The incorporation of Mg2+ and Sr2+ into calcite overgrowths: influences of growth rate and solution composition

The incorporation of Mg*’ and Sr*+ into calcite overgrowths: influences of growth rate and solution composition ALFONSO MUCCI and JOHN W. MORSE Department

of Oceanography, Texas A & M University, College Station, Texas 77843

(Received April 22, 1982; accepted in revisedform

November

1, 1982)

Abstract-The concentrations of Mg*+ and Srr’ incorporated within calcite overgrowths precipitated from seawater and related solutions, determined at 25”C, were independent of the precipitation rate over approximately an order of magnitude. The saturation states used to produce this range of precipitation rates varied from 3 to 17 depending on the composition of the solution. The amount of Mgr’ incorporated in the overgrowths was not directly proportional to Me/CZa*’ in solution over the entire range (l-20) of ratios studied. Below a ratio of 7.5, the overgrowth was enriched in MgC03 relative to what is predicted by the constant distribution coefficient measured above a ratio of 7.5. This increased MgCO3 correlates with the relative enrichment of adsorbed Mg2’. Above a ratio of 7.5 the concentration of MgCO, in the calcite overgrowths followed a classical thermodynamic behavior characterized by a constant distribution coefficient of 0.0123 (~0.008 std dev). The concentration of SrCOr incorporated in the overgrowths was linearly related to the MgCO, content of the overgrowths, and is attributed to increased solubility of SrCO, in calcite due to the incorporation of the smaller Mg” ions. The kinetic data indicate that the growth mechanism involves the adsorption of the cations on the surface of the calcite prior to dehydration and final incorporation. It is suggested that dehydration of cations at the surface is the rate controlling step. IN’IXODUCIlON IN RECENT years there has been a continued interest in the chemistry of “foreign” ions present in the com-

mon carbonate minerals; especially calcite and aragonite at earth surface temperatures (HOLLAND, 1966; TSUSUE and HOLLAND, 1966; CROCKET and WINCHESTER, 1966; KINSMAN and HOLLAND, 1969; KITANO and OOMORI, 197 1; KATZ et aI., 1972; KATZ, 1973; KITANO and OKUMURO, 1973; RAINSWELL and BRIMBLECOMBE,1977; DE KANEL and MORSE, 1978; WHITE, 1977, 1978; LORENS, 1981). Most of these studies were done in solutions significantly different in composition from seawater. Generally, both the saturation state and the concentration of the cations of interest changed significantly during precipitation. No attempts will be made to review this extensive literature, as many of these results are not applicable in solutions similar to seawater in composition, which were of concern in this study. WEYL ( 1965) was a pioneer in demonstrating the significance of overgrowths and surface history to the behavior of carbonates in seawater. His work indicated that M8+ was the most important cationic coprecipitate in calcite formed in the marine environment. However. his results did not give any information as to the amount of M&+ incorporated into the calcite crystal lattice. In fact, of the major divalent cations in seawater Ca”, Mg2+ and S?‘, all can be accommodated in ‘various amounts in the crystal structure of calcite. Several notable attempts have been made tc determine the distribution coefficient of Mg2+ (WINLAND, 1969; KATZ, 1973; F&ZHTBAUERand HARDIE. 1976; BERNER, 1978) and S?’ (HOLLAND et al.. 1964: HOLLAND, 1966; KATZ ez

al., 1972: LORENS, 198 1) in calcites precipitated from solutions containing various concentrations of Caz+, Mp and S?’ ions. The distribution of Mg2’ (or S$‘) in calcite is defined as the Mp (or SF) to Ca2+ molar ratio in the solid divided by their molar ratio in solution. WINLAND (1969) analyzed calcites formed from a direct precipitation experiment. Known amounts of reagent-grade CaC03 and MgC03 were dissolved with hydrochloric acid and diluted to a constant volume. The solutions were cooled to about 10°C and excess sodium bicarbonate was added short of forming a precipitate. The solutions were allowed to warm to about 20°C over a period of eight hours with the slow formation of a CaCO3 precipitate during warming and degassing. WINLAND (1969) noted that x-ray diffraction patterns of the precipitate did not suggest zoning even though the Mp to Ca2+ concentration ratio in the reacting solution was not maintained. He also indicated that the composition of the precipitate was in equilibrium with the final solution since it fitted an homogeneous distribution coefficient equation better than the DOEIZNER-HOSKINS(1925) (or heterogeneous) relation. Results of his experiments indicate that the distribution coefficient was independent of the M8+ and Ca2’ concentration ratio in solution and was equal to 0.0 19 with a mean error of 0.001 at 20°C. KATZ (1973) determined the distribution coefficient of M$’ in calcite over a wide range of temperature, 25 to 90°C. Calcite crystals were grown in a closed system by recrystallization of synthetic and natural aragonite crystals in the presence of various CX12-MgC12 solutions with and without NaCl. By

217

218

A. Mw.5 and J. W. Morse

using this procedure the saturation state of the so-

lution was maintained at 1.5 or less. Unlike WINLAND(19691, KATZ (1973) found that, since the solution composition was allowed to vary, the distribution of Mg2+between the calcite and the solution was heterogeneous, closely following the DoemerHoskins distribution law. His rest&s also indicated that neither an incmase in the absolute Ca*+concentration, nor the pmsence of 0.48 m NaCI in the solution had any et&t on the incorpo&on of Mgr in calcite. Results of his experiments indicate that the distribution coe&ient of Me” is quai to 0.0573 (k.0017 std dev) at 25°C with a strong temperature depe+xxe. FUCHTBAUERand HAIZDIE(1976) precipitated Mg-cakitesby adding Na&a to mixtwes of M#&CaClzof various ionic s&en@, while txying to keep the Mga+ to Caz’ concentration ratio in solution as near to constant as possible during prclcipitation. Their study was also carried out over a wide range of tempemtuma ( 13.5” to So’c). Results indicate that the distribution coeikent of MgzCbetween the Calciteprecipitateandthe*tionwas~ndentof the Mgr+ to Ca*+concentration ratio in sohrtion and the ionic stmngth. The value of the distribution coefficient of Mga+at 25°C extrapolated from. their data is equal to 0.031 (z!AlO5). Comparison of the results of the pnviousty desuibedin~~thusshowa~~of agmement.BERNHR(1978)hassummakedobMrved concentrations of Mg in cakite precipitatedfkomseawater under a variety of conditions which he had IIccumu&tedover the yean and pfevi&y repowd vaiuar BERNER(1975, 1978) thatthese results substantiate the idea that both sohktion chemistry and reaction kinetka infhtenee the concentration of Mg incorpomted into the sohd. This is in agreement with the ~It&&ons reached by THORSTBEtONoILdptUA&ZR(1977, 1978)basedontheir comic saturation model. However, LAFoN (1978) and GARRBLSand WOLLAST(1978) have arguedthatthecompositionofnaturahyoccukngMgcalcites with as much as 12 to 15 mole % MgCOr may be in equilibrium with the sohttioss From which they formed. HOLLAND et a(. ( 1!%i4) and HOLLAND

