Behaviour of calcium carbonate in sea water

Behaviour of calcium carbonate in sea water* P. E. CLOUD, U.S. Geological

Survey,

JR.

Washington

25, D.C.1

Abstract-Anomalies in the behaviour of calcium carbonate in natural solutions diminish when Best values found by traditional oceanographic methods for the apparent considered in context. solubility product constant K’caco, in sea water at atmospheric pressure are consistent mineralogically-at 36 parts per thousand sa1init.y and T-ZS”C, K ‘sragonite is estimated as 1.12 x 1OV values are 0.98 x 1OW for aragonite and J~‘oa~eiteas 0.61 x 10e6. At 30°C the corresponding and 0.53 x lO-‘j for calcite. Because the K’ computations do not compensate for ionic activity, however, they cannot give thermodynamically satisfactory results. It is of interest, therefore, that approximate methods and information now available permit the est’imation from the same basic data of an activity product constant Kcaco, close to that found in solutions to which Debye-Hiickel theory applies. Such methods indicate approximate Ksrsgonite 7.8 x 10e9 fol lower. surface sea water at 29°C; Kcailcite would be proportionately Field data and experimental results indicate that the mineralogy of precipitated C’aCO, depends primarily on degree of supersaturation, t,hus also on kinetic or biologic factors that faciliThe shallow, generally hypersaline bank waters t,ate or inhibit a high degree of supersaturation. west of Andros Island yield aragonitic sediments with 018/016 ratios t’hat imply precipitation mainly during t’he warmer months, when the combination of a high rat,e of evaporation, increasing salinit,y (and ionic st’rength), maximal temperatures and photosynthetic removal of CO, result in high apparent supersaturation. The usual precipitate from solutions of low ionic st,rength is calcit’e, except where the aragonibe level of supersaturation is reached as a rest& of diffusion phenomena (e.g. dripstones), gradual and marked evaporation, or biologic intcrrcntion. Published data also suggest, the possibility of dist,inct chemical milieus for crystallographic variations in skeletal calcium carbonate. It appears that in nature aragonite precipitates from solutions t’hat are supersaturated wit’h respect to both calcite and aragonite and calcit,e between saturation levels for the two species. Such a relation is consistent, wit,11 Os~wa~n's rlllc of successi\re rcact,ions. Aragonit,r of marine origin persist’s in cont’act with supersaturated interstitial solutions at, ordinary, temperature and pressure. Conversion to calcite follows transfer to solrltions undersaturated with respect to aragonite or upon exposure to the moist atmosphere.

STUDIES of the shoal marine environment west of Andros Island, Bahamas, lead to t’he conclusion that anomalies in the behaviour of calcium carbonate diminish when the relevant details from different disciplines are viewed concurrent,ly. Because of current interest in the subject, I have been urged to summarize this view in a more generally available form than the detailed account from Tvhich it originally emerged (CLOUD et cd., in press). Such condensation, naturally, involves omissions and incompletely documented generalizations for which I can but ask the indulgence of the critical reader pending his examination of the complete account. With apologies to the pure solution chemists, I will also stretch part of the theory of electrolytic solut’ions far beyond its accepted range, and will invoke the results of

* Publication authorized by the Director, U.S. Geological Survey. This pa,per has l>rofitetl from crit,ical reading by R. M. GARRELS, P. BARTON and F. DASIELS, and from the met,icrdous typing and editorial services of DORIS Low. t Present address: Dept. of Geology, Universit,y Minnesota, Minneapolis 14. 867

868

P.E.CLOUD,JR.

Better methods would be welcome, but the problem some inelegant computations. does not cease to exist in their absence, and the manipulations employed give values that are more directly comparable with thermodynamic results than the procedures usually applied to sea water. If my crude efforts achieve no more than to provoke better qualified persons to attack the problem t,hey will not have been in vain. -4s to the problem: although it has been repeatedly demonstrated that calcite is the equilibrium crystallographic form of calcium carbonate at atmospheric temperature and pressure, the precipitation and persistence in sea water of its relatively unstable polymorph aragonite respond as if to an equilibrated system. At the same time empirical, quasi-stoichiometric methods used to estimate the apparent solubility product constant of calcium carbonate (K&ol) in the sea give a wide span of computed values and indicate high supersaturations, even in some waters where other evidence suggests recurrent solution. Curiously also, non-skeletal calcium carbonate characteristically takes the form of aragonite in sea water and of calcite in open fresh waters (cave and other dripstones are another problem). To explain these and other anomalies, it has been variously suggested that the precipitation and persistence of aragonite depends on’some particular relation to temperature, salinity, pH or some special biologic or minorelement control. High apparent supersaturation is increasingly attributed to chemical or organic complexing. In this summary, evidence relating to the origin of aragonite sediments in the area west of Andros Island is first briefly outlined as background for the part that follows. Then the chemical state of this water and its mineralogic implications is evaluated, maintaining distinction between the crystallographically and thermodynamically different polymorphs of calcium carbonate. BAHAMAK

