. Annu. Rev. Earth Planet. Sci. Lett., 2:101 150. Millero, F.J., 1979. The thermodynamics of the carbonate system in seawater. Geochim. Cosmochim. Acta, 43: 1651-1661. Millero, F.J. and Schreiber, D.R., 1982. Use of the ion pairing model to estimate activity coefficients of the ionic components of natural waters. Am. J. Sci., 282: 15081540. Milliman, J.D., Gastner, M. and Mfiller. J., 1971. Utilization of magnesium in coralline algae. Geol. Soc. Am. Bull., 82: 35-50. Morrow, D.W. and Ricketts, B.D., 1986. Chemical controls on the precipitation of mineral analogues of dolomite: The sulfate enigma. Geology, 14:408 410. Morse, J.W., Mucci, A. and Millero, F.J., 1980. The solubility of calcite and aragonite in seawater of 35fic salinity at 25°C and atmospheric pressure. Geochim. Cosmochim. Acta, 44: 85-94. Morse, J.W., Zullig, J.J., Bernstein, L.D., Millero, F.J., Milne, P., Mucci, A. and Choppin, G.R., 1985. Chemistry of calcium carbonate-rich shallow water sediments in the Bahamas. Am. J. Sci., 285: 147-185. Mucci, A., 1983. The solubility of calcite and aragonite in seawater at various salinities, temperatures, and one atmosphere total pressure. Am. J. Sci., 283: 780-799. Mucci, A., 1986. Growth kinetics and composition of magnesian calcite overgrowths precipitated from sea-
320 water: Quantitative influence of orthophosphate ions. Geochim. Cosmochim. Acta, 50: 2255-2265. Mucci, A., 1988. Manganese uptake during calcite precipitation from seawater: Conditions leading to the formation of a pseudokutnahorite. Geochim. Cosmochim. Acta, 52: 1859-1868. Mucci, A. and Morse, J.W., 1983. The incorporation of Mg ~+ and Sr '-'+ into calcite overgrowths: influences of growth rate and solution composition. Geochim. Cosmochim. Acta, 47: 217-233. Mucci, A. and Morse, J.W., 1984. The solubility of calcite in seawater solutions of various magnesium concentration, I~=0.697 m at 25°C and one atmosphere total pressure. Geochim. Cosmochim. Acta, 48: 815-822. Plummer, L.N. and Busenberg, 1982. The solubilities of calcite, aragonite and vaterite in CO2-H20 solutions between 0 and 90 ° C and an evaluation of the aqueous model for the system CaCQ-CO2-H20. Geochim. Cosmochim. Acta, 46: 1011-1040. Plummer, L.N. and Sundquist, E.T., 1982. Total individual ion activity coefficients of calcium and carbonate in seawater at 25 ° C and 35%c salinity, and implications to the agreement between apparent and thermodynamic constants of calcite and aragonite. Geochim. Cosmochim. Acta, 46: 227-258. Reddy, M.M. and Nancollas, G.H., 1976. The crystallization of calcium carbonate, IV. The effect of magnesium, strontium and sulfate ions. J. Crystal Growth, 35: 3338.
Reddy, M.M., Plummer, L.N. and Busenberg, E., 1981. Crystal growth of calcite from calcium bicarbonate solutions at constant Pco~ and 25°C: A test of a calcite dissolution model. Geochim. Cosmochim. Acta, 45: 1281-1289, Sass, E., Morse, J.W. and Millero, F.J., 1983. Dependence of the values of calcite and aragonite thermodynamic solubility products on ionic models. Am. J. Sci., 283: 218-229. SjSberg, E.L., 1978. Kinetics and mechanism of calcite dissolution in aqueous solutions at low temperatures. Acta Univ. Stockholm., Stockholm Contrib. Geology, 32 ( 1): 1-92.
Takano, B., 1985. Geochemical implications of sulfate in sedimentary carbonates. Chem. Geol., 29: 393-403. Takano, B., Asano, Y. and Watanuki, K., 1980. Characterization of sulfate ion in travertine. Contrib. Mineral. Petrol., 72:197 203. Walter, L.M., 1986. Relative efficiency of carbonate dissolution and precipitation during diagenesis: A progress report on the role of solution chemistry. In: D.L. Gauthier (Editor), Roles of Organic Matter in Sediments Diagenesis. Soc. Econ. Paleontol. Mineral., Spec. Publ., 38: 1-11. Wray, J.L. and Daniels, F., 1957. Precipitation of calcite and aragonite. Am. Chem. Soc. J., 79: 2031-2034.
309
[3]
THE SOLUBILITY OF CALCITE AND ARAGONITE IN SULFATEFREE SEAWATER AND THE SEEDED GROWTH KINETICS AND COMPOSITION OF THE PRECIPITATES AT 25 °C ALFONSO MUCCI, RENI~ CANUEL and SHAOJUN ZHONG Department of Geological Sciences, McGiU University, Montreal, Que. H3A 2A7 (Canada) (Received January 22, 1988; revised and accepted September 6, 1988)
Abstract Mucci, A., Canuel, R. and Zhong, S., 1989. The solubility of calcite and aragonite in sulfate-free seawater and the seeded growth kinetics and composition of the precipitates at 25 ° C. Chem. Geol., 74: 309-320. The influence of dissolved sulfate ions on the precipitation kinetics and composition of calcite and aragonite overgrowths was investigated in artificial seawater solutions at 25 ° C using a constant disequilibrium technique. The solubility of both minerals in sulfate-free seawater was determined to allow the expression of the saturation state of precipitating solutions as the ratio of the CaCO.~ ion concentration product to their stoichiometric solubility constant. SO~- ions inhibit both calcite and aragonite precipitation. The inhibition in seawater is not selective and therefore, whether in the presence or absence of 0.028 m S042- , aragonite precipitates more rapidly on aragonite seeds than calcite does on calcite seeds from supersaturated solutions. SrC03 incorporation in aragonite is independent of the precipitation rate and unaffected by S042- ions in seawater. The distribution coefficient of Sr 2+ in aragonite overgrowths precipitated from seawater is 0.97 _+0.08 while a value of 1.05 ± 0.09 is obtained in sulfate-free seawater. In contrast, M g C Q incorporation in calcite overgrowths is influenced significantly by the presence or absence of SO~ions in seawater. While the distribution coefficient of Mg 2+ in calcite is also independent of the precipitation rate, its value increases from 0.0172 ± 0.0022 in seawater to 0.0230 ± 0.0026 in sulfate-free seawater, corresponding, respectively, to the precipitation of 8.1 ± 1 and 10.5 ± 1 mole% magnesian calcites.
1. I n t r o d u c t i o n
It has often been reported that sulfate ions in solution inhibit both calcite dissolution and precipitation rates (Reddy and Nancollas, 1976; SjSberg, 1978; Busenberg and Plummer, 1985; Walter, 1986). In addition, several investigators (Kitano, 1962; Kinsman and Holland, 1969; Kitano et al., 1975; Walter, 1986) have indicated that the presence of sulfate ions favored the formation of aragonitic CaCOs rather than calcite, while the work of Bischoff and Fyfe
0009-2541/89/$03.50
(1968) demonstrated the inhibition of aragonite-calcite transformation by dissolved S O l - . Seawater contains a relatively high concentration of sulfate ions ( ~ 28 mM at a salinity of 35). The presence of SO~- and Mg 2+ ions is believed to be responsible for the preferential nucleation of aragonite rather than the more stable calcite in highly supersaturated seawater solutions. However, during the diagenesis of organic-rich sediments and following oxic and suboxic diagenesis, dissolved sulfate is used by bacteria as an alternate electron acceptor in the process of organic matter oxidation. Under the
© 1989 Elsevier Science Publishers B.V.
310
appropriate conditions (i.e. closed or restricted sedimentary environment) this metabolic reaction may result in the depletion of SO~- ions in the sediment pore waters. In addition, sulfate reduction is accompanied by the production of alkalinity which often leads to an increase in the pore-water saturation state with respect to CaCO3 and the precipitation of carbonate minerals. Calcite rather than aragonite is most often found as an authigenic phase in these sediments (Bathurst, 1975). Marine waters with low S042- concentrations are also believed (Baker and Kastner, 1981; Kastner, 1984; Baker and Burns, 1985) to be conducive to dolomite formation, but this idea has not yet received widespread recognition as a possible dolomitization mechanism (Morrow and Ricketts, 1986; Hardie, 1987). The purpose of this paper is to provide quantitative data on the growth kinetics of both calcite and aragonite in "normal" and sulfate-free seawater and determine the composition of the calcium carbonate overgrowths precipitated from these solutions at 25 ° C.
2. M a t e r i a l s a n d m e t h o d s
2.1. Solids, solutions and experimental procedures Baker ® "Instra-analyzed flux reagent" grade calcium carbonate was used for all experiments in which calcite was studied, both for the solubility measurements and as a seed material for the precipitation reactions. Procedures for cleaning and size separating this material have been described previously (Mucci, 1986). Aragonite synthesized in the laboratory by the method of Wray and Daniels (1957) as modified by Katz et al. (1972) at a temperature of 70 ° C was used without further treatment for all experiments in which aragonite was studied. Xray diffraction spectrometry (XRD) and scanning electron microscopy examination (SEM) of this material indicate the absence of vaterite and the presence of < 1 wt.% calcite. The cal-
cite and aragonite have surface areas of 0.52 and 3.40 m 2 g- 1, respectively, as determined by the Kr-BET method of de Kanel and Morse (1979). Aged artificial seawater was used for all the experiments. The artificial seawater of salinity 35 was prepared to include all major elements of natural seawater including F - according to the method of Kester et al. (1967), modified to fit the analysis of Millero (1974). The sulfatefree seawater was prepared according to the same procedure but excluding sulfate. The ionic strength of the solution was balanced to equal that of a standard seawater (It=0.697 m) by increasing the amount of NaC1 added to the preparation. The seawater solution was stored in a polypropylene container for two months to reduce the dissolved phosphate concentration ( < 1/IM) prior to use. Solubility measurements were carried out in a closed system at 25°C over long periods of equilibration by a procedure described at length by Morse et al. ( 1980 ) and Mucci ( 1983 ). Equilibration was followed from both undersaturated and supersaturated solutions. Precipitation reactions were conducted in an open system in solutions of close to constant composition over a wide range of precipitation rates. Constancy of composition was maintained during the length of the precipitation by the use of a chemo-stat technique through the simultaneous injection of two titrants in equal amounts by a dual syringe pump. The mixture of the two titrants reproduced the exact composition of the precipitating solution plus an excess in calcium, strontium and carbonate alkalinity to compensate for the calcium carbonate precipitation. The temperature of the precipitating solution was held constant at 25 (+0.05)°C by circulating water through a jacketed 400-ml glass reaction vessel from a constant-temperature bath. The Pco~ of the solution was kept nearly constant at ~ 3000 ppm or 10 -2.5 atm. by bubbling a CO~-N2 gas mixture of known composition. A detailed description of the chemo-stat, its working concept and titrant composition have been presented pre-
311
viously (Mucci and Morse, 1983; Mucci, 1986). The precipitation was carried out on ~ 0.20.7 g of seed material in 300-400 ml of seawater solution until > 10 -4 mol of carbonate was precipitated. The amount of carbonate precipitated was calculated from the weight of the standardized titrant used during the length of the experiment and the carbonate alkalinity of the solution before and after the precipitation reaction (according to equation 1 in Mucci, 1986).
2.2. Calculation o[ the stoichiometric solubility constants The stoichiometric solubility constant of calcite (or aragonite) in an electrolyte solution is defined as: K c* (or a) = [Ca 2+1 [COW-]
(1)
where [i] is the equilibrium total ion concentration. While [Ca 2+ ] was measured directly by titration, [CO~- ] was derived from the carbonate alkalinity, the equilibrium pH of the solution and the carbonic acid second apparent dissociation constant. The carbonate alkalinity was obtained from the titration alkalinity from which the boric acid contribution was subtracted. Analytical procedures for the determination of Ca 2+ and titration alkalinity have been described previously (Mucci, 1986). The equilibrium pH was measured using a combination glass electrode which had been previously calibrated against U.S. National Bureau of Standards (N.B.S.) buffer solutions (6.862, 7.413 and 9.180). The first apparent dissociation constant of boric acid and the second apparent dissociation constant of carbonic acid in sulfate-free seawater were derived from their standard seawater values (salinity 35, 25 °C) as they appear in Millero ( 1979 ) and the ratio of the total ion activity coefficients of the appropriate ions in solution. Thus,
K~(sf) =Ki3tsw) 7(B(OH)~-)sw ~(B(OH)4)si
=2.16-10 -9
(2)
and
Y(co,~ - )~ Y(HCO:r )~w
=8.61"10 -1°
(3)
where Ki~ and K.~ are the first apparent dissociation constant of boric acid and second apparent dissociation constant of carbonic acid, respectively, in seawater (sw) and sulfate-free seawater (sf); and 7 ( f s are the total ion activity coefficients. Total ion activity coefficient estimates were calculated from the ion pairing model of Millero and Schreiber (1982).
2.3. Steady-state solution composition and saturation state of the precipitating solution The steady-state ion concentration product, [Ca 2+ ] [CO~- ], of the precipitating solution was calculated using the steady-state pH measured during the experiment, the carbonate alkalinity and the Ca 2+ concentration of an aliquot withdrawn at the end of the precipitation. For most measurements the steady-state saturation of the solution was reached within 15 min. and remained constant for the length of the experiment. The steady-state pH was monitored by two sets of electrodes which were calibrated against N.B.S. buffer solutions. The stoichiometric solubility constants of calcite (or aragonite) determined in this study in sulfatefree seawater and by Mucci ( 1983 ) in "normal" seawater were used to calculate the saturation state of the precipitating solution with respect to calcite (or aragonite), according to: ~2c(or a)
--
[Ca 2+1 [COW-] /~c(or a)
(4)
Uncertainties associated with estimates of the saturation state, which might have been introduced by u s i n g the N.B.S. buffer calibration
312
system in strong electrolyte solutions (e.g., liquid junction potentials, etc.), are minimized since b o t h / ~ c (or a) and the steady-state C a C Q ion concentration product of the precipitating solution were determined using the same electrode and p H scale. The Sr 2+ concentration of the solutions was measured before and after the reaction by flame atomic absorption spectrophotometry (AAS) using aqueous standards in a seawater matrix for calibration. Precision of the AAS analysis is estimated to be better than + 5%.
2.4. Overgrowth composition The Mg content of the calcite overgrowths and the Sr content of the aragonite overgrowths were determined following the acid digestion of a known amount of reacted solid by AAS. The mole fraction of MgC03 and SrCO3 in the overgrowth was calculated from the results of the AAS analyses, the amount of carbonate precipitated and the amount of material dissolved for analysis (Mucci and Morse, 1984 ). No correction was introduced to compensate for the presence of residual solution salts as the overgrowths were rinsed with calcite-equilibrated distilled water after being filtered out of the parent solution (Mucci, 1986). The reacted solid was also examined by X R D spectrometry to determine its composition and mineralogy and identify other carbonate mineral phases which might have precipitated along with the calcite and aragonite. Powder packs were prepared and irradiated using a Siemens ® model D-500 X-ray diffractometer. The Cu-K~ wavelength radiation was used as a source and the diffraction spectra was recorded using a proportional counter detector. 3. R e s u l t s
3.1. Solubility of calcite and aragonite in sulfate-free seawater Analytical results and calculated CaCO3 ion concentration products of sulfate-free seawater
solutions equilibrated with calcite and aragonite are presented in Tables I and II. The C a C Q ion concentration products measured after 188 days of equilibration differ by no > + 7% and their averages are taken to represent the stoichiometric solubility constants of calcite and aragonite in sulfate-free seawater at 25°C and 1-atm. total pressure. This is also taken to be an indication that the solubilities are reversible since measurements were carried out from both undersaturated and supersaturated solutions. Stoichiometric solubility constant values are 4.07+0.12 and 5.94+0.08 for calcite and aragonite, respectively. These values were used in eq. 4 to calculate the saturation state of the precipitating solutions in the overgrowth experiments.
3.2. Influence of dissolved sulfate on the precipitation kinetics of calcite and aragonite in seawater Selected parameters and results from single precipitation reactions are presented in Tables III-V. The precipitation rate data were fitted to the following empirical kinetic equation: R = k ( ~ c ( .... } - - 1 ) n
(5)
or to its logarithmic form, l o g R = l o g k + n l o g ( t 2 c ( .... I - 1 )
(6)
where R is the precipitation rate normalized to the reactive surface area (mol hr.-~ m - 2 ) ; k is the rate constant; n is the empirical reaction order; and (t2~(or at - 1 ) is the degree of supersaturation with respect to calcite or aragonite. The precipitation rate data for calcite and aragonite in both "normal" and sulfate-free seawater are plotted in Fig. l a and b as a function of the logarithm of their respective supersaturations. The kinetic parameters n (order of reaction ) and log k (rate constant) for each phase in both solutions are presented in Table VI. As was expected from results of previous studies, the precipitation rate of calcite in sulfate-bearing or "normal" seawater is slower
313 TABLE I
Calcite stoichiometric solubility constant in sulfate-free artificial seawater, It = 0,697 m at 25 °C Approach
Equilibration period
[Ca 2+ ]
At (meq kg- ~)
PH T
(mmol k g - ~)
M
(107 mol ~ kg ~)
(days) UND SUP
2 2
10.23 10.19
0.844 1.670
7.800 7.552
3.97 4.83
UND SUP
13 13
10.13 9.87
0.891 1.849
7.773 7.531
3.95 4.96
UND SUP
188 188
10.11 9.74
0.923 1.349
7.749 7.613
3.90 4.22
UND SUP
231 231
10.13 9.72
0.930 1.224
7.753 7.654
3.97 4.15
UND SUP
248 248
10.24 9.82
0.963 1.233
7.730 7.648
3.98 4.17
UND = approach to equilibrium from undersaturation; S U P = approach to equilibrium from supersaturation. TABLE II
Aragonite stoichiometric solubility constant in sulfate-free artificial seawater, I, = 0.697 m at 25 c C Approach
Equilibration period
[Ca `-'+ ] (mmol kg- ~)
A, (meq kg ')
pH Na~s
K: (107 mol e kg e)
(days) UND SUP
2 2
9.83 9.67
1.068 1.866
7.930 7.660
6.28 6.44
UND SUP
13 13
9.68 9.54
1.087 1.654
7.918 7.650
6.16 5.50
UND SUP
188 188
9.77 9.29
1.158 1.671
7.850 7.696
5.84 5.96
UND SUP
231 231
9.78 9.23
1.172 1.533
7,853 7,733
5.95 5.85
U ND SUP
248 248
9.87 9.41
1.206 1.496
7,839 7,750
6.02 6.02
For explanation of UND and SUP see footnote to Table I.
than in sulfate-free seawater at all supersaturations investigated. The rates differ by a factor of 2 (at f2c=3.5) to 3 (at t2c=17). Extrapolation of the least-squares fit lines to lower saturations indicates that the precipitation rate of calcite in both solutions would be identical at a f2c = 1.4. A similar pattern is observed for aragonite in Fig. lb. Above a saturation state, f2a, of 1.7 the precipitation rate of aragonite in sulfate-free seawater becomes faster than in "nor-
mal" seawater. Rates differ by a factor of 4 at ~a = 4. The rates of calcite and aragonite precipitation were also plotted as a function of the supersaturation with respect to calcite in "normal" (Fig. 2a) and sulfate-free (Fig. 2b) seawater to allow comparison of the relative precipitation rates of the two phases in the same solution at identical CaCO:~ ion concentration products. In both supersaturated solutions ara-
314 T A B L E III Kinetic data and composition of aragonite overgrowths precipitated from artificial seawater at 25 : C Experiment No.
[Sr 2+ ] (ppm)
[Ca 2+ ] (mmol kg -1)
pH TM
~l~ a
Amount of carbonate precipitated (104mol)
-log(rate) (mol hr. ' m -2)
SrCO:~ (mole%)
Zl Z2 Z3 Z4 Z5 Z6 Z7 Z8 Z9 ZIO Zll Z12 Z13 Z15
6.7 6.7 6.6 6.7 7.0 6.9 7.0 6.7 --5.1 7.5 8.3 8.1
8.83 8.84 8.84 9.24 9.67 9.87 10.08 9.21 9.66 9.62 9.27 9.75 9.96 9.81
7.896 7.829 7.737 7.714 7.852 7.637 7.718 7.785 7.787 7.873 7.937 7.653 7.672 7.595
3.77 2.87 2.33 -3.38 1.44 -2.84 2,14 3.93 4.95 1.64 1.65 1.19
15.6 15.4 15.6 14.5 11.4 12.0 11.1 14.8 4.26 12.6 14.3 3.18 3.22 0.841
4.010 4.290 4.584 -4.009 4.950 -4.318 4.539 3.703 3.437 4.776 5.066 5.588
0.88 0.75 0.81 0.80 0.73 -0.83 0.68 --0.63 ----
D~")
1.05
0.89 0.98 0.99 0.92 1.08
0.84
1.03
*Calculated from the AAS analysis and the [Sr 2+ ] / [ C a 2. ] ratio of the final solution.
TABLE IV Kinetic data and composition of calcite overgrowths precipitated from sulfate-free artificial seawater at 2 5 C Experiment No.
R1 R2 R3 R4 R6 R7 R8 R9 R12 R13 RI4 RI5 RI6 R17 R18
[Ca ~+] (mmol kg 1)
10.47 10.31 10.75 10.98 11.00 11.10 10.84 10.70 10.93 10.76 10.52 10.52 10.57 10.38 9.56
pH TM
7.934 7.876 7.869 7.887 8.055 7.772 7.990 7.805 7.848 7.704 7.854 7.857 8.032 7.767 7.645
£2,.
12.25 9.54 10.01 9.96 13.63 6.60 12.52 7.43 9.02 4.73 -9.62 14.9 5.51 3.43
Amount of carbonate precipitated (104 mol)
-log(rate) (mol hr. -I m -~)
10.5 11.7 9.28 9.70 9.53 8.68 9.83 10.4 9.59 9.51 11.4 11.2 9.05 24.8 22.3
2.964 3.242 3.245 3.100 2.806 3.902 2.957 3.666 3.390 4.023 -3.251 2.664 4.321 4.907
*Calculated from the AAS analysis and the [Mg 2+ ] / [ C a `-,+] ratio of the final solution.
MgCO:~ (mole%) ( × 1 0 ~) AAS
XRD
9.6 10.1 10.2 8.7
-9.0 ---
2.11 2.20 2.31 1.99
11.7 9.7
11.0 9.0
2.80 2.21
--
12.0
--
11.8 11.2 10.3 9.0 10.0
-11.0 9.5 --
2.74 2.52 2.30 1.98 2.18
--
315
TABLE V Kinetic data and composition of aragonite overgrowths precipitated from sulfate-free artificial seawater at 25 °C Experiment No.
[ Sr~+ ] (ppm)
[Ca 2+ ] (mmol kg -1)
PH
R28 R29
8.0 7.7 7.9 8.5 8.5 8.5 8.2 8.7 8.2 8.2 8.0 8.2 8.0 8.1 7.7 7.8
9.79 9.73 9.60 10.87 10.90 10.32 10.13 10.53 10.05 10.39 10.33 10.33 10.16 10.18 9.96 10.24
7.699 7.697 7.683 7.531 7.554 7.665 7.644 7.617 7.629 7.580 7.539 7.660 7.672 7.705 7.562 7.562
R30 R31 R32 R33 R34 R35 R36 R37 R38 R39
R40 R4I R43 R44
-2 O0
i (a)
I
T M
~a
Amount of carbonate precipitated ( 104 mol )
-log(rate) (molhr.-~ rn 2)
SrCQ (mole%)
D~r
3.20 2.76 2.70 -1.82 2.82 2.59 2.25 2.33 1.97 -2.85 3.04 3.44 1.73 1.77
12.4 12.8 12.1 8.25 8.50 10.5 10.3 8.59 10.9 9.39 9.16 10.1 10.3 10.5 10.9 9.47
3.724 ?1.022 4.029 -4.597 3.848 4.012 4.373 4.105 4.513 -3.736 3.638 3.529 4.970 4.824
0.91 0.91 0.93 1.05 1.01 0.89 -0.93 0.82 0.90 1.02 I).85
1.00 1.03 1.02 1.21 1.17 0.97 -1.01 0.91 1.04 1.19 0.96
-
-
,
I
i
-
-0.93 --
,
~
t
-
i
,
o 'E
5
/ j
J
j
g
o
~ o
o!~o
'
o!8o
,'.oo
•
-4 O0
x-
g
r
!
1
E
o 4o
-
-1.08 --
,!2o
log ( Q c - I )
-500
-o'7o
o! . . . .
!. . . .
11o o!1o I o 9 ( f ~ a - I)
o!3o
o!5o
Fig. 1. Precipitation kinetics of calcite (a) and aragonite (b) overgrowths in "normal" and sulfate-free seawater at 25 ~C as a function of their respective supersaturations; . . . . calcite in seawater (from Mucci, 1986); A =calcite in sulfate-free seawater; • = aragonite in seawater; o = aragonite in sulfate-free seawater.
gonite precipitates more rapidly on aragonite seeds than calcite does on calcite seeds.
3.3. Influence of dissolved sulfate ions on the composition of calcite and aragonite overgrowths precipitated from seawater Results of the X R D analysis of calcite and aragonite overgrowths precipitated from sulfate-free seawater confirmed that the mineral-
ogy of the overgrowths is controlled by the nature of the seed material. Only magnesian calcite overgrowths could be detected when calcite seeds were used while the presence of no other phase could be detected on the reacted aragonite seed material. When it could clearly be resolved from the calcite seed material the position of the magnesian calcite overgrowth peak was used to determine its MgCO:, content using the data of Goldsmith et al. (1961) as
316 T A B L E VI Kinetic parameters for calcite and aragonite precipitation rates in seawater and sulfate-free seawater at 25 ~C, expressed according to the empirical rate law described in eq. 6 Mineral
Solution
n
- log k
Source
Calcite
seawater
2.8
6.29
1.9
5.41
Mucci and Morse ( 1983 ); Mucci ( 1986, 1988 ) Burton and Walter (1987
SOs-free seawater
3.0
6.08
this study
seawater
1.6 1.7
4.63 4.39
this study Burton and Waiter (1987
S04-free seawater
2.2
4.43
this study
Aragonite
I (a) ,~;/
//
L o
-6.oo
g
~ / 7 '
oo
-400 /
/
jj
J
f
-6 O0
-]
E
E
-5 oo
J o!oo
'
oL2o
'
log
I
046 ('~C-
060
. I. . . .
0!80
!,o
o6o' log
I}
(~'c-
o!6o
,!oo
I)
Fig. 2. Precipitation kinetics of calcite and aragonite overgrowths in "normal" (a) and sulfate-free (b) seawater at 25 C as a function of the supersaturation with respect to calcite. Symbols are as in Fig. 1.
presented by Milliman et al. (1971). These results are presented in Table IV. Since the calcite and aragonite overgrowths were precipitated from solutions of nearly constant composition despite the ongoing precipitation, the partitioning of a trace component, X 2+ (e.g., X 2 + = M g 2 + , Sr2+), between the CaCO:~ solid and the solution can be defined by the H e n d e r s o n - K r a c e k (Henderson and Kracek, 1927) or homogeneous distribution coefficient: c
c
Mx2./Mca2+ D ~ + -- Mx~'+M / ca~+L L
(7)
where Mx2+ and Mca2+ denote the molar concentrations of X 2+ and Ca 2+, respectively, in the C a C Q overgrowth (c) and the parent solution (L). The distribution coefficients measured in this
study are independent of the precipitation rate. The averaged distribution coefficient of Mg 2+ in calcite overgrowths precipitated from sulfate-free seawater is 0.0230_+0.0026, corresponding to the precipitation of a 10.5_+ 1 mole% magnesian calcite. This value is significantly higher than the one obtained in "normal" seawater (0.0172 _+0.0022 or 8.1 + 1 mole% MgCO~) and cannot be explained by the small variation (i.e. < -+ 1% ) in the ratio of the activity of Mg ~+ to Ca 2+ ions between the two solutions. Processes occurring at or near the surface of the growing crystal are probably responsible for the discrepancy. This hypothesis will be expanded upon further in Section 4.3. On the other hand, the averaged distribution coefficients of Sr "'+ in aragonite overgrowths precipitated from seawater (0.97_+0.08) and sulfate-free seawater ( 1.05 _+0.09 ) are statisti-
317
cally identical. This is consistent with the fact that the ratio of the ion activity of Sr 2+ to Ca s + in the parent solutions are the same ( < +_ 1% ). The values of D ~ . determined in this study are in excellent agreement with the results of Kinsman and Holland ( 1969 ) and Kitano et al. (1971). Kinsman and Holland (1969) measured the distribution coefficient of Sr s + in aragonite precipitated from seawater between 16 ° and 96~C. A value of 1.13 + 0.04 at 25°C can be interpolated from the linear relationship they observed between the distribution coefficient and temperature. Kitano et al. (1971) determined a distribution coefficient of 1.1-2_ 0.1 following the precipitation of aragonite from a Ca (HCO:~) 2-MgC12-SrCls solution at 20 ° C. 4. D i s c u s s i o n
4. I. Solubility of calcite and aragonite in sulfate-free seawater Calcite and aragonite do not precipitate as pure C a C Q from seawater and yet, the calcite and aragonite solubility behavior in seawater closely follows that predicted from thermodynamics of pure calcium carbonate (Plummer and Sundquist, 1982; Mucci, 1983). The same conclusion can be shown to apply to their solubility behavior in sulfate-free seawater by comparing the C a C Q ion activity coefficient product in sulfate-free seawater estimated from model calculations with the value derived from solubility measurements. The C a C Q ion activity coefficient product can be calculated from the ratio of the thermodynamic to the stoichiometric solubility constants according to: 0 t ~r 7~c~,~*~)'(co~ ~= (Ksp/K~p)
02
(8)
where K °" is the thermodynamic solubility constant of calcite or aragonite at 25 °C and 1-atm. total pressure (pK(,! = 8.48 _+0.02, p K ° = 8.34 ± 0.03; from P l u m m e r and Busenberg, 1982; Mucci, 1983; Sass et al., 1983); and 0 is a correction factor used to convert the total ion ac-
tivity coefficient product from kg 2 sw m o l - " to kg 2 H 2 0 mo1-2 (0=0.964). CaCO3 ion activity coefficient product values calculated from eq. 8 and the calcite and aragonite solubility constants (0.0076 ± 0.0006 and 0.0072 +_0.0006, respectively) compare thvorably with a value of 0.0065 _+ 10% estimated using the ion pairing model of Millero and Schreiber (1982).
4.2. Influence of dissolved sulfate ions on the relative rates of calcite and aragonite precipitation from seawater Based on the results described in Section 3.2 it appears that a decrease of the sulfate content of seawater following sulfate reduction in anoxic sediments would not result in a change in the mineralogy of authigenic C a C Q . These results apparently contrast with conclusions reached by previous investigators (Kitano, 1962; Kitano et al., 1975; Walter, 1986) who noted that dissolved SO~ ions favor aragonite precipitation by selectively inhibiting calcite precipitation, thus leading to a reversal in the relative precipitation rate of the two minerals following addition of SO4- to the precipitating solution. However, the high [Mg ~+ ] / [ C a ~+ ] or [Mg 2+ ] in seawater is responsible for inhibiting calcite precipitation even in the absence of SO4-, effectively shadowing its influence on the relative precipitation rate of the two minerals. Calcite and aragonite inhibition by SO4- is likely due to a combination of' its adsorption on the surface and incorporation in the mineral structure. Sulfate adsorption at Ca ~+ sites and partial binding of CaSO ° ion pairs on the surface of calcite and aragonite require that the CaS04 bond be broken and SO~- be desorbed from the surface before growth can occur. Furthermore, several workers (Kitano et al., 1975; Takano et al., 1980; Busenberg and Plummer, 1985; Takano, 1985) have shown that SO4 ions are more easily incorporated in calcite than aragonite. In tact, Busenberg and Plummer
318 (1985) observed that the incorporation of SO~-- in synthetic calcites increased with the precipitation rate and sulfate concentration, and resulted in an increase in the solubility and unit cell dimensions of the sulfate-bearing calcites. Processes accompanying sulfate reduction in marine sediments, such as the release and accumulation of dissolved phosphate (Walter, 1986) and organic matter decomposition products (Berner et al. 1978) and an increase in the Pco~ in the interstitial waters (Reddy et al., 1981; Given and Wilkinson, 1985) may influence the mineralogy of marine carbonate cements. In fact, Walter (1986) observed that phosphate ions are selective in inhibiting aragonite precipitation from seawater. Temperature (Burton and Walter, 1987) and substrate mineralogy also certainly play important roles
reaction and may behave differently on the surface of the mineral. The incorporation of sulfate in calcite may help to relieve or accentuate lattice strain induced by the incorporation of Mg 2. in magnesian calcites. Mg 2+ ions (0.65 ~,) are smaller than Ca 2+ ions (0.99 A) while S O l - ions ( ~ 1.19 ,~3) are significantly larger than CO~ions (~0.63 A~). Whether or not this mechanism could buffer Mg 2+ incorporation in calcite precipitated from seawater is uncertain but warrants further investigation. Takano (1985) claimed that S O l - ions in inorganically precipitated calcite occupy lattice sites replacing carbonates and that high-Mg calcites can incorporate more sulfate than if Mg-free.
4.3. Influence of sulfate ions on the composition of calcite overgrowths precipitated from seawater
Based on a comparison of calculated and estimated C a C Q total ion activity coefficients it appears that the solubility or calcite and aragonite in sulfate-free seawater closely follows the predicted solubility of pure CaCO:~. Quantitative seeded precipitation kinetic data of calcite and aragonite in "normal" and sulfate-free seawater indicate that dissolved SO42- ions inhibit the precipitation of both minerals in seawater. The inhibition is not selective and could not be responsible for a reversal in the mineralogy of authigenic CaCO:~ precipitated following a decrease in S O l - concentration in sediment pore waters as a result of microbiologically-mediated sulfate reduction. SrCO., incorporation in aragonite overgrowths is unaffected by the removal of SO] ions from seawater. The distribution coefficient of Sr 2+ in aragonite precipitated from seawater is independent of the precipitation rate and approximately equal to 1. In contrast, MgCO3 incorporation in calcite is influenced significantly by the presence or absence of SO~- ions. The distribution coefficient of Mg 2+ in calcite precipitated from sulfate-free seawater is 34% larger than in "normal" seawater.
One of the more interesting findings of this study is the strong influence which dissolved SO]- ions appear to exert on the amount of Mg 2+ incorporated in magnesian calcite overgrowths. In the absence of SO42- ions the distribution coefficient of Mg 2+ in calcite precipitated from seawater increases by 34%. This result has significant implications with respect to the genesis of inorganic high-Mg calcites and dolomite in marine sedimentary environments. However, as the activity ratio of Mg 2+ and Ca 2+ in the two solutions does not vary by more than + 1% we can only speculate on the processes responsible for the increased incorporation of Mg 2+ in calcite in the absence of sulfate. Variation of the distribution coefficient can only be explained partially in terms of an increase ( ~ 4 . 4 % ) in the concentration of MgCOI~ ion pairs relative to CaCO ° between the two solutions. These species may be important intermediate "reactants" in the precipitation
5. Conclusions
319
Acknowledgements The authors wish to express their gratitude to Marco Ciarlo for his technical assistance during most stages of this study. Financial support for this study was provided by the Natural Sciences and Engineering Research Council of Canada through grants Nos. U0432 and E1446. This project also benefited from the award of an equipment grant by the Faculty of Graduate Studies and Research, McGill University. Finally, one of us (S.Z.) wishes to acknowledge the financial support of the Government of the People's Republic of China. References Baker, P.A. and Burns, S.J., 1985. Occurrence and formation of dolomite in organic-rich continental margin sediments. Am. Assoc. Pet. Geol. Bull., 69: 1917-1930. Baker, P.A. and Kastner, M., 1981. Constraints on the formation of dolomite. Science, 213: 214-216. Bathurst, R.G.C., 1975. Carbonate Sediments and Their Diagenesis. Elsevier, Amsterdam, 2nd ed., 658 pp. Berner, R.A., Westrich, J.T., Graber, R., Smith, J. and Martens, C.S., 1978. Inhibition of aragonite precipitation from supersaturated seawater. Am. J. Sci., 278: 816837. Bischoff, J.L. and Fyfe, W.S., 1968. Catalysis, inhibition, and the calcite-aragonite problem, I. The aragonitecalcite transformation. Am. J. Sci., 266: 65-79. Burton, E.A. and Walter, L.M., 1987. Relative precipitation rates of aragonite and Mg calcite from seawater: Temperature or carbonate ion control? Geology, 15:111114. Busenberg, E. and Plummer, L.N., 1985. Kinetic and thermodynamic factors controlling the distribution of SO~- and Na + in calcites and aragonites. Geochim. Cosmochim. Acta, 49: 713-725. de Kanel, J. and Morse, J.W., 1979. A simple technique for surface area determination. J. Phys. E, Sci. Instrum., 12: 272-273. Given, R.K. and Wilkinson, B.H., 1985. Kinetic control of morphology, composition, and mineralogy of abiotic sedimentary carbonates. J. Sediment. Petrol., 55: 109119. Goldsmith, J.R., Graf, D.L. and Heard, H.C., 1961. Lattice constants of the calcium-magnesium carbonates. Am. Mineral., 46: 453-457. Hardie, L.A., 1987. Dolomitization: A critical view of some current views. J. Sediment. Petrol., 57: 166-183.
Henderson, L.M. and Kracek, F.C., 1927. The fractional precipitation of barium and radium chromates. J. Am. Chem. Soc., 49: 739-749. Kastner, M., 1984. Control of dolomite formation. Nature (London), 310: 410-411. Katz, A., Sass, E., Starinsky, A. and Holland, H.D., 1972. Strontium behavior in the aragonite-calcite transformation: an experimental study at 40-98 °C. Geochim. Cosmochim. Acta, 36: 481-496. Kester, D.R., Duedall, J.W., Connors, D.N. and Pytkowicz, R.M., 1967. Preparation of artificial seawater. Limnol. Oceanogr., 12: 176-179. Kinsman, D.J.J. and Holland, H.D., 1969. The co-precipitation of cations with CaCO:,, IV. The co-precipitation of Sr 2+ with aragonite between 16 ° and 96 ° C. Geochim. Cosmochim. Acta, 33: 1-17. Kitano, Y., 1962. The behavior of various inorganic ions in the separation of calcium carbonate from a bicarbonate solution. Bull. Chem. Soc. Jpn., 35: 1973-1980. Kitano, Y., Kanamori, N. and Oomori, T., 1971. Measurements of distribution coefficients of strontium and barium between carbonate precipitate and solution - Abnormally high values of distribution coefficients measured at early stages of carbonate formation. Geochem. J., 4:183 206. Kitano, Y., Okumura, M. and Idogaki, M., 1975. Incorporation of sodium, chloride and sulfate with calcium carbonate. Geochem. J., 9:75 84. Millero, F.J., 1974. The physical chemistry of seawater. Annu. Rev. Earth Planet. Sci. Lett., 2:101 150. Millero, F.J., 1979. The thermodynamics of the carbonate system in seawater. Geochim. Cosmochim. Acta, 43: 1651-1661. Millero, F.J. and Schreiber, D.R., 1982. Use of the ion pairing model to estimate activity coefficients of the ionic components of natural waters. Am. J. Sci., 282: 15081540. Milliman, J.D., Gastner, M. and Mfiller. J., 1971. Utilization of magnesium in coralline algae. Geol. Soc. Am. Bull., 82: 35-50. Morrow, D.W. and Ricketts, B.D., 1986. Chemical controls on the precipitation of mineral analogues of dolomite: The sulfate enigma. Geology, 14:408 410. Morse, J.W., Mucci, A. and Millero, F.J., 1980. The solubility of calcite and aragonite in seawater of 35fic salinity at 25°C and atmospheric pressure. Geochim. Cosmochim. Acta, 44: 85-94. Morse, J.W., Zullig, J.J., Bernstein, L.D., Millero, F.J., Milne, P., Mucci, A. and Choppin, G.R., 1985. Chemistry of calcium carbonate-rich shallow water sediments in the Bahamas. Am. J. Sci., 285: 147-185. Mucci, A., 1983. The solubility of calcite and aragonite in seawater at various salinities, temperatures, and one atmosphere total pressure. Am. J. Sci., 283: 780-799. Mucci, A., 1986. Growth kinetics and composition of magnesian calcite overgrowths precipitated from sea-
320 water: Quantitative influence of orthophosphate ions. Geochim. Cosmochim. Acta, 50: 2255-2265. Mucci, A., 1988. Manganese uptake during calcite precipitation from seawater: Conditions leading to the formation of a pseudokutnahorite. Geochim. Cosmochim. Acta, 52: 1859-1868. Mucci, A. and Morse, J.W., 1983. The incorporation of Mg ~+ and Sr '-'+ into calcite overgrowths: influences of growth rate and solution composition. Geochim. Cosmochim. Acta, 47: 217-233. Mucci, A. and Morse, J.W., 1984. The solubility of calcite in seawater solutions of various magnesium concentration, I~=0.697 m at 25°C and one atmosphere total pressure. Geochim. Cosmochim. Acta, 48: 815-822. Plummer, L.N. and Busenberg, 1982. The solubilities of calcite, aragonite and vaterite in CO2-H20 solutions between 0 and 90 ° C and an evaluation of the aqueous model for the system CaCQ-CO2-H20. Geochim. Cosmochim. Acta, 46: 1011-1040. Plummer, L.N. and Sundquist, E.T., 1982. Total individual ion activity coefficients of calcium and carbonate in seawater at 25 ° C and 35%c salinity, and implications to the agreement between apparent and thermodynamic constants of calcite and aragonite. Geochim. Cosmochim. Acta, 46: 227-258. Reddy, M.M. and Nancollas, G.H., 1976. The crystallization of calcium carbonate, IV. The effect of magnesium, strontium and sulfate ions. J. Crystal Growth, 35: 3338.
Reddy, M.M., Plummer, L.N. and Busenberg, E., 1981. Crystal growth of calcite from calcium bicarbonate solutions at constant Pco~ and 25°C: A test of a calcite dissolution model. Geochim. Cosmochim. Acta, 45: 1281-1289, Sass, E., Morse, J.W. and Millero, F.J., 1983. Dependence of the values of calcite and aragonite thermodynamic solubility products on ionic models. Am. J. Sci., 283: 218-229. SjSberg, E.L., 1978. Kinetics and mechanism of calcite dissolution in aqueous solutions at low temperatures. Acta Univ. Stockholm., Stockholm Contrib. Geology, 32 ( 1): 1-92.
Takano, B., 1985. Geochemical implications of sulfate in sedimentary carbonates. Chem. Geol., 29: 393-403. Takano, B., Asano, Y. and Watanuki, K., 1980. Characterization of sulfate ion in travertine. Contrib. Mineral. Petrol., 72:197 203. Walter, L.M., 1986. Relative efficiency of carbonate dissolution and precipitation during diagenesis: A progress report on the role of solution chemistry. In: D.L. Gauthier (Editor), Roles of Organic Matter in Sediments Diagenesis. Soc. Econ. Paleontol. Mineral., Spec. Publ., 38: 1-11. Wray, J.L. and Daniels, F., 1957. Precipitation of calcite and aragonite. Am. Chem. Soc. J., 79: 2031-2034.