( 1966) de-

termined the amount of Sr taken up by calcite during the direct precipitation from ammoniacal CaClrSIC& solutions by the slow reti of COr, over a large range of temperature. They found that the distribution coefficient of Si2’ between the calcite and the sohttion decmas& from about 0.14 at 25°C to 0.08 at 100°C. These resuhs were codnt with analyses of cave formations and the waters from which they deposited. KATZ et af. ( 1972) using a procedure similar to the one QTZ (1973) used for the determination of the distribution coefficient of MgZ’, determined the distribution coeikient of S#+ between 40 and 90°C. They found that the amount of Sr incorporated in

the calcite recm from aragonite followed *be heterogeneous distribution law and was essentially unchanged by the pmsencc of NaCI. Their resuhs indicate that the distributioncoe&ie~t of Sr between the calcite and the solution was increased significantly with increasing calcium concentration. They attributed this behavior to an enhancement of the naystalIizotion rate. The vahres of the distribution coefficient at O.Olm Ca** (seawater con~n~tion) are 0.055 at 4O“C and 0.059 at 98°C. himENs ( 198I) investigated the coprecipitation at SP with calcite on calcite seeds using radioactive tracers in CaC&NaCl soiutions at 25°C. The distribution coeBcient of Sr between the calcite overgmwth and the sdution was determined as a functioa of precipitation rates+using a pH-stat system similar to that of MORSE(1974), although not automated. His results indicate a substantial increase in the distribution coefficient with increasing rate of precipitation. The distribution coe&cient varied from 0.03 too.1 w~e~~~~~rn 1 to 12X 10-9mole mg-’ tin-‘. L~RENS (198t) dezno~~~U-S&that his resultswere consistent with the resultsof Katz et al. (1972). Xtis thereforedililcult to amertain if the coprecip nation reactions in these studies approximate equilibrium conditions, and if not, to what extent reaction rate in3uences the distribution of mater& between &id and soiution phases. Since many natural proceases do not occur at equilibrium, the successtkl application of Wwatory derived distribution coefficients requires a knowledge of the signigcance of the in#uence of reaction rates. As for most s&s in so&ion, there normally exists a range of supersaturation which permits growth but where the nucleation rate is nearly zero. This applies to the formation of CaCOl in seawnter and related sohrtions. This featum of the system allows for the sseded precipitation of CaCGr over a wide range of precipitakn ratea. It was possibk to precipitate catcite on calcite seeds tiom seawater and similar solutions, and anaiyxe the composition of the overgrowth. Precipitation experiments were carried out under strictly controlled conditions in an open systern on reagent grade calcite seed powders The effects of ftiors such as Mg2+to Ca” concen~tion ratio in solution, solid to solution ratio, amount of overgrowth, growth rate, stirring rate and Pcch on the composition of overgrowths were investigated at 25’C and one atmosphere total pressure. This was undertaken to determine the distribution coefficient of e and St? in c&&e overgrowths and examine the influence of the presence of Mg2+ in the solid on Srr” uptake. The major goals were to resolve the arguments which have been put forth concerning the e&&t of precipitation rate on the comPosition of the oveqgrowths and elucidate the factors governing the precipitation of Mg-cakites in a nonbiopsy controlled marine en~~nment at ambient temperature and pressum.

Incorporation MATERIALS

of Mg, Sr into calcite

AND METHODS

Solids. solurlons and gases

All CaCO, overgrowth precipitations were carried out with Mallinckrodt brand “Analytical Reagent” grade C&O, as a seed material. This C&O, is greater than 99% calcite as determined by x-ray diffraction. The calcite is present as rhombs of approximately 3 microns m size with a surface area of 0.55 rn’ g-’ of CaC03 as determined by the K.r-BET method of DE KANEL and MORSE (1979). The surface area of the solid was used to calculate the mean overgrowth thickness and the reaction rates. Gulf Stream near-surface seawater was used for all natural seawater expenments. It was collected 5-7 km off the coast of Miami, Florida. The natural seawater was aged for several months and adjusted to 35% salinity according to a method described previously (MORSE er al.. 1980). This procedure was used to minimize biological activity and lower phosphate concentration since PO, is a known calcite precipitation Inhibitor. After aging. the seawater contained undetectable amounts of reactive phosphate (co.1 rg PO,-P/ I ). Artificial and altered Ca*’ and Mg” concentration seawater solutions were prepared to include all major elements of natural seawater including F- according to the method of USTER et al. (1967). modified slightly to fit the analysis of MILLERO (1974). The Ca2’ and Mg*’ concentrations were varied to determine the effect of absolute CZa*’concentrations and Mg” and Ca*+ ratio on the composition of the overgrowths. For all solutions the ionic strength was balanced to equal that of 35% salinity standard seawater (IT = .697 m) by either increasing or reducing the amount of NaCl added to the preparation according to: A[CZa2’]W, + AIMgz’lhcsCI,= 3A[Na+]-

(1)

Precipitation of an overgrowth was carried out in an open system on calcite seeds in solutions of close to constant composition. Constancy of composition was achieved by using a chemo-stat technique which is a modification of the pH-stat methods used by MORSE (I 974) to study CaCO, dissolution kinetics. The working concept of the cbemo-stat technique, which is to maintain a constant-state of disequilibrium. comes from consideration of the precipitation reaction of a given Mg-calcite from a solution (1 - x)ca2+ + xMg2+ + C@- 5 Ca,_,Mg$JO@olid)

(2)

This reaction indicates that CZa2’,Mgz’ and CO:- concentrations must remain invariant if a constant state of disequilibrium is to be maintained. In an open system. the CO:- concentration can be held constant by keeping a,.,(pH) and Pm constant. It is possible to keep the Pcol constant by bubbling a CO2N2 mixture of constant composition through the solution during the precipitation reaction. Most precipitation reactions were carried out at a Pa near 0.310%. The Pm of the gas mixtures was determined by infrared analysis to within a precision of *2’% of the reported value (Air Products Specialty Gas). In order to maintain a constant saturation state, this higher than atmospheric Pco, was chosen for most of the growth experiments. BERNER (1975) indicated that at lower Pco, values, variations in the solution Pco2 could result from an inability of the CO, equilibration to keep up with relatively fast reaction rates which occur at high supersaturations. In slower precipitation rate experiments, atmospheric PCO,was used in several runs. The compositional data was identical with the runs carried out at higher Pcol. Since the gas mixtures were generally waterfree. the gas was bubbled through distilled water before being bubbled through the reacting solution to prevent evaporation in long term runs.

‘19

BERNEP.(1975) maintained the pH of his reacting solutions constant through the use of the carbonate pH-stat described by MORSE (1974) by injecting a NaOH titrant solution to cause the conversion of CO2 to CO:-. In this study. a constant pH was maintained by addition of excess carbonate alkalinity in the form of HCO; and CO:- from a mixed NaHCOr-NazCO1-electrolytes solution. PC@ reequilibration probably requires less time by this method. Calculation of Pm from the steady-state pH and measured carbonate alkalinity values in the reaction solution agreed to within +O.Ol’%atm. Precipitation

reaction

The composition of the reacting solution was maintained by 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 reacting solution plus an excess in calcium and carbonate alkalinity (ACa2’ = %AAc) to compensate for the Mg-calcite precipitation. Detailed composition of the two titrant solutions are given by Muccl ( 198 I). Generally, by “tuning” the carbonate alkalinity of one of the titrant solutions according to the precipitation rate, it was possible to hold the Ca” concentration in the reacting solution to 23% of the initial concentration. No excess Mgz+ was mixed into the titrant solutions since in all cases the amount of Mgz+ incorporated in the overgrowth is very small compared to the amount in the reacting solution. The reaction vessel consisted of a water-jacketed glass vessel with a volume of approximately 450 ml. The temperature of the reacting solution was maintained constant at 25.0 i 0.1 “C by recirculating water from a constant temperature bath. Stirring of the solution was provided by a four-bladed glass propeller-type stirrer, powered by an adjustable electric motor mounted above the reaction vessel. This type of stirring system was used to prevent grinding of the CaCO3 particles against the bottom of the reaction vessel by a magnetic stirrer, which could result in an increase of the reactive surface area and in consequence cause a variation of the reaction rate. Initially, the calcium carbonate ion molal product of the solution was adjusted to the steady state supersaturation at which the precipitation reaction was to be carried out. This was achieved by addition of small amounts of Na2C0, and presaturation of the solution with the COs-N2 mixture (-0.3 10?J0or 340 ppm). Prior to the addition of the seed material. an aliquot of the solution was withdrawn and stored for later comparison of the initial and final concentrations of Ca2+, Mg?, S? and CO:- in solution. The initial volume of the reacting solution was generally approximately 350 ml. A known amount of calcite seed material was then added to the solution and the syringe pump was activated. The volume of titrant solution injected into the reacting solution and the pH were monitored on a dual pen recorder. By varying the amount of seed material and the injection rate, a large (order of magnitude) range of precipitation rates could be achieved. The lower limit of precipitation rate (or IMP) at which the reaction was conducted was dictated by the length of time required to produce enough overgrowth for precise analysis. The upper limit was restricted by the risk of spontaneous nucleation which would create more surface area for precipitation. Preliminary experiments demonstrated that a major problem with the pH-stat system was the magnitude of electrode drift over long periods of time. In most cases, to grow enough Mg-calcite on the seed material the precipitation reaction was carried out over a 2 to 24 hour period. During this long time interval, electrode drift was significant enough to yield questionable kinetic data. To circumvent this difficulty, the controller system was usually bypassed and a constant rate of addition of the titrant solutions was set on the syringe pump. By means of this procedure, shortly after

A. Mucci and J. W. Horse

220 TABLE1

?huurd of

rtolchlovtric in altered

calcite

mnmter &+I

-

solutioos 10.28

after EDTA tiuation of the total divalent canons and subtraction of Ca” and Sr?’ concentrations. The p&on of Me” determinations was approximately 1%. The concennation of Sr’+ in solution was presumai to be 3.0 4 IO-’ kg-’ for all solutions. A few solutions were analyrcd by atomic abaorp&m spcctrophotomcuy before and atIer the precipitation reaction. In all cases the Sti+ concentration did not vary by more than 3% of the initial solution concentration. This d&ation is of t&e same order of magnitude

aolubillty conatante magn~ium

concbatratlon

(IT - 0.697

x 1O-3 oola

kg-‘)

m. at

25’C

asthe&mamdpra&ionoftheaaaiy&.Alkalinitieswen detemCned by the GRAN (1952) method with a precision

K; ( &+I

/ m2+1

1 2.5 5.13 7.5 10.0

1 aol

1n

mole -2 kg -’ x 2.93 3.50 4.39 5.40 6.23

+ T 5 T 5

lOi

.07 .02 .20 .05 .08

of 0.2% using a Na2C0, soluuon for standardizauon of the acid. The second apparent disaocianon constant of the carbonic acid system, X2, for the altered seawater solutions was calculated using the method of BEN-YAAKOV and c&DIURRR (1973). Calculations of the total ion activity cc&k&t of the borate ion, ?#(OH)r), usingthe ion mr( 1982) indicate ing equations of MlLLJiRO and !~UEBER rbar its value does not vary over tbc compoaitron ranges used in this study. The d&o&ion constam of boric acid, K;, dctcrminai by LYMAN(1957) for 35% 5alinity seawater was uaai to calculate boric rid contribution to the tit&on aUlinity. The method most frrquently used 10 express the saturationsaoeofasdutioncwithrrspecttoasolidphPseisas aratiooftheioamolalpfoductoftbcdiPociPtadsaltto the stoichioatattic solubility constant. For CaCO,. the exprrssion for sanuation state. O, is delined as:

where Kg is the quilibrium noichiomeuic volubility ot calcite in the par&u&r solution studied. Tbc stoicbiometric solubility constant of calclte in 35% manta at 2S’C &m&nod by MoRse n al. (1980) was lusl for saturati state c&ulu&s in normad natural and Thesto&hioaI&c sc&tnlity constant syntbuic B. of a&ire in the altered w contention seawater solutions wcm dcrermined by Muca aad MORSE (1983) in a clo&sys&mbytkesamegenualmethodde.saibaIby MOU% et al. ( 1980). over extended squibbration periods. Tbcakitc~sobAb&yaJnstantvaiucsuscd in this study are aasembkd in Table I. Overgrowth composmon

Steady state sdution amqxuiritm and saturaion state Tbc stady state ion molal prodwzt (mc+m&) rcactiqsoIutionfbfagivmrprscipit6tianmtcwncalcuIatat~theotadystatepH~dmingtheupaimen:mccu&lato

of the

t&o&a wae m by EGTA titratio; using a mod&c&m (xc Mvan, 1982) oftbe tocbnique of T’suNoGueroI.(1966).lQrpre&iMoftlla~ was 0.4%. A mam balaafx akad&oo m psrformsd to wYifyiftJlccak&Ilncososatntioointhefiorlsdutionwas awaamtioo aul the amount cqu8ltotbesumoftlKtinithl addadduriqthe expaimsnt minus the u#wIL1 plecipi~aaQLdumarboaate.Tberslrr#lUimtaithatia akYiumamceetr8tioninthc plmor(auMtbeakuhU finals&tionagmadtowi&n2%orbaQaoftbecoacentmtiondcten&&bytiUUion.Thcsumoftbeunccrtainticslrrrtal+tbthemYbahaccakuhtionmorsthan makca up for this smaU dmiation. Md* was determined

The concentrations of Na’. K’, Mt’ and Si’ lncorpomtrd in the overl(rowth and absorption layers. plus the ruidud solution sahs, were detcmined by atomic absorg tion opcctroobatomctry ti acid di(mtion of a known amount ofthc driunpned carbonate material. Blanks and standa& were prqmraj using equivalent amounts of the scedmuerialtoobtainasimilarmawixe&ctThccstimafed preciaions on the atomic absorption analysis were 22% for Mg+, Na’. K’ and *3% for w+. To avoid the dissolution of any of the ovqrowth, the reacted mat&al was not rinsed a!?cr 6itration from the reacting solution. Thaeforc, a con-on from the residual salts to the mauurcd w* and SP concentrations had to be subtnctcd Using Na+ and K’ as tmcers, their concentntion on the reacted material and their dative conccnmuion ratios in g kg-’ in the rcspaXive solutions. it was poaib&to &male tbc r&dual solution salt contribution. WHITE ( 1977, 1978) has invcadgated the copmcipitation of Na’ and K’ in aragonite and Na’ in calcite at 25°C. In-n Of his tiults B Ihat the overgrowth contribution of Na’ ions to tbc analysis is negiigibbz compared to the residual solution salts. WHITE’S ( 1977) study also shows that K’ distribution coe&cient in aragooite is 20 times smaller than Na’. This may indicate that K’ could be an even btttcr tracer of &dual solution salts than Na’ if his results can be extrapobted to calcite. Agreement be-

Incorporation

of Mg, Sr into calcite

twetn corrected Mg2’ and d+ concentrations using Na’ and K’ as tracers was generally better than 3%. This is excellent considering the manipulation of the atomic absorption determinations which already carry a 2 to 3% uncertainty. In all cases, the solution salt contribution derived using Na’ as a tracer was smaller than the one calculated using K+. Based on this consideration and the fact that there is close to 45 times more Na’ than K' in solution making the determination of Na+ easier and more sensitive, the results using Na+ as a tracer were used in all subsequent calculations. It was assumed that Ca*‘, Mg+ and S?+ were the only cations that had a significant concentration in the overgrowth and that their charge was balanced by carbonate ions (they were present as carbonates). The mole fractions of MgC09 and SrC03 in the overgrowth were computed from the corrected Mg2’ and S? concentrations, the amount of carbonate precipitated and the amount of material dissolved for the analysis. A step by step explanation of the calculation procedure is given by Mum1 ( 1981). The dried reacted CaCOJ was also examined by x-ray d&action spectrometry in an attempt to detect the presence of other carbonate mineral phases which might have precipitated along with the Mg-calcite. The calcite peak (20 = 29.45) on the diffraction spectra was clear and sharp since nearly 90% of the sample was composed of reagent calcite seed. In most cases, the Mg-calcite overgrowth was also identified by the presence of a diEaction peak shouldering the calcite peak. Mg-calcites containing less than -8 mole W MgC03 were not clearly seen due to the overlap of the seed material signature. The position of the Mg-calcite peak was used to determine its molal MgCO, content from the idealized GOLLSMITH and GRAF (1958) calibration curve. The Mg content of the MB-calcite overgrowths determined by this procedure was in good agreement with the wet chemical analysis within the analytical uncertainties (+-10%). No other mineral phases such as aragonite or vaterite were detected. If calcite precipitates from a very large solution reservoir, or in this case, a solution reservoir where the composition of the solution was maintained constant to keep up with the precipitation, the distribution coefficient of Mp (or S$‘) between the calcite and the solution can be defined by the Henderson-Kracek (or homogeneous) distribution coefficient: -. D&+=

M&+/M!&+ ML,+/M&+

where Mw+ and Mu+ denote the molar concentrations Mg2’ and Caz+ in the calcite (C) and the solution (L).

(4) of

RESULTS

Selected parameters and results from representative single measurements are presented in Tables IIa through IIg. They include steady state pH and corresponding saturation state of the solution, amount of carbonate material precipitated, the precipitation rate and the calculated distribution coefficient of MgZ’ and S?’ in the calcite overgrowth. The complete data set which includes initial seed weight, precipitation period, the final Ca2+ and M$+ concentrations, carbonate alkalinity of the reacting solution and mole fraction of MgC03 and SrCO3 in the overgrowth, in addition to the parameters presented in this paper, is available upon request or from Muccl (1981). The concentrations of MgC03 and SrCOs incorporated in the overgrowths precipitated from a solution of constant composition are independent of

221

the precipitation rate, over the range of precipitation rates investigated. Precipitation rates were varied by approximately an order of magnitude, while the saturation state varied from a value of 3 to 17 depending on the solution composition. It is, therefore, possible to average data for overgrowths formed from a given solution composition. The initial solution Mg+ to Ca2’ concentration ratio and corresponding averaged overgrowth composition, mole fraction, distribution coefficient and standard deviations are summarized in Table III. The total ion activity coefficients of M$+, Ca2+ and S? in the solution, estimated using the ion pairing equations of MILLERO and SCHREIBER (1982).are not affected significantly by the variations in M$+ and Ca’+ concentrations. The total ionic strength of the solutions were matched by varying NaCl concentration and, the concentration of the major anions, sulfate in particular, was not altered. Consequently, any variation of the distribution coefficients, D&Z+ and D&Z+, from one solution to the other can be attributed to the variation of the concentration or activity of the cations in solution, if calcite is the only growth phase. Mg2’ and Sr2+ coprecipitation

The distribution coefficients of Mg between the solid and the solution for each overgrowth experiment are plotted as a function of the Mp to Ca2+ concentration ratio in solution in Fig. 1. The averaged o’w+ and mole fraction of MgCO3 in the overgrowths are also assembled in Table III. These results indicate that the amount of MgC03 incorporated in the calcite crystal lattice during precipitation is controlled by an exchange equilibrium. The data also indicate that the absolute concentration of Ca2+ or Mp ions in solution does not affect the amount of MgC03 incorporated in the overgrowth, but rather that the overgrowth composition is dominantly controlled by the M$+ to Ca2+ concentration ratio in solution. All solutions were undersaturated with respect to nesquehonite, the most stable hydrated magnesium carbonate (MgC03 does not form directly from low temperature aqueous solutions). Above a Mg2+ to Ca2+ solution concentration ratio of approximately 7.5, the distribution coefficient D&+ becomes nearly constant (see Table III and Fig. 1). This constant value of D&Z+, 0.0123 (20.008 std dev), is reflected by the proportional increase in the mole fraction of MgC03 in the overgrowth with increasing (Mg2’/Ca2’) .wl’nand is probably controlled by the conditions of an equilibrium of chemical potentials between the solid and the solution. However. below the ratio (M~‘/Ca2’),,.. z 7.5, the distribution coefficient increases exponentially with decreasing (Me/Ca2+),,.,. If the mole fraction of MgC03 in the overgrowth is plotted as a function of (M$‘/ CaZ+)lorn(Fig. 2), the MgC03 incorporated above that predicted by D$$z+ for (Me/Ca2’)_,.. 2 7 S, is then represented by the deviation from the line extrapolated to [M$‘lWr, = 0.

79 .._

71 II

JJ 13 14 15 I7 18 19 20 II ?i ‘5 511 64

bXll

u

---

96 97 98 99 loo 101 102 103 104 105 106 107 108

1.839

1.904 7.902 ?.a92

7.879 7.979

7.901

7.907 7.951 7.873

7.870 7.825 7 .a26

7.90: 7.944 7.838

Y.611 1.92

9.93 10.00

7.51 9.93 11.55 9.31 8.66 Y.34 x.44

8.61 7.61

9.7, 11.48 7.97

81

4.23 3.66 4.48 4.11 3.6b 3.79 5.GG 4.49 5.35 3.89 3.34 4.20 4.72 5.12 3.19 5.76 4.58

________(ale

__---_.

PJJ

7.728 7.682 7.736 7.711 7.703 7.697 7.764 7.737 7.780 7.705 7.661 7.718 7.lM 7.777 7.663 7.006 7.741

G

Carbonate

7.00 6.43 7.04 7.27 6.31 --.

8.96 b.65

6.90 6.90 6.66 b.81 7.25 7.46 7.00 7.15 8.25

7.33 7.44 a.52 6.81 6.16 6.63 7.84 7.45 0.36 a.41 7.66 7.52 9.73 a.53 a.03 8.62 6.40

x 104)

Precipirsted

-1

3.459 1.178 3.747 3.604 3.892 3.900 3.464 3.373 1.678 1.646 1.618 j.046 3.471 1.472 3.4hB 1.768 __-

3.560 3.851 3.415 3.55a 3.702 3.706 3.268 3.412 3.127 3.b37 5.a45 3.558 3.334 3.194 4.064 2.981 3.414

(sole hr

.

-2

)

1.69 I.13 1.6b ".__.

i .ba

i.94 1.96

1.57

1.30 1.24

1.35 1.46

1. 39

2.57 2.80 3.19 3.06 2.?2 2.7s 2.62 3.00 2.76 2.61 2.45 2.54 2.92 2.43 2.54

(102)

0.305 O.u)? 0.3oR .__

0.299

0.275 0.24) 0.330 0.317 0.392

0.2t2

0.222

0.220 0.1M 0.398 0.192 0.214 0.193 0.211 0.195 0.220 0.205 0.160 0.187 0.193 0.220 0.150 0.222 0.164

54 60 62 63 68 69 70 76 78

48 49 51 53

exp

I

7.883 7.930 1.852 7.915 7.906 7.890 7.944 7.871 ?.I%1 ?.832 7.911 1.915 7.902

PH

7.809 (7.77) 7.853 7.871

i.818

4.97

7.755

7.177 7.825 7.854 7.799 7.7>3 7.819 ?.tiaO 7.765 7.745 7.792

110

111 112 113 114 115 116 117 118 119 120 121 122 123 124 125 -..

JO.43 LO.44 IO.51

11.32 8.50 10.48 10.43 9.65 11.46 8.86 8.37 7.82

9.82

5.65 6.82 7.57 5.89 5.50 6.46 8.22 5.13 4.51 5.47 6.16 5.99 7.24 7.85

6.14

7.894

109

___ x

7.46

8.04

6.42 6.30 6.62 6.59 6.71 6.84 7.50 7.63 6.70

5.45

7.82

7.24

6.21

7.40

6.74

of

104)

6.60 6.78 5.19 6.53 6.67 6.49 6.82 b.25 6.07 b.7b 7.02 7.1) 1.05

(mole x 104L_

rrecipi‘alrd

Carbonate

Amount

(mole

Precipitated

J.439 0.201

,_.~

1.74 1.6, 1.89 1.81 1.511

3.755 I.903 3.471 J.467 1.477 .-.

1.88 1.65

1.759

1.90

J.615 1.328

L.Ji 2.00 2.07 2.12

(22,

“C 2,

1.66 2.11 2.16 2.18 1.79 2.29 2.26 2.38 2.13 2.17 2.23 2.00 2.09 ___~

0.283 0.260 0.317 0.3lb (1.261,

U.JO1

0.388 0.328 0.396 0.349 0.350

“.J68

0’: 2+ bC

0.234 0.245 0.247 0.239 0.28G 0.270 0.268 0.239 0.299 0.272 0.249 0.173 0.253 0.266

0.195

J.7b 1.84

2.07

3.473

3.362 3.335 3.757 3.474

(sw,c hr-' n?)

-log rete

3.218 3.100

3,072 3.583 3.295 3.155 3.584 3.733 3.439 3.012 3.799 4.043 3.655 3.464 3.553

C9C’O 6TE’0 SEE’0 OLE’O C9E’O O&E’0 YLZ’O ZSE’Q EZE’O LEE’0 SZE’O mc ‘0 EbZ’0 CEE’O

CSE’O OZ.1 02’1 82’1 ES’1 LE.1 ix.1 OZ’T ET’1 oz.1 01’1 02’1: 9E’T VZ’I ZT’T

llZ’I

E8b’E (98’S 09Z’E 892’? ET.9 9LE’E LZL’E 965-E 062’9 L9C.C XL ‘C E99.C LOO’? 119-c 9VI’9 SBS’E

ZO9’C ZSf’E S9V.E TOtl’E TO(I’E LO’I’E LOV’E 9SS’E ZE’I’E ZSS’E 00+/-c v7S’C ZOtl’E

(,OI

La.9 99’S $9’5 9T.P

x WV--

59.r 09’E OC’tl 9Z’P

80Z’8 OllZ.8 OBZ’8 (tOZ’8 LT0’8 8T0’8 KEO’B E+IO’B L90’8 LEO’8 LEO”8 620’8 SOT’8

Z’IT 1’21 6’91 88’6

9’TT E’ZI 9’01 2’71

916’1 926-L 900’8 bLB’1

106-L X6.1 068-L 296’1

99’f 69’9 IL’9 V7’L L9’L 99’9 VG’L ZE’L x-9 66’9 ST’L 9L’S LB’S

(uo~aPlauaouoa

651 8ST LSI 951

L91 991 591 791

E”OT & ‘OK E’OT

C’OT 59’2 59-z E9’Z 99’2 99-t 99’2 $9’2 8Z’S 8t’S

192’0

91’1

ET’1

pa.Taarq

692’0

+zs3

ZLZ’O 261’0 CLZ’O LZZ’O 89i’O 662’0 6fZ’O 661’0

ST”1 60’1 9C’T 91’1 61’1 ST’T ZZ’T ZT’T TE’I EZ’I 02’1 91’1 21’1 %:: T6Z’O 6ZZ’O P9Z’O

910’9 WE-E ZLS’C ZP9’E 06S ‘t T91’9 8Ei.E 698-E ZS9’E Z6S‘E ELB’E TbS’E OZO’ 9 899’E 061 ‘E CCL’E 98s’E

z911.0 L9E.O OOt1’0 TOS’O tos*o LOS’0 96’1’0 SStl’O 097’0 589’0 905’0 LE9’0 SSS’O

(901

x elm

EE’S ET’9 t9.s IC’L tx’f 16’S 68’9 Lb’S Z9’9 9E.9 65’S 10’9 Eb’S tr9.5 91’9 68’11 CL’S

-

*/LT ELT ZLT TLI ciL1 69T 89t L9f 991 S9T 991 9E SE

CL.0 19’El 28’6 f9’ZI LO’ZT 16‘1 20-n L9’6 Ob’ZT ZT’ZT 69’8 28’TT 99’6 99’ZT T9’rr 61’01 9E’TI

OSB’L ELb’L Sb8’f 996-L bC6.i 9PE’L CZb’L 268-L 09b’L 6E6’L 006’1 LEb’L LLB’L 056’1 086’1 Zlb’L 1Eb’L

Zfll T1T 091 bC1 BE1 LET 9EI SET %I EC1 7x1 XI OCI 621 EL1 LZI 921

224

A. Mucci and J. W. Morse TABLE

iII

l'he inlrlal soliltion and corresponding averaged overgrowth compos~tlon efOm the seeded precipitation of calcite at 2S'C

Mole $ tions

XqCO,

D&Z+

%31e ? SrCO,

cl02,

-

a

0

la

:7 ,., 2; i. 7 17 z.5 7 L 4

2.5; 5.1 5.1: i.5 10.3: lEc 2oc

2.‘42.22 2.08r.18 1.632.20 3 892.16 ;:2ot.o7 :.23C.O8 1.S42.17 ;.12~.07 L.35t.25

2.7'0.2 4.9t0.4 7.7t0.9 5.8?0.8 3.220.5 11.2?0.: 7.3iO.8 11.9'0.6 21.Ot3.9

.195=.023 .246=.032 .29?i.O40 .326t.040 .270t.020 .344t.a24

.I_2.02 .20t.03 .24*.03 26i.03 ‘2Zi.02 .27'.02

-1 'Synthetic seavater, iCa2+. , = 10.23 x 1o-3 mle kg SOl'il bAged Gulf Stream seawater -1 'Synthetic seawater, iCa2+ 1 = 5.i4 x ?Om3 Role kq sol'n rvalucas are standard deviations

In ail the precipitation

the S? eon3-4% of its initial concentnuion of 9.0 X IO+ mole kg-’ solution. Also, the solutions were always undersaturated with respect to strontianite. As in the case of the MgCO:, incorporation, the amount of SrCOj incorporated in the calcite overgrowths was found to be independent of the precipitation rate, within the range investigated. This finding is in dkgrwment with the results presented by L.OREN~( 198 1.) who found a strong increase in the uptake of Sti+ with increwing rate of precipitation on calcite seeds in M$+ free solutions. centration

was maintained

experiments, to within

FIG. 1. The distribution coefficient of magwsium, of various magnesium to calcium ratios.

solutions

The average coefficient, D&z+, and mole ‘4~of SrCO, determined from each solution composition investigated in this study are assembled in TabIe XII. Results of this study indicate that the measured distribution coefficient of Sti’ between the calcite overgrowth and the soiutios at a given !St’+/Ca2’fti, is influenced by the amount of MgCOX incorporated in the overgrowth. This trend is better observed in a plot of L&Z+versa % xbnsCch, in Fig. 2. The variation of L&+ is directly proportional to the increase of the MgC03 content of the overgrowth. Extrapolation of the data to a Mg-free calcite yie!ds a @$I+

D&A+1 in calcite overgrowths precipitated

fram

ln~~~tion

of Mg, Sr into calcite

Holland

(f966)

n Koir $al_ (I9721

0 This study

%)c

M9CO3

FIG. 2. The distribution of strontium, D$+, in calcite overgrowths as a function of the MgC03 content.

at 25°C equal to 0.146 which is the by HOLLAND (1966). The variation tion coefficient of S?+ in M&calcite cipitated from a solution having = 8.75 X lob3 can then be described equation:

D& dete~in~ of the distribuovergrowths prea (S?+/Ca2+),, by the following

D&l+ = 0.146 + I .833xr@cQ

(5)

from seven data points (70 independent measurements) ranging from ~~03 = 0 to 12%. Prec~~~tut~~n kinetics A ciassieal approach to reaction kinetics is to determine the empirical order of a reaction and its rate constant. The form of equation which has been used most often to describe reaction kinetics of calcite dissolution (MORSE, 1978; MORSE and BERNER, 1979) and precipitation (NANCOLLAS and REDDY, 197 1; REDDY and NANCOLLAS, 197 1; WIECHER~ et al., 1975) is: R = k(Q - 1)