ARAGONITE

MUDS

Field relations !l’he area west of Andros Island, Great Bahama Banks (Fig. l), illustrates natural conditions at a marine site of high calcium-carbonate withdrawal not obviously related to biological secretion (such as an organic reef). A persistent salinity gradient rises eastward from 36 or 37 parts per thousand dissolved solids at bank edge to a maximal bankwise average exceeding 39 and to local peaks of 42 to 46 parts per thousand in the lee of Andros Island. The calcium and carbonate ioils, however, do not maintain constant relations to chloride as with simple Instead, alkalinity varies inversely with chloride, indicating a loss concentration. of 0.6 t’o 0.S mequiv./l. of CO, 2-- during movement of the water from the western bank edge to Andros Island. When, moreover, ratios of titration alkalinity and of analytical calcium to chloride are plotted against increasing chloride, they show regular and nearly equal decreases across the bank-proper that imply progressive concurrent removal of Ca2+ and COs2+ from Dhe bank waters (Fig. 2). The anomalous relationships indicated from bank edge outward in Fig. 2 may reflect complexing or rate phenomena, the uncertain ancestry of waters at the surface of the straits and the edge of the banks, or some combination of these things. They do not detract from the ,regularity of change across the bank itself.

Behaviour

of calcium

carbonate

869

in sea water

Computations from analytical data of the rate of formation of CaCO, accord with figures calculated from sediment thickness and radiocarbon age differences, allowing for some transportation seaward off the banks. This concordance, and other data discussed in full elsewhere (CLOUD et al. in press), indicates that the deposits formed where found and eliminates need to call on particulate transport of

25N

240 N

Fig. 1. Area bet)ween Andros Island and Miami, showing location number of stations occupied along each.

of traverses

and

their components from Andros Island or other places. Indeed, major differences are found between bank sediments at different sites of deposition. Oiilite and skeletal limesands form underwater dunes in areas of strong currents, whereas the sediment beneath the sluggishly moving hypersaline water in the lee of Andros Island consists of aragonite needles 2 to 8 p long, fecal pellets and subordinate skeletal debris. @igin LOWEKSTAM and EPSTEIN (1957) have studied the oxygen isotope ratios of oiilitic sands, aragonite muds and aragonite needles precipitated by algae with reference to the question of origin of the bank sediments. According to the 01*/016 ratios found after salinity adjustments by these authors (1957, p. 372), equilibrium temperatures should be about 24 to 25.7°C for the oiiids, 27.6 to 31.7” for the sedimentary aragonite needles and 23 to 40” for the algae. As LOWESSTAM

P. E. CLOUD,

8Sfl

JR.

and EPSTEIN also indicated, such results suggest the formation of known algal aragonite out of equilibrium with 018/016 ratios of surrounding sea water, but formation of the ooids at equilibrium in cooler bank-u~argin waters. Inasmuch as the computed equilibrium temperature range for the sedimentary aragonite needles falls near the middle of the range for those of algal origin, these authors also suggest an algal origin for the bulk needle-sediments.

j.130

1.08~.

I\\ \

At/Cl-

P

SaTiO/Opi STRAITS OF FLORIDA AND OCEANIC SURFACE

,--

,

,

FUNK OUTER EDGE BANK 1 w.

1:: 7 1’

’

20 Fig.

2. Ratios

,__-___-v--wATERS

21

,.

,.

.

22

MID-BANK I 23

‘0 INNER BANK i 24

of combining equivalents of calcium and titration chloride plotted against increasing chloride.

wfZ_ alkalinity

cito

On the basis of field and conventional analytical evidence, however, I prefer an alternative interpretation, also recognized by LOWENSTAMand EPSTEIN, which regards the isotope data as not inconsistent with equilibrium formation of the sedimentary aragonite muds according to possible or even likely average water temperatures and salinities at and near their sites of deposition during the mont>hs of most active CaCO, sedimentation. To be sure, the early morning temperatures in the tidally flushed bights midlength of Andros Island (Fig. 3), and the average temperatures of surface oeean waters in the region (Fig. 4), barely fall within the range of estimated eql~ilibriunl averages during the mont,hs of demonstrable nlaximum precipitation and the given averages also imply formation over a range of temperatures. But the likelihood of general formation of the sedimentary aragonite needles at isotopic equilibrium increases when allowance is made for

Behaviour

of calcium

carbonate

871

in sea water

34 32 ,030 28

w E 1 l;

26

2

22

5 w I-

2c )-

24

IEII

1

II

I

JAN. FEB. MAR. APR.

I

I

MAY JUNE JULY

I

II

I

AUG. SEPT. OCT. NOV.

1

DEC.

Fig. 3. Early morning temperatures at Middle Bight, Andros Island, through the year 1939, compared with suggested equilibrium temperature range for oxygen isotope ratios found in sedimentary aragonite needles. (Water temperatures from SMITH, 1940; oxygen isotope data from LOWENSTA>I and EPSTEIN, 1957.)

a ---_--W a 26-

L

JAN. ’ FEB. ’ MAR.’ APR.”

I

I

MM ’ JUNE ’ JULY ’ AUG. ’ SEPT.’ OCT. