Manganese Uptake During Calcite Precipitation from Seawater: Conditions Leading to the Formation of a Pseudokutnahorite

Gewhimica et Cosmoehimic~ Acta Vol.52, pp. 1859-1868 0016-7037/88/$3.00 + .OO Copyright Q 1988 Pcrgamon Press plc.printed in U.S.A.

Manganese uptake during calcite precipitation from seawater: Conditions leading to the formation of a pseudokutnahorite

ALFONSOMUCCI Department of Geological Sciences, McGill University, 3450 University Street, Montreal, Quebec, Canada, H3A 2A7

(Received September 17, 1987; cicceptedin revisedform April 6, 1988)

Abstract-Manganoan magnesian calcite overgrowths were precipitated from artificial seawater at 25°C on inure calcite seeds using a constant disequilibrium technique. The coinposition of the overgrowths, and more specifically their Mn, Mg, Na and Sr content, was determined as a function of the precipitation rate and Mn*+ concentration in the parent solution. X-ray di!Traction patterns indicate that the overgrowths produced were one-phase multicomponent solid solutions, and contained up to 40 mole% MnCO,. The amount of Mn coprecipitated with calcite decreased with increasing precipitation rate. Thus, overgrowths are not in exchange equilibrium with the solutions from which they precipitated. A kinetic model is proposed which adequately describes the composition of the overgrowths in terms of the relative precipitation rate of an 8- 10 mole% magnesian calcite and a “pseudokutnahorite” from seawater. The possible existence of a pseudokutnahorite in marine sediments and its implications are discussed. The concentration of Mg in the overgrowths decreased with increasing MnCOp content, but their Mg:Ca ratio remained nearly constant. Strontium and Na incorporation was strongly dependent on the number of available non- lattices sites.

INTRODUCTION positions lying below the solvus observed at high temperature were obviously metastable, but found them to reinain un- IT HAS BEEN SUGGESTED(EMERSON et al., 1980; SAYLES, changed in contact with their supematant liquid for at least 1981, 1985; BOYLE, 1983; DE LANGE, 1986) that the for- six months. Furthermore, DE CAPITANI and PETERS (198 1) mation of an authigenic mixed Mn-Mg-Ca carbonate phase suggested that it was probably impossible to demix metastable may explain the observed supersaturation of deep-sea car- (Ca, Mn)C03 solid solutions under laboratory conditions bonate-rich sediment pore waters with respect to calcite in because the process was too slow either due to high activation the reduced manganese zone. If this is true, it may have a energies or low diffusion rates. As is often observed, theoretical significant influence on the accumulation of calcite and the modeling of solid solutions fails short of predicting their actual diagenesis of carbonate-rich deep-sea sediments. behavior in natural systems because of kinetic restrictions. In marine sediments, MnC03 has been found to occur in Estimates of the distribution coefficient of Mn( II) in calcite solid solution with calcite up to 50 mole percent (LYNN and have been derived from field and laboratory measurements. BONATTI, 1965; CALVERTand PRICE,1970; PEDERSENand Observed and selected values range between approximately PRICE,1982). BOYLE (1983) has observed that the Mn/Ca 2.5 and 20 (BODINE et al., 1965; MICHARD, 1968; ICHIKUNI, ratio of foraminifera tests increases significantly below the 1973; KUMAGAI, 1978; PINGITORE, 1978; TEN HAVE and manganese redox boundary. He argued that the increases are HELJNEN,1985 ) , but these studies give little indication of the due to the formation of manganese carbonate overgrowths compositional range over which the solid solution can form. and that manganese carbonate coatings may be a significant Furthermore, L~RENS (198 1) and PINC~ITOREet al. (1988) sink of manganese in deep-sea sediments. Likewise, MICHARD found that the distribution coefficient of Mn(II) in calcite (197 1) and THOMSONer al. (1986) noted the importance of varies with the precipitation rate. They observed that the dis- Mn( II) adsorption on calcium carbonate surfaces to the dif- tribution coefficient decreases with increasing precipitation fusive flux of Mn( II) from anoxic sediments and it has been rate. The value determined by L~RENS (198 1) at the slowest suggested that the process may be responsible for the scarcity precipitation rate (i.e. close to calcite saturation) is approx- of manganese nodules in calcareous sediments (PIPER and imately 50. Unfortunately, most of these studies were con- WILLIAMSON,1977; BOYLE, 1983 ) . The formation of mixed ducted in solutions whose compositions differ significantly Mn-Ca carbonates has also been inferred or demonstrated from seawater and may not be applicable directly to a seawater (MANHEIM, 1961; SUES& 1979; ELDERFIELDet al., 1981; system. In fact, FRANKLINand MORSE (1983) obbrved that FRANKLINand MORSE, 1983) to explain the apparent non- the adsorption behavior of Mn( II) on calcite in dilute solu- equilibrium behavior of rhodochrosite in marine sediments. tions and in seawater is clearly different. They demonstrated Recently, in a theoretical study of the system CaC03- that this distinction was most likely due to the presence of MnCOs-H20, MIDDELBURGet al. (1987) indicated that the magnesium ions in seawater. carbonates should form solid solutions over a very limited In view of the discrepancies which exist in the literature compositional range. However, GOLDSMITH and GRAF and the questionable applicability of previous laboratory (1957) and FUBINI and STONE(1983) were successful in pre- studies to a seawater system, the factors governing Mn (II) cipitating a complete series of well-crystallized solid solutions incorporation in calcite must be clarified before we can draw between calcite and rhodochrosite at room temperature. any conclusions concerning the influence of Mn( II) on the GOLDSMITHand GRAF ( 1957 ) noted that at least those com- solubility behavior of calcite in seawater. In this paper I pre- 1859 1860 A, Mucci sent and discuss the results of a laboratory study on the in- determined by MUCCI (1983) in S = 35 seawater at 25°C (4.39 corporation of Mn( II) in magnesian calcite overgrowths pre- X IO-’ mole* kg-’ SW) was used to calculate the saturation state of cipitated from seawater at 25°C. the solution with respect to calcite, aa defined by:

MATERIALS AND METHODS where K,* is the equilibrium stoichiometric solubility of calcite in All the calcite overgrowth precipitations were carried out using seawater. Baker “Instra-analyzed flux reagent” grade calcium carbonate as a Saturation calculations using pH measurements done on both scales seed material. This material was washed in deionized distilled water and the appropriate constants agreed in most cases to within a few and size separated (3-7 pm) by settling. The CaCQ was freeze percent or better. Results presented in this paper were calculated dried, X-rayed (>99% calcite) and its surface area (0.52 m2 g-i) was from pH measurements based on the NBS scale, as they may be determined by the Kr-BET method of DE KANELand MORSE(1979). more consistent when used with the stoichiometric solubility constant Aged artificial seawater was used for all the experiments. The ar- of calcite determined by MIJCCI(1983). tificial seawater of salinity 35 was prepared to include all major ele- The SrZf and Mn2+ concentrations ofthe solutions were measured ments of natural seawater including F- according to the method of before and after the reaction by flame atomic absorption spectro- KESTER ef al. (i967), mod&d to fit the analysis Of MILLER0(1974). photometry (AAS) using aqueous standards in a seawater matrix for Mn(II) was added to the artificial seawater solution prior to each ~libmtion. Precision of the AAS analysis is estimated to be better experiment in the form of a concentrated (2000 ppm) Mn(II) so than 3% for Mn2+ and 5% for Sr”. lution. This concentrated solution was prepared by dissolution of MnClz .4HrO in artificial seawater. Four sets of experiments were conducted, corresponding to the following initial Mn( II) concentra- Overgrowthcomposition tions in the precipitating solution; 1,5, 10 and 25 ppm. The manganese, magnesium, strontium and sodium content of Precipitation reactions were carried out in an open system in so- most of the overgrowths was determined following the acid digestion lutions of close to constant composition over a wide range of precip of a known amount of reacted solid by AAS. The mole fraction of nation rates iO-*.5to 104,5 mole m-* hr- i. Constancy of composition MnCOs, MgCOp and SrCOs and the Na content of the overgrowth was maintamed during the length of the precipitation by the use of were calculated from the results of the AAS analyses, the amount of a chemo-stat technique through the simultaneous injection of two carbonate precipitated and the amount of material dissolved for titrants in equal amounts by a dual syringe pump. The mixture of analysis (MUCCI and MORSE, 1983). No correction was introduced the two &rants reproduced the exact composition of the precipitating to compensate for the presence of residual solution salts as the over- solution plus an excess in calcium, manganese(II), strontium and growths were rinsed with distilled water equilibrated with calcite after carbonate alkalinity to compensate for the manganoan magnesian being filtered out of the parent solution (Mucc~, 1986). calcite p~ipi~tion. The temperature of the p~ipi~ting solution Some of the reacted solids were also examined by X-my apron was held constant at 25 (kO.05 )“C by circulating water through a spectrometry to determine their mineralogy and identify other car- jacketed 400 ml glass reaction vessel from a constant temperature bonate mineral phases which might have precipitated along with the bath. The PcoZ of the solution was kept nearly constant at -3000 manganoan magnesian calcite. Powder packs were prepared and ir- ppm or lO-2.5 atm. by bubbling a C02/Nz gas mixture of known radiated using a Siemens model D-500 X-ray dilBactometer. The composition. A detailed description of the chemo-stat, its working Cu-I& wavelength radiation was used as a source and the diffraction concept and titrant compositions have been presented previously spectra were recorded using a proportional counter detector. (MUCCI and MORSE, 1983; Muccl, 1986). Since the distribution coefficient of Mn ( II) in calcite is strongly dependent on the precipitation rate, titrant solutions of widely varied Mn(II) content had to RESULTS be prepared for each set of experiments, Selected parameters and results from single precipitation The precipitation was carried out on about 0.6-0.7 g of seed material in 300 to 400 ml of seawater solution until more than 7 X lo-’ reactions are presented in Table 1. During the precipitation moles of carbonate were precipitated. The amount of carbonate pre- experiments, the Mn2+ concentration of the solution often cipitated was calculated from the weight of the standardized titrant varied with time from its initial value. It was very difficult to used during the length of the experiment and the carbonate alkalinity fine-tune the com~sition of the tin-ants in order to hold the of the solution before and after the pre~pi~tion reaction (according to Eqn. 1 in MLJKI, 1986). precipitating solution initial [ Mn’+] invariant. For this rea- son, the average between the initial and final solution Mn2+ Steady state solution composition and saturation state concentrations is reported. This also allows the reader to ap- preciate the variability of this parameter during the experi- The steady state ion concentration product, [ Ca*+] [ CO:-], of the ment, since the initial solution concentrations were respec- precipitating solution was calculated using the steady state pH measured during the experiment, the carbonate alkalinity and the calcium tively 1.0, 5.0, 10 and 25 ppm in each set of experiments. concentmtion of an aliquot ~~~~ at the end of the precipitation. For most measurements, the steady state saturation of the solution Overgrowth composition was reached within fifteen minutes and remained constant for the length of the experiment. The steady state pH was monitored by two When a solid is precipitated from an infinite reservoir or sets of electrodes which were calibrated against both NBS buffer so- a solution whose composition remains constant despite the lutions and Hansson “TRIS” buffer (8.074 at 25”C, salinity 35). Analytical procedures for the determination of Cat+ and titration precipitation reaction, the partitioning of a trace component, alkalinity have been described previously (MU-, 1986). X( e.g. X = Mn2+, Mg*+, Sr’+), between, for example, calcite The first apparent dissociation constant of boric acid, 4, deter- and a solution can be defined by the Henderson-Kracek mined by LYMAN (1957) or HANSSON(1973) and refittedby MIL- (HENDERSON and KRACEK, 1927) or homogeneous distri- LERO( 1979) was used to calculate the boric acid contribution to the bution coefficient: titration alkalinity. Finally, the steady state carbonate ion concentration was derived from the carbonate alkalinity of the final solution, M&+/M&+ the steady state pH and the second apparent dissociation constant Dcx1+ = (2) of carbonic acid, ir;, determined by MEHRBACHei al., (1973) or ~~=~~M~~~ HANSON (1973) as refitted by MILLERO(1979), accordingto which pH scale was used. The stoichiometric solubility constant of calcite where Mxz+ and Mcaz+ refer to the molar concentrations of Calcite precipitation and Mn uptake 1861

TABLE 1. Sumnary of experimental parameters and composition of overgrowths precipitated from seawater at 25°C.

EXP. [Mn2+1 [++I &+I' pHNaS 0, MC03' d(211) -log R(MC03) XMnC03 'MgCO3 XSrC03 (ppm) (mole hr-1 m-2) (X) (%I (%I (A)

356 5.36 6.9 10.25 7.92 9.59 11.7 3.238 7.6 8.0 0.18 357 4.91 7.2 10.42 7.90 9.49 10.6 3.392 8.1 8.9 0.17 - 359 4.84 7.3 10.56 7.89 8.61 11.4 3.524 8.1 9.1 0.16 - 360 - 10.79 7.85 7.77 8.10 3.669 361 5138 7.1 10.67 7.86 7. 31 9.44 3.796 1o:o 8.3 2.988 362 5.25 7.7 10.64 7.83 6.87 8.35 3.945 10.5 9.4 0.14 - 364 5.27 7.79 5.68 8.86 4.188 365 5.67 718 10.9010.77 7.81 5.39 7.31 4.338 11:5 1o:o 0.12 - 366 4.39 7.0 9.84 8.08 14.2 13.0 2.893 5.4 8.4 0.20 - 367 5.01 7.4 10.29 8.02 13.1 13.2 3.051 7.5 8.5 2.999 368 4.61 7.2 10.28 8.10 16.5 12.7 2.763 5.9 8.3 0.21 - 369 4.76 7.1 10.49 8.10 18.1 10.1 2.638 5.9 7.9 370 10.4 - 10.10 7.99 12.9 12.7 3.051 371 10.6 7.5 10.48 7.98 11.5 11.7 3.107 16.5 7.5 0.16 - 372 9.87 7.5 10.29 7.96 10.5 12.5 3.250 15.5 a.0 0.15 373 10.5 7.4 10.19 8.06 15.2 14.0 2.766 14.5 7.3 0.16 2.990 374 9.25 7.4 10.17 7.94 9.53 13.6 3.384 15.0 8.4 0.15 - 375 10.2 7.4 9.96 8.00 13.1 13.6 2.913 14.0 7.7 0.18 - 376 10.3 - 10.18 8.05 16.9 13.7 2.637 377 9.97 7.5 10.50 7.83 6.56 11.0 3.799 17.5 8.9 0.12 - 378 10.2 7.5 10.25 7.90 7.42 11.9 3.660 17.0 8.6 0.13 - 379 9.95 10.52 7.83 6.16 11.5 3.934 18.0 8.7 2.972 380 10.7 8.2 10.79 7.78 5.21 7.02 4.193 23.0 9.1 0.11 381 27.5 a.2 11.12 7.90 9.44 13.3 3.230 35.0 5.6 0.088 2.953 382 27.9 - 10.97 7.98 12.7 14.6 2.883 383 24.2 - 10.67 8.09 17.6 14.2 2.608 27.5 6.1 2.971 384 28.6 - 10.49 7.75 4.57 11.2 3.919 385 23.5 8.4 10.92 7.77 5.09 11.7 3.927 3g:o 7.0 0:059 2.943- 386 26.5 7.91 9.05 13.2 3.230 35.0 5.9 2.955 387 24.5 812 10.7810.70 7.97 12.3 11.6 2.964 29.5 6.3 0.11 2.966 388 27.9 - 10.86 7.74 3.59 9.54 4.308 389 25.1 8.4 10.98 7.70 3.63 a.55 4.313 40.0 7.2 0.053 2.941 390 24.5 8.3 11.17 7.96 10.7 9.22 3.228 33.0 6.1 0.11 2.958 391 28.5 - 10.43 8.08 14.4 12.7 2.740 393 25.2 8.0 10.78 8.08 14.3 12.0 2.733 28.0 5.6 0.11 2.964 395 1.08 8.0 10.63 8.01 12.6 9.12 3.250 1.67 9.3 0.21 - 396 1.22 - 10.69 8.07 15.4 9.07 2.969 1.67 8.8 397 1.10 811 10.189.95 8.07 10.4 12.5 3.362 1.59 9.0 3.007 398 1.11 7.97 9.44 9.89 3.662 1.92 9.7 0.22 3.004 400 1.18 - 10.79 7.89 8.04 8.12 3.950 402 1.22 8.4 10.61 7.85 6.55 7.56 4.301 2147 10:5 0:24 3.001- 403 0.96 e:3 10.399.09 a.15 16.6 12.5 2.612 404 1.05 7.91 a.27 8.81 3.957 1195 1o:o 0.26 3.006 405 0.93 8.3 10.41 7.94 9.28 11.1 4.728 1.68 10.5 0.26 3.001 406 0.95 10.35 8.05 12.9 10.4 3.098 1.31 8.8 408 1.01 : 9.99 8.26 16.6 11.1 2.642

' mmoles kg-' ; ’ moles x 104.

X2+ and Ca2+ in the overgrowth (c) and the solution (L). composition of the precipitated phase. On the other hand, By analogy to the original definition, X2+ represents a mi- the activity coefficients of the aqueous species are not expected crocomponent and Ca ‘+, the carrier phase. This definition to vary significantly. In a recent review article on the behavior of the distribution coefficient was introduced as a more con- of ionic solid solutions in aqueous solutions, DRIESENS venient expression of the Berthelot-Nemst distribution law (1986) formulated the distribution laws for regular and com- (MCINTYRE, 1963). plex solutions, at equilibrium, as a function of their stoichi- D$+ should be independent of the solid composition as ometry. However, in practice it remains difficult to evaluate long as the concentration of the trace component is small. solid activity coefficients and interaction parameters. Fur- The overgrowths produced in this study contain as much as thermore, as solid solutions precipitated in this study are not 40 mole% MnC03. Under these conditions, Mn*+ no longer formed under equilibrium (stable or met&able; they are qualifies as a microcomponent of the solid. Consequently, precipitation rate dependent) conditions, these formulations superimposed on the precipitation rate effect, the distribution cannot apply. coefficient of Mn2+ as defined in Eqn. (2) is subjected to the The influence of substantial compositional changes of the influence of solution or solid composition. This translates solid can be suppressed or eliminated if one considers the into a series of four slightly converging lines (one for each bulk solid as the carrier phase, and a distribution coefficient set of solution [Mn”]) describing the variation of D&,2+ is defined as: with the precipitation rate. Variations of D&z+ with solution composition are most likely due to sharp variations in the t&z+ = XMoCO~ activity coefficient of the “calcite-carrier” phase resulting from [Mn2’]/[Ca2’] the incorporation of large amounts of MnC03. In other words, the carrier component no longer approximates the where xMvlnCOais the mole fraction of MnC03 in the solid. behavior of an ideal solid solution. A similar conclusion was Since the solution [Ca*+] was almost invariant in all the reached by HOLLAND(1966) after he pointed out that BODINE experiments performed in this study, this expression would et al. (1965) observed that the distribution coefficient of Mn2+ be equivalent to the Berthelot-Nemst distribution coefficient in calcite precipitated at 175°C was sensitive to the bulk if Eqn. (3) was described in terms of activities and applied 1862 A. Mucci to a system containing only trace amounts of Mn’+. The variation of k&z+ as a function of the precipitation rate is presented in Fig. 1. The least squares ftt line drawn through all the data yields the following rate dependence expression:

kcMn2+ = -3.407 x log R(hfCt&) - 2.31 (4) where R( IMCO3) is the overall rate of carbonate precipitation in mole hr-’ md2. The correlation coefficient of the least squares fit is 0.95, As was observed by LBRENS (1981) and PINGITOREei al. (1988) the incorporation of Mn2+ in calcite increases with a decrease in the precipitation rate. Absolute values of the distribution coefficient obtained in this study and the two previous studies are in fairly good agreement Log R(MC0,) (8.7, 12.9 and 9.5, respectively) for a precipitation rate of I mg/min-m2 or 6 mmole/hr-m2. However, the agreement FIG. 2. Variation ofthe Hende~n-bank di~~bution coefficient appears to be fortuitous as results of this study are otherwise of strontium in calcite, L&2+,as a function of the logarithm of the different at other rates. This is certainly related to the differ- carbonate precipitation rate for artificial seawatersolutions of various Mn*+ concentrations; (- -), 0 ppm (from MUCCI,1986); (A), -1 ence in the composition of the precipitating solutions and ppm; (O), -5 ppm; (O), -10 ppm; (A), -25 ppm. particularly the presence of magnesium in seawater (FRANK- LIN and MORSE, 1983). The amount of Mg incorporated in the overgrowths de- words, the Henderson-bank dist~bution coefficient of Sr”* creases with increasing MnC03 content. However, the Hen- increases with the precipitation rate but decreases with the derson-Kracek distribution coefficient of Mg2+, L$,Q+ (Eqn. [ Mn2+] content of the solution or solid. 2), appears to be independent of the solid composition, In The influence of precipitation rate on the distribution coef- other words, the amount of Mg2+ incorporated in the over- ficient of strontium in calcite has been explained by a multiple growth is proportional to the Ca2+ content of the solid, - 1 site mechanism. FINGITORE and EASTMAN( 1986) suggested mole% MgC03 for every 10 mole% CaC03. The value of the that strontium incorporation in calcite increases as the num- dist~bution coefficient, 0.02 1 f 0.003, is larger than the dis- ber of available non-lattice sites or crystal defects increases tribution coefficient determined recently in Mn-free calcites with the precipitation rate and distortion of the calcite lattice (0.0 17 f 0.02, MUCCI and MORSE, 1983; 0.0 17 ? 0.02, BUR- following incorporation of foreign ions (e.g. Mg2+, MUCCI TON and WALTER, 1987; 0.019 r 0.01, OGMORIet al., 1987). and MORSE, 1983). On the other hand, because of its large This may indicate that Mg*+ ions are also incorporated in radius, SrZ+ (1.13 %I)should be less susceptible to lattice sub- what would otherwise be Mn*+ sites. Nevertheless, the dis- stitution in Mn2+ sites (-0.80 A) than in the larger Ca2+ crepancy remains within the unce~inty of the data. sites (-0.99 A) ( KRETZ, 1982f. Furthermore, Mn*’ should The variation of the Henderson-Kracek distribution coef- be abie to compete efficiently with Sr2+ for occupancy of ficient of Sr2+, as expressed by Eqn. (2), is presented as a deformation and non-lattice sites in calcite. The last two ar- function of the precipitation rate in Fig. 2. Included in the guments could explain why the solubility of Sr2+ in the over- figure is the log-linear relationship derived by MUCCI (1986) growth decreases with increasing incorporation of MnC03. for the distribution coefficient of Sr*+ in a Mn-free magnesian Results of this study conflict with the conclusion reached by calcite precipitated from artificial seawater. Four parallel lines, ICHIKUNI (1973), following his analysis of two manganese- one for each solution [ Mn2+] are obtained. These, in turn, bearing travertines. Differences in the precipitation rate of are parallel to the Mn-free magnesian calcite line. In other the travertines could easily account for the observed opposite trend. Likewise, the amount of sodium incorporated in the overgrowths increases with the precipitation rate (see Fig. 3). It also increases slightly with increasing Mn con~ntration in the solution and solid. These results are consistent with those of BUSENBERGand PLUMMER(1985) who proposed that sodium incorporation in calcite is dependent on crystal defect density, which increases with precipitation rate and deformation induced by the substitution of Ca2+ by other cations. The absolute sodium content of the overgrowths are well within the range of values reported by other authors f WHITE, 1978, ISHIKAWA and ICHKUNI, 1984; BUSENBERG and I I I PLUMMER,1985; OKUMURAand KITANO, 1986) for the same -LOO -s.,o 4.00 -..,0 Log R(MCO,) sodium activity in solution. Values of d( 112),,,, or d( 104& for most of the solids i%G. 1. Variation of the distribution coefficient k&z+, defined in were obtained by measuring spacings of the overgrowths Eqn. (3) as a function of the logarithm of the carbonate p~cipi~tion rate for artificial seawater solutions of various Mn*+ concentrations: against that of the pure calcite seed. All the d-values presented (A), -1 ppm; (O), -5 ppm; (O), -10 ppm; (A), -25 ppm. in Table 1 are in excellent agreement (+.005 A) with the Calcite precipitation and Mn uptake 1863

I i I DISCUSSION It is now well established that the distribution coefficient of Mn2+ in calcite at low temperatures is kinetically con- I \--. I trolled. In fact, almost all distribution coefficients obtained -2 at 25°C are influenced by crystal growth rates and several +m *0 models have been proposed to explain experimental obser- a A vations (e.g. OHARAand REID, 1973). LAHANNand SIEBERT w 0 . A (1982) attempted to explain the formation of magnesian cal- L P -I,000 %:0

loooI_-3.00 -3.60 -4.00 Log R(MC0,)

FIG. 3. Sodium content of manganoan calcite overgrowths precipitated from artificial seawater solutions of various Mn’+ concentrations. The symbols are as in Fig. 1. d(211)-isopleth of the CaCOr-MgCOr-MnC03 system presented by GOLDSMITHand GRAF ( 1960). In general, d-values obtained in this study are slightly larger than those extracted from the isopleth. The incorporation of small amounts of Sr’+ in the overgrowths is probably responsible for this ob- I. / 3 * 1.1. I .I 0.60 0.60 0.60 1.00 1.20 servation. Nevertheless, the excellent agreement provides ev- Log (n,- I) idence that the solids precipitated in this study can be considered as one-phase solid solutions of at least three end- member minerals (i.e. calcite-magnesite-rhodochrosite).

Precipitation kinetics The precipitation rate data were fitted to a rate law of the following form:

R = k(C&- 1)” (5) or to its logarithmic form:

logR=logk+nlog(Q,-1) (6) 6.. I.. I where R is the precipitation rate normalized to the reaction I .50 2 .oo P.50 3.00 1.54 surface area (mole hr-’ m-‘), k is the rate constant, n is the Logu-Llt- I) empirical reaction order and (R, - 1) is the degree of supersaturation with respect to calcite. The usefulness of this empirical rate law has been illustrated by numerous investigators (MORSE, 1983; INSKEEPand BLOOM, 1985; MUCCI, 1986; -2.00- BURTON and WALTER, 1987; and many others). Results of n the precipitation rate measurements are represented in Fig. s -2.00 . 4a by a plot of log ( A4C03) versus log (a, - 1). Results plot s as four distinct and intercepting lines, one for each of the :, four sets of experiments corresponding to the different so- z lution [ Mn *+I / [ Ca’+] investigated. In almost all cases, pre- : -6.00 . cipitation rates for a given !A, are faster than for the precip- -I itation of calcite from seawater in the absence of Mn2+ (MUCCI, 1986; BURTON and WALTER, 1987). These obser- vations suggest that more than one stable phase and/or mixed I. 1. 1. 1. I. reaction kinetics might be involved. However, as indicated 030 0.60 0.60 1.00 1.20 previously, X-ray diffraction analysis ofthe precipitates clearly Log WC- I) demonstrates the formation of a solid solution rather than a FIG. 4. Overgrowth precipitation kinetics in artificial seawater so- mechanical mixture. Yet, it can lx shown that the kinetic lutions of various Mn*+ concentrations. (a) Total carbonate as a function of calcite supersaturation, (b ) manganese carbonate com- data can be used to model the composition of the overgrowths ponent as a function of kutnahorite supersaturation, and (c) calcite in terms of the precipitation rate of a magnesian calcite and component as a function of calcite supersaturation. Symbols are as a phase whose stoichiometry would be similar to kutnahorite. in Fig. 2. 1864 A. Mucci cites in seawater based on a mass transfer kinetic model. R(MnC03) = R(MC03) X XM”~~. (9) They proposed that the rate of mass transfer between the The stoichiometric solubility constant of rhodochrosite in solution, the calcite surface and the calcite crystal is propor- seawater determined by JOHNSON( 1982; 3.24 X 10m9mole 2 tional to the activity of the end-member components (i.e. kgV2 SW) at 25°C was used to calculate the saturation state CaC03 and MgCOj) in the “reactant” phase. According to of the precipitating solution with respect to rhodochrosite. the model, the distribution coefficient for a given cation is The log R( MnC03) versus log (Q, - 1) plot yielded four expected to decrease or increase with increasing supersatu- widely separated parallel lines, one for each set of experiments. ration, depending on the solubility of the pure end-member This indicates that either the precipitation of rhodochrosite phases and the activity of the solid phases. in seawater does not obey this particular rate law or that It appears, in the light of more recent data (MUCCI and rhodochrosite is not the manganese carbonate end-member MORSE, 1983; MUCCI, 1986, 1987; BURTON and WALTER, component ofthe solid solution. In the latter case, kutnahorite 1987; OOMORI et al., 1987), that the model does not apply could be considered as the end-member component. This to “inorganic” magnesian calcites since their composition is choice. has no strict theoretical basis but is made as a feedback independent of the precipitation rate. The magnesium con- to the model. The~~yna~~iy, it does not matter what centration in abiotic magnesian calcites is controlled strictly phase is chosen as the end-member of a solid solution series, by the [ Mg2’]/[Ca2’] of the parent solution (Mucc1 and as long as there is no unmixing in the system. MORSE, 1983; see MORSE, 1985, for discussion) and possibly The thermodynamic solubility of kutnahorite was calcu- by the nature of the crystallographic faces on which growth lated from its Gibbs free energy of formation (-466.2 kcal occurs ( REEDER and GRAMS, 1987 ) . mol-‘; GARRELSet al., 1960) at 25°C and one atmosphere The model proposed by LAHANN and SIEBERT(1982) may total pressure and the free energies of formation of Ca2+, be suitable to describe the incorporation of MnzC in calcite. Mn” and CO:- ( ROBIE et al., 1979), resulting in a value of As is observed, it predicts that the in~~mtion of Mn2+ in pru” = 19.84. The saturation state of the precipitating solu- calcite should decrease with increasing precipitation rate. The tions with respect to kutnahorite was calculated using total model is also consistent with a mechanism in which the ion activity coefficients derived from the ion pairing equations amount of Mn2+ incorporated in calcite is limited by the rate of MILLERO and SCHREIBER(1982; yT( Ca2+)yT( Mn2+)- of adsorption of Mn2+ on the growing surface. A similar rf(CO:-) = 2.82 X lo-“). The saturation state of the so- mechanism was proposed by MUC~I (1986) to explain the lutions is obviously dependent upon the precision of the ac- increased calcite p~ipitation-i~bition effectiveness of or- tivity coefficient estimates. For example, the total ion activity thophosphate ions with decreasing precipitation rate. coefficient product derived from the estimates of PLUMMER Since the rate laws for individual mass transfer reactions and SUNDQUIST(1982; for Y~( Ca2’) and Y=( CO:-)) and of between the solution, the surface and the solid are not readily KLINKHAMMER(1980; for yT(Mn2+)) differ by more than available it is not possible to apply the model proposed by 90%. However, the conclusions drawn from these calculations LAHANNand SIEBERT( 1982) in its mechanistic form. How- are not afleeted by this uncertainty. Saturation state calcu- ever, it may be possible to derive a rate law which would lations indicate that all solutions are supersaturated with re- describe the incorporation of MnC03 in calcite in terms of spect to both rhodochrosite and kutnahorite. All the solutions the MnC03 activity product in the solution in order to satisfy are 10 to 40 times more supersaturated with respect to kut- a more general expression of the model (Eqn. 40 in LAHANN nahorite than rhodochrosite. and SIEBERT, 1982): The log R( MnC03) versus log (Q, - 1) plot yielded one line reflecting a simple linear correlation (Fig. 4b). According (7) to Eqn. (6) the slope of the least squares fit line through these nearly 40 data points would correspond to the empirical re- or action order of the kutnahorite precipitation reaction. Its dMn/dt value in this case is 1.24. The intercept, which corresponds xMnco3= dMn/dt + dCafdt to the log of the rate constant is equal to -7.59. The correlation coefficient of the fit is .995. These values are tentative R(MnCOs) RWnCW (8j since the solubility of kutnaho~te in seawater has not yet = R(MnC03) + R(CaCOJ) = R(MC03) been measured. More importantly, as will be discussed fur- where dCa/dt and dklnldt or R(CaC03) and R(MnC03) ther, the solubility of a “disordered” kutnahorite or “pseu- are the precipitation rates of the end-member minerals, calcite dokutnahorite” would have been more appropriate to use and rhodochrosite, respectively, in seawater. since ordering is probably not obtained under the experi- The precipitation rate of rhodochrosite in seawater has not mental conditions. In addition, the presence of ordered kut- been measured. However, if the model described above does nahorite in the solid would imply the existence of two distinct apply, it should be possible to extract the rate data from results phases, which is not supported by the X-ray diffraction data. of this study. If we assume that the precipitation of rhodo- The solubility of “pseudokutnahorite” in aqueous solutions chrosite in seawater follows a growth law identical to Eqns. is unknown but the rate law described above would still apply (S)or(6),thenaplotoflogR(MnC03)versuslog(Q,- I) although the absolute value of the rate constant and reaction should show a linear relationship. order would certainly be different. The precipitation rate of MnCO-, was derived from the The p~ipi~tion rate of the CaCO3 component of the total carbonate precipitation rate and the mole fraction of precipitate was plotted as a function of the supersaturation MnC03 in the precipitate, as given in Table 1, according to: with respect to calcite, where: Calcite precipitation and Mn uptake 1865

R(CaC03) = R(MC03) X Xcaco,. (lo) the solubility of a “pseudokutnahorite” precipitated from seawater can be confirmed The log R( CaC03) versus log (Q, - 1) plot gave results similar The decomposition of the overall reaction rates into two to those in Fig. 4a, but with lines closer together. If it is as- individual rate laws does not mean that they actually represent sumed, for the purpose of the model, that “pseudokutna- “elementary” reactions. It only serves to apply the model. horite” is the end-member component, as Fig. 4b would sug- The interpretation of the results of this study may not be gest, then the rate of the calcite-CaCOs component of the unique nor is it truly original. LERMAN( 1965 ) attempted to precipitate and not that of all the CaC03 should follow the use dolomite as an end-member solid component instead of calcite precipitation rate law. Based on a mole balance of the magnesite when modeling the thermodynamic properties of carbonate ion, the rate of calcite precipitation would become: magnesian calcite solid solutions.

R(c-CaC03) = R(MC03) X (1 - 2 X XMvlnco,). (11) Kutnahorite or ‘bseudokutnahorite” The possible occurrence of the mineral kutnahorite in When plotted as a function of log (4 - l), Fig. 4c is obtained. suboxic and anoxic marine sediments has been reported by It reproduces almost exactly the calcite ( -8 mole% magne- several authors ( CALVERTand PRICE, 1970; PEDERSEN and sian calcite) precipitation rate law observed previously by PRICE, 1982: and references within). However, positive MUCCI(1986) in seawater in the absence of Mn*+ (n = 2.9 identification of kutnahorite has been hampered by the lack vs. 2.8, log k = -6.28 vs. -6.29 and corr. coeff. = 0.974). of ordering reflections in the X-ray powder diffraction spectra. Thus, it appears that the total carbonate precipitation rate Kutnahorite is an isotype of dolomite and in fact has often data represented in Fig. 4a can be decomposed into two in- been referred to as manganese dolomite (FRONDEL and dependent rate laws characterizing the precipitation of calcite BAUER, 1955). The ordering reflections that distinguish the and “pscudokutnahorite” from seawater. It could imply that mineral kutnahorite from a disordered solid solution of the solids which form under such conditions are likely to be me- same composition are weak and cannot always be detected chanical mixtures rather than solid solutions. However, the with confidence by X-ray powder diffraction (GOLDSMITH X-ray diffraction patterns clearly indicate that the solids are and GRAF, 1957; GOLDSMITH, 1983). They are better mea- solid solutions and this lends support to the applicability of sured by single-crystal techniques (PEACOR et al., 1987). the kinetic distribution coefficient model described by Eqns. CAP~BIANCO and NAVROTSKY(1987) indicated that dolo- (7) and (8). mite-type order within Mn-bearing calcite is unlikely at low If it can be interpreted, for modeling purposes, that “pseu- temperatures typical of the ocean floor because of kinetic dokutnahorite” rather than rhodochrosite is the manganese hindrance. It was suggested that naturally-occurring phases carbonate end-member component of the solid solutions, near CaMn(C03)* in composition should be called “pseu- then: dokutnahorite” ( CAPOBIANCOand NAVROTSKY, 1987 ) or “disordered kutnahorite” ( PEACOR et al., 1987), since it is R(MCOJ) = 2R(MnC03) + R(c-CaC03) (12) more likely that they formed without any dolomite-type or- or dering. Natural samples of kutnahorite show some solid solution R(MCOs) = 10-7~2g(Qk- 1)‘.24 + 10-6.28(R, - l)*.‘. (13) with various end-members, including CaC03, MgCOs and MnC03 ( REEDER, 1983 ) , but solubility must be very limited While the manganese content of the solid solutions can bc since the presence of an ordered kutnahorite would imply calculated by substituting Eqn. (13) into Eqn. (4), it can also that two separate phases are present in the solid. Indeed, be derived from the application of the mass transfer mode1 FRONDELand BAUER (1955) and GOLDSMITH(1983) have or from the relative precipitation rate of the two major com- identified and analyzed what appeared to be mechanical ponents, according to Eqn. (8): mixtures of kutnahorite and calcite containing several pcr- 10-‘.59(fjk _ 1)1.*4 cent Mg. Glnco, = 10-7.29(fik_ 1)'.24+ lo-6.28(~, _ q2.9 (14) PEACOR et al. (1987) demonstrated that compositions of most natural manganoan carbonates reported in the literature from which k&z+, as defined in Eqn. (3), can be calculated are well represented by the ternary system CaC03-MnC03- if the solution [ Mn*+] and [ Ca*+] are known. According to MgC03 since they commonly contain between 5 and 10

Eqn. (14) xMulnC0~would tend towards a value of 0.5 (i.e. mole% MgC03 in substitution. Manganoan calcite over- CaMn( C09)*) when the solution is in equilibrium with cal- growths precipitated in this study all contain between 5 and cite. Since the composition of the solid solutions is kinetically 10 mole% MgC03. Their compositions are plotted on a three controlled it is also conceivable that the presence of specific phase triangle (Fig. 5) and are found in the region where growth inhibitors (e.g. phosphate, organic matter, which are GOLDSMITH and GRAF (1960) observed a single phase to often abundant in sediment pore waters) to either or both occur at high temperatures. At 25°C these solid solutions are end-member components could significantly alter their com- most likely metastable but, as was observed for CaC03- position. Again it should be emphasized that constants for MnCOJ solid solutions (GOLDSMITH and GRAF, 1957; DE the empirical precipitation rate equation describing the CAPITANIand PETERS, 198 1) , will probably not demix under CaMn(COs)2 component growth kinetics depend on the laboratory conditions. value of the stoichiometric solubility constant of “pseudo- Constancy of Of,,*+ for the overgrowths precipitated in kutnahorite” in seawater that is used to calculate Qk. Con- this study indicates that Mg*+ substitution for Ca*+ is not stants which appear in Eqn. (13) and (14) are tentative until affected by the formation of this hypothetical “pseudokut- 1866 A. Mucci

CaCO, laboratory. After more than one year of equilibration, active uptake of Mn2+ and variations of the CaC03 ion concentration product are still observed but tend towards values lower than manganese-free calcite solubility products. Kutnahorite solubility measurements in distilled water and seawater have recently been initiated.

CONCLUSIONS As was observed by LORENS(198 1) and PINGITORE et al. (1988)) results of this study indicate that the amount of manganese incorporated during calcite precipitation decreases with increasing precipitation rate. CaMgKO,), Co MII(CO,)~ A kinetic distribution coefficient model based on the individual precipitation kinetics of the manganoan calcite solid PIG. 5. Composition of manganoan magnesian calcite overgrowths solution end-member components was applied. Results of precipitated in this study onto a partial CaCOr-MnCOx-MgCOster- nary phase diagram in comparison with the experimental solvi of the model could be interpreted to identify “pseudokutna- GOLDSMITHand GRAF (1960) at 500” and 700°C. (a) one phase horite” as the possible manganese-bearing carbonate end- region; (b) two phases region. The formulae CaMg( CO& and member. The precipitation kinetics of the proposed solid so- CaMn(COa)z indicate chemical composition and do not imply or- lution end-member components, calcite and “pseudokut- dering in the structure. nahorite”, modeled individually according to empirical rate equations, could also be used to determine both the overall nahorite”-calcite solid solution. In fact, an equivalent number carbonate precipitation rate and the composition of the pre- of Ca’+ sites remains available whether the solid solution is cipitate. It should be noted that the model does not imply with “pseudokutnahorite” or rhodochrosite. From the ex- that individual reaction rate laws are representative of ele- trapolation of data presented in this study it is expected that mentary reactions. a 50 mole% MnC03 manganoan calcite or “pseudokutna- The kinetic model predicts that close to calcite saturation horite” precipitated from seawater would contain - 5 mole% a manganoan calcite containing approximately 50 mole% MgCOJ, such as the mixed carbonate recovered by PEDERSEN MnC03 should precipitate from seawater solutions (e.g. ma- and PRICE (1982) in Panama Basin sediments. rine sediment pore waters) containing >1.5 ppm dissolved Mn2+, not unlike the mixed carbonate or “pseudokutnahor- IMPLICATIONS ite” recovered by PEDERSENand PRICE (1982) in Panama Basin sediments. The occurrence of mixed Mn-Ca-Mg carbonate solid so- The amount of magnesium precipitated with the over- lutions, some of which may approximate the composition of growths decreases with increasing MnC03 content. How- kutnahorite, has been used to explain the apparent non-equi- ever, the ratio of Mg:Ca in the solid remains constant at librium behavior of carbonate-rich deep-sea sediment pore about 1: 10. waters with respect to calcite and rhodochrosite. In view of The distribution coefficient of Sr2+, D&Z+, increases with past reports and the results ofthis study, the existence of such the precipitation rate, but decreases slightly with the [ Mn2+] solid solutions or a “pseudokutnahorite” mineral phase in content of the solution or solid. Competition between Mn2+ marine sediments appears highly probable. Sediment pore and Sr2+ for lattice and non-lattice sites in calcite can explain waters in the suboxic zone contain highly variable dissolved this observation. manganese concentrations (OS-400 pmole kg-’ or up to Finally, the amount of Na incorporated in the overgrowths -22 ppm; MIDDELBURG et al., 1987). If the pore waters increases with the precipitation rate and the Mn concentra- are saturated with calcite, Mn2+ concentrations greater than tion. This is consistent with previous findings, which indicate about 27 Hmole kg-’ (- 1.5 ppm), at 25°C and one atmo- that sodium incorporation in calcite is dependent on crystal sphere, would suffice to result in a supersaturation with respect defect density. The number of defects increases with precip- to kutnahorite. In carbonate-rich sediments thin coatings of itation rate and deformation resulting from the substitution manganoan calcite would form but these would be unde- of Ca2+ by other cations (e.g. Mn’+, Mg2+, Sr’+). tectable by X-ray diffraction. On the other hand, in a carbonate-poor sediment the presence of few carbonate nuclei Acknowledgements-The author wishes to thank Dr. John W. Morse could allow the formation of X-ray distinguishable manga- for reading and commenting on an earlier version of this manuscript noan calcites or “pseudokutnahorite”. The precipitation of and Dr. Robert F. Martin for pointing out pertinent literature. The manuscript benefited greatly from the comments of Drs. Rob N. J. rhodochrosite would probably require very high [Mn’+] / Comans, Nicholas E. Pi&ore, two anonymous reviewers and the [ Ca2+] in seawater. patience of the handling editor. Controversy still exists concerning the solubility of these Financial support for this study was provided by the Natural Sci- solid phases in deep-sea sediments (see BOYLE, 1983 ) and ences and Engineering Research Council of Canada through grants as to whether or not they can account for the observed sat- #UO432 and E 1446. Completion ofthis project was also made possible through the award of an equipment grant by the Faculty of Graduate uration state of pore waters in the reduced manganese zone. Studies and Research, McGill University. Solubility measurements of calcite in seawater in the presence of varying concentrations of Mn2+ are in progress in this Editorial handling: S. E. Calvert Calcite precipitation and Mn uptake 1867

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