(6)

where R is the precipitation rate normalized to the reacting surface area (mole hr-’ me*), k is the rate constant, n the empirical reaction order and (Q - i ) can be considered as the degree of su~~tu~tion or ~~ui~ib~urn which accounts for the variation of the stoichiometric solubiiity of calcite in the solution studied. In the logarithmic form this equation becomes: logR=logk+nlog(Q-

1)

(7)

The value of the rate constant and empirical reaction order can he determined by the intercept and slope, respectively, of a plot of log R versus log (Q - 1) since the relation is hnear. The data for the precipitation of Mg-calcites on reagent grade calcite

seed powders in the various Mg” to Ca’” concentration ratio seawater solutions have been processed in this manner. The vaiue of log k and n determined from the least squares fit of the data of each solution composition are gathered for comparison in Table IV. Typical results for the precipitation kinetics of calcite from natural and synthetic seawater are presented in Fig. 3. A slight variation in the rate constant between these two solutions of the same Mg2”/Ca2’ ratio could be explained by the presence of phosphate ions in the synthetic seawater. The source of the phosphate is the reagent grade chemicals used to prepare the solution. The computed reaction orders for the precipitation of a Mg-caicite in each one of the solutions are similar, ranging from 3.07 to 3.70. The data indicate that there is no marked trend of the variation of the reaction order with solution composition. However, if one of the data points is ignored (n = 3.5 at M$+/ Ca2’ = 2.5) an increase of the reaction order with increasing Mgr+ to Ca*’ concentration ratio in solution is observed. The increase in the reaction order may possibly be attributed to the increased Me concentration dependence on the precipitation of a higher Mg-calcite. It has also been observed (MORSE and BE~NER, 1979; WALTER er al., 1979) that the reaction order for the precipitation of calcite and aragonite from natural seawater is noticeably increased by the presence of inorganic calcium carbonate precipitation inhibitors such as PO?- ions. Mti” ions may behave in a similar manner. The kinetic data indicate that there is a strong negative linear relation (see Table IV and Fig. 4) between log k and the M2’ to Ca2+ concentration ratio in solution. Since all solutions for which the kinetics of calcite precipitation were investigated contained the same concentration of Ca2+ (- 10.28 X 10s3 mole

226

A. Mu& and J. W. Morse TABLE IV

Kinetic data ir, the precipitation of calcite overgrowths on reagent grade ealclte seeds in various solutions Log k (bfg2+l/tC+2+l)sol’na

"

Rd

0 1

-5.09

5.13b 2.5

-6.62 -5.85

3.07 3.31 3.11

.9931 ,987s .9936

5.i3c 7.5 10.3

-6.53 -6.98 -7.78

3.13 3.30 3.70

.9871 .9855 .9933

a[Ca*+lSol,n

-1

= 10.28 x 1o-3 mole kg

solution

bAged Gulf Stream seewater 'Syntheticseawater d Correlationcoefficientof the log rate versus log (n-1)least squaresfit

The e* of several experimental factors on the composition of calcite 03qmwth precipitated iiom a given sohtion have been in . Apart from the pmcipithon rate which was the sin&2 most importolatplanmetatoexamine,the@kctofP~,soIid to sok&ion ratio, stirring rate and Bmount of overgrowth were also examined. With the exception of three experimcn& aii preciphation reactions were CarrjGd out at the same Pa, 0.310% or lo+’ atm. The experiments were do~ataP~~~to~~~cC~~

pressure. Air pumped from outside the laborator was bubbled into the reacting solution. These three overgrowth experiments were done in aged seawater at 25°C. The analysis of the overgrowths precipitated during these three experiments yields a fl’&+ equai to 0.0160 2 mean error .0007 which is equivalent 10 the results of the other experiments conducted in the same solution, I&Z+ = 0.0163 (t.0020 std dev), in which the PCO, was maintained at 0.3 10%. The e&et of the solid to soiution ratio on the composition of the calcite overgrowths was investigated in ail the solutions studied. The sotid to solution ratio was varied in the reacting solution by changing the amount of initial seed materid at the beginning of the precipitation reaction. The range of solid to solution ratio examined varied from 0.0020 to 0.0047. Over the range investigated the solid to solution ratio did not exhibit any marked influence on the composition of the calcite overgrowths, L&a+ or D&z+, The solid to solution ratio did not affect the precig itation rates which were found to be equivalent at a given saturation state when normalized ta the reactive surf&~ area which is calcuh~ted from the initial seed weight and the krypton BET surface area det~i~~ons. The reacting solution was stirred by a propellertype w stirrer, powered by a variable speed motor. The stirring was necessary to keep the seed material in suspensian so that all surfaces were exposed to precipitation and to assure mixing of the injected &rants to prevent locaked spontaneous nucleation near the injection ports. The stirring rate was varied over a range between 60 and 300 rpm. The stirring rates for each precipitation reaction were not reported here since no variation in the overgrowth composition or caption rate was observed. DE BOER (~9~7)~s~~~~~nto~~~ speci& surfWe areas could occur by a process of crys-

c

-3.000 -

-a-

Aged

-o-

synthetic

Gulf

Strram

aeauator

seawater

FIG. 3. Kinetics of~recipitation of a msga&um calcite from a natural and synthetic seawater @,$$+j Ca2+ = 5.13) at 25T.

Incorporation of Mg, Sr into calcite

eT

210.26

227

x 10m3 mole

kg-’ 8w

F * sg

-6.50-

FIG. 4. The effect of solution composition, (Mg2+/Ca2+)mI.n, on the rate constant for the precipitation of various magnesium calcites at 25°C. tal breeding during stirring. This would cause an increase in the surface area available for precipitation and hence an increase in the precipitation rate. This phenomena was not observed in this study as the precipitation rate was constant during each run and the steady-state pH did not vary. NANCOLLAS and REDDY ( 197 I), WIECHERS ef al. ( 1975) and LORENS (198 1) also reported that stirring rate did not influence the CaC03 precipitation rate in their studies. In all the precipitation experiments reported, 3.6 to 9.8 X 10T4 mole of carbonate was precipitated on the seed material. Within this range, no observable variation of L& or LI$$was detected. However, several experiments were carried out where considerably less (< 1 X 10e4 mole) overgrowth was precipitated. Result of the analysis of these smaller quantities of overgrowth indicated that the distribution coefficient of MgZ+ between the solid and the solution increased as the amount of overgrowth precipitated on the seed material was decreased. The observed trend can be explained by considering the fact that the reacting material was not rinsed after being filtered out of the reacting solution. The proposed explanation is consistent with findings from this study and previous studies (BR;~;TTERer al., 1972: MILLER, 1973) which indicate that in the solutions studied, absorbed layers on the surface of calcite contain a larger M$+ to Ca’+ ratio than the overgrowth layers. As less overgrowth is precipitated on a given amount of seed material. the contribution of the adsorbed layers to the M$ analysis becomes more significant. The same effect should be observed if the solid to solution ratio is increased significantly while the total amount of overgrowth precipitated is the same. However, if enough overgrowth is precipitated the contribution of M$* from the adsorbed layers becomes negligible and the calculated D&Z+ are representative of the overgrowth only.

Dl!XUS!SION ikig2+ and Sr2+ incorporation in overgrowths The results obtained for the incorporation of M$ into calcite overgrowths from solutions with a Mg:Ca ratio less than 7.5 suggest that surface related phenomena may be important in determining overgrowth compositions. The most likely process involved is adsorption, which could significantly aher the Mg:Ca ratio at the solid-solution interface. MILLER and his associates (B~~I-~ER et al., 1972; MILLER and WERR, 1972; MILLER, 1973; MILLER and SASTRI,1973, 1974; MILLER and PAREKH, ! 975: MILLER and RAJAGOPALAN, 1976) have conducted a number of investigations on the surface exchange of 4%Ia2’ and Me with calcite in NaClCX12-MgCl* and artificial seawater solutions. Their work on Mgr+ sorption on calcite, especially the studies done by B~~I-I-ER etal.(1972), MILLER (1973) and MILLER and PAREKH (1975), has been chosen for special attention based on such factors as: 1) similarity of the material used, compared to this study; 2) the wide range of Me to Ca2* solution concentration ratios investigated: 3) the exhaustive amount of research done on the subject by the same group over the years; 4) the fact that their results exhibit the characteristics of a classical adsorption isotherm. MILLER and PAREKH (1975) demonstrated that the exchange isotherms on reagent grade calcite powder conducted with and without 3.5% w/v NaCl were almost identical. It is reasonable, based on these observations, to assume that the adsorption behavior of Mg2+ and Ca2’ on calcite surfaces in seawater and related solutions would also be similar. The data of BGTTER et al. ( 1972) and MILLER (1973) is pre-

A. Mucci aad

228 TABLE The

surface

calcite

of

in

or

solutions

a&orbed of

v Layer

varying

composition 2+

Hg

and the crlculatcd distfibucim cosfficfent 2+ cm the surface. (From Grllttcr ec&., np Mm1er.

fIng2+llIca2+1),l,,

of

colwmltrarion 1972;

1973)

(IMg241/lC*2+1) surf

= 2+ ‘kg

0

0 .3

0.64

0.64

2 T .4 3? -5

1.05 1.06

0.525 0.3f3

L +

1.1

0.167

871 15 + 2 25 T 3

L.4 1.1 1.1

0.175 0,073 0.044

37 ‘i: 4 54 i 5

1.4 1.6

0.037 0.030

671

sented in Table V. It was use4 to calculate a distribution coefkient on the surface of calcite, AC&+. according to Eqn. (8),

where A’&++is the distrhtion coefhcient of Mgz+ in the adsorbed or surf&e layers of cakite. A plot of A&z+ versus the w+ to W” co%Wentmtum ratio in soiution (Fs 5) tewsis that,at lower(MgZ+/Ca2+~.. thealcite stuBiceadsort%,on 8 relativesale, more Me?’ ions than at h@er (M~/Ca2’)ti. ratios. UsingFig 5. the dimibutian czwffkicnt of M$ in the adsorbed or s&ice layers, A&+, was estimated at the same M$+ to Ca2* solution concentration ratio at which the h¶$+ distl$bution Sxl&kient in the overgrowth D&Z+, was determined. The values, L%&+ and A&Z+ at the corresponding (Mgr+/

FIG. 5. The Wbution co&Sent mapuium to arlcium conantration MOLLER (1973).

J. W. Morse Ca”+~.,, are gathered in T&e VI. A pbt of D&D versus the corresponding AC&+ such as in Fig. 6 exhibits a linear correlation between these values (R = .997). This hiding indicates that. as M&LER (1973) had suggmed previously, the surface or adsorbed layers may act as matrices for the growth of mixed carbonate phases. Results presented in Fig. 7 suggest that the influence of the adsorbed layers an the incorporation of M$+ in the overgrowth is more important below a (Mgfc/Ca2+)~, of 27.5. This correlated with the reiative enrichment of M$+ in the surface or adsorbed layers of calcite displayed by the variation of A&Z+ (Fig. 3) Wow the Mg’+ to ca”” solution concentration ratio of 7.5. Above a (MgZ+/ Ca2+)&, 2 7.5, where the reiative composition of the &orb& layers does not change significantly, the incorporation of Me ions in the overgrowth follows the classical thermodynamic behavior characterized by a constant distribution coefkient. Although the relation fxtween adsorption of Mt$ on the surface of calcite and its incorporation in the crystal lattice can be clearly demonstrated, it is difficuh to determine what reaction mecha&ns are invoked. The simplest case can be made for situations where the “fore@” ion is present at low concentrations. under these conditions the adsorption isotherm is generally close to linear, as only a very small Percentage of surface adsorption sites an2 involved. Then (91 where Mt and Mz represent the adsorbate at low concentrations and the major solid component it can replace, mapectively. and A* is the proportionality constant for adsorption in the linear region oniy.

of magnaium, AC&+,on the surface of caicite as a function of the ratio of the solution, according to B&~R

et al. (1972)

and

Incorporation of Mg, Sr into calcite

stant. This is because A* would decrease linearly as

TABLE VI The dxstributlon coefficient of Ng2* in ttie a&orbed or surface layers estunated frorr Flgu e 9 ant! the distribution coefflclent of Ng S-+ in the overgrowth precipitated from the various solutions investigated

mg2+l:~Ca2+1)sol,n A.&z+

D&2+

(x 10')

0 1

.64

2.5 5.13 7.5 10.3 10.0 20

.41 .23 .14 .11 .11 .06

229

2.74 + 2.08 7 1.70: 1.20 T 1.23 7 1.32 5 1.35 5

.23 .18 .23 .07 .08 .O? .25

+ -

1.24 -08

It is also probable that a constant distribution coefficient will exist between the adsorbed surface layer and the solid forming from it, when only trace amounts of the “foreign” ion are being incorporated. Then

where Awl is the distribution coefficient between the surface and the solid for M,. The observed distribution coefficient f&J between the liquid and the solid is then simply Dy, = ‘4*X&f,

(11)

In this study we found that this general approach worked well even for the more complex situation where A* was not a constant, as the piateau region of a Langmuir type adsorption isothetm was ap proached. Once the plateau is reached, Dy, would be expected to decrease if the relation between the adsorbed surface layer and the growing solid Ant, is con-

remained

constant.

Our observation is that this linear decrease in DM, does not occur and in fact Dw, is constant as

(M1 M1

increases above Mg:Ca > 7.5. The fact that

Iiq

I&, is constant above a solution MgCa ratio > 7.5 demands that either MgCa ratio on the surface must change or that Au,, the distribution coefficient between the surface and growth layer, must vary with solution composition. The previously cited work of M6ller and his associates indicates that for calcite the surface MgCa ratio, at solution Mg:Ca ratios > 7.5 to the maximum ratio used in this study, should be close to constant at about a value of 1. It is also difficult to envision a simple mechanism by which the changing solution Mg:Ca ratio should cause a major change in X$&W. The only major variable not considered is the composition of the solid. For minor (< 1%) changes in soiid composition it’s reasonable to expect that there would be little resultant variability in adsorption affinities or partitioning between the surface and growth layer. However, in this study Mg-calcites with up to 2 1 mole 9%MgC03 were produced. It is possible that these major changes in the composition of the calcite result in changes in the surface affinity for Me and/or the surface-solid distribution coeficient. There are two observations that lend support to this hypothesis. The first is that BERM% (1966) found for sediments containing natural Mg-cahAes from Florida Bay, no limiting M&a ratio on the surface and much higher MgXa ratios than MiiIler and his associates found for pure calcite. Although Bemer’s results may have involved some dissolution and pre-

-

-

0.10

0.20

0.40

0.60

A;,,+ FIG. 6. The relation between L&Z+and A&+ as a function of the magnesium to calcium concentration ratio in solution.

.A. Mucci and J. W. Morse

seawater

solution

0

natural

0

synthetic

seaw-ster

solutions

x

svnthetic

seawater

solutions

, nole

44

-i

FtG. 7. The mole IiWdon ofMgCC$ in c&c&eovergrowths as a fimction of the magnesium to calcium concentration mtio of the so&ions firorn which they precipitated.

cipitation of solid, they are still strongIy indicative that Mg&tea may indeed have a higher al?inity for Mg2’ on their surtks than c&&e. The second observatin is that the distribution czoefficientfor Srr+ in calcite irmreams with mcm%siRg Mgcos content of the calcite (this study). This indkxues that the incorporation of major ~n~~~a~ of MgCOs caicite can change distribution coefficients. However, it would be msonabie to expect a decrease in Xc&+ with increasing MgCOj content, not the increase which would be demanded if the suribce composition were constant. It is also difhcuit to imagine why h&+ should change in response to changing solution composition when eEect.ively blanketed from the solution by a surface of constant composition. In summary, we feel that the following is the most probable explanation for observed behavior of D&h. In solutions with an Mg:Ca ratio < 7.5 the greater afimity of calcite, and the Mg-cakzites formed from these solutions, for Mg2+ relative to Ca2* Ieads to a much faster rise in the surface MgCa ratio than that in solution. This results in a non-constant and higher L&a+ than found at higher Mg:Ca solution ratios. In sdutions with these higher h4gCa ratios an isotherm plateau is not reached as found for adsorption on pure calcite. but rather the surface Mg:Ca ratio slowly increases as a result of adsorption on Mgcalcites with a higher affinity for Mg:‘. A “feedback” relationship between the growing solid and solution

is thus set up through the interface, in which a Mgcalcite adsorbs more M$+ permitting the formation of higher Mg-caicite which adsorbs yet more Mg’+ until an equilibrium is established between the solution, interface and growing Mg-caicite. This interaction probabiy takes place during the formation of the first few atomic layers of growth. aRer whit% a homogeneous solid is produced. The results for S$’ incorporation in the Mg-caicite overgrowths indicated a dependence of L&Z+on the MgCOll content of the overgrowth. A possible expianation for this behavior lies with the different size of the cations involved, Ca*+, Mg2+, St?’ and the distortion they produce in the crystal lattice of calcite. Of the three cations the Mg2+ ion has the smallest ionic radius (0.65 a; followed by the Ca’- ion (0.99 A) and the Si’ ion ( I. t3 ‘41. The existence of the tat&e distortion was demonstrated by ~OLDSMITI? es al. ( 196 1) who measured lattice constants of synthetic Mg-caicite shorter than those for calcite. It was aiso found that the c axis dimension is more affected than the a axis, supporting FOLK’Sf 1974) contention that the primary effect of Mg” incorporatton is to shorten the c-axis dimension. Incorporation of a M$+ ion in the c-axis would cause the CO? layers to be pulled closer together and produce a deformation of the lattice. near the Mg. The deformation would create a site above it where a cation larger than Ca2+, such as S?, could be more easily accommo-

231

Incorporation of Mg. Sr into calcite dated. As more MgC03 is incorporated into the crystal kittke of calcite, the number of sites available for S?+ ions to settle in is increased and this is reflected by the huger amount of SrC03 incorporated. Precipitation kinetics The inhibitory effect of Mg?+ on the calcite precipitation rate demonstrated by the large decrease of the rate constant is significantly greater than the effect of ion-pairing of Mg2+ with CO:- ions, since the kinetic data was expressed as a measure of disequilibrium (Q -’ l), which takes into account the variation of the stoichiometric solubility. The effect on the rate constant is dominantly the result of the inhibition of the precipitation reaction at the surface of the crystal. PYTKOWICZ (1965, 1973) reported the lack of any effect of Mg2’ upon the rate of seeded aragonite precipitation and postulated that if M$+ retards the induction period for the nucleation of carbonate it must do so by inhibiting aragonite nucleation and not crystal growth. This contrasts with the calcite precipitation behavior which has been shown to be greatly influenced by the presence of M2+ in solution. A possible explanation lies with the reported observations that Mg2’ ions are preferentially adsorbed on the surface of calcite compared to aragonite (e.g. DE GROOT and DUYVE, 1966; BERNER, 1966; MILLER and PAREKH, 1975). The growth inhibition can then be attributed to the impingement of hydrated Mg2 ions (Mg(H20)$‘) on the crystal lattice of the seed material on active growth sites, such as kinks. It would thereby inhibit the spreading of monemolecular steps on the crystal surface. REDDY and WANG (1980) have recently repotted a similar finding for the inhibition of seeded calcite growth by Mp in dilute solutions. Their results indicated that the change in the precipitation rate constant was consistent with the hypothesis of inhibition by adsorbed Mg’+ if a Langmuir type adsorption isotherm was assumed. A curious finding of their research was that no coprecipitation of MgC03 occur at Mp concentrations of less than 2 X 10e4 M. No attempt was made to determine the composition of Mg-calcites formed at higher M?+ concentrations, although they reported significant changes in solution M2+ content, which is interpreted as indicating the formation of Mg-calcites. This process of inhibition is likely as a consequence of the smaller size, higher charge density and resullant stronger hydration of Mg2’ in comparison to Ca2+ ions. NANCOLLASand PURDIE (1964) indicated that the dehydration of reactants on the surface of a growing crystal is often a rate controlling step in crystallization processes. DE BOER (1977) found a strong positive correlation between the growth rate of calcite from a solution containing M$+ ions at a given saturation state and the Ca2’ concentration in the precipitating solution. He attributed this finding to the faster dehydration rate of Cai* ions adsorbed

on the point, growth longer growth

calcite surface. Thus, from a statistical standa strongly hydrated M2+ ion at an active site will remain in position for a relatively time than a hydrated Ca2+ ion, preventing until dehydration and final incorporation. CONCLUSIONS

Results of this study have served to demonstrate that the amount of MgCO3 and SrCOs incorporated in calcite overgrowths precipitated at 25°C from seawater and related solutions is independent of the saturation state of the solution or the precipitation rate, over a wide range. These results support the hypothesis that the Mg-calcite composition determined during the precipitation from a given solution is representative of the composition of the solubility controlling phase. The composition of the calcite overgrowths is determined by the Mg2’ to Ca2+ concentration ratio in the solution from which they precipitated. The amount of MgC03 incorporated in the overgrowths is also influenced by the surface or adsorbed layers. The composition of the adsorbed layers on calcite is controlled by the solution composition. The influence of the adsorbed layers on the incorporation of MgC03 in the overgrowth is greater at (w/Ca2+)=r. < 7.5. This correlates with the relative Mg2+ enrichment on the surface of calcite below that ratio. The concentration of St-CO, incorporated in the overgrowth is dependent on the MgCOs content of the precipitate. The variation of L&a+ is proportional to the mole fraction of MgC03 in the calcite overgrowths. The variation may be the result of the creation of cationic sites larger than Ca2+ by the distortion of the crystal lattice produced by the incorpo ration of the smaller Mg2* cations. It is conceivable that the incorporation of SrCO3 in the calcite crystal lattice may serve to relieve some of the stress caused by the incorporation of Me and consequently partially compensate for the destabilizing effect of M$* incorporation. However, the amount of SrCO3 incorporated in the calcite overgrowths is so much smaller than the amount of MgCO, incorporated that its presence does not a&t the x-ray di5?action speetra of the Mg-calcite since results of the atomic absorption analysis and the x-ray analysis are in very good agreement. The kinetic data for the precipitation of the various Mg-calcite indicate that M$+ strongly inhibits precipitation and the growth mechanism involves the adsorption of the cations on the surface of calcite prior to dehydration and final incorporation. They also indicate that the dehydration of the cations at the surface may be the rate controlling step. The results of this study have potential application to the study of carbonate sediments and their diagenesis. Most significantly they indicate that the composition of Mg-calcite cements may dominantly be influenced by the Mg to Ca ratio of the solution from

232

A. Mucci and J. W. Morse

which they precipitate and that I&+ is dependent on

the MgC03 content of the calcite. Acknowledgements-We thank Drs. Garrett Bras-s,-Irunes Carpenter, Frank M&o, Giite oatlund and Claes Rcoth for their many help&J sugger&ons and comments, and Sara Sotolongo for her assismnce in the analytical work. The authors also wish to acknowledge Amatai Katz, Robert Lorens and Michael Reddy for ctitically reviewing this paper. Thisresmrchwas~bytheNationaiScienaFoundation Marine Chemistry Program Grant 0CE79-19242.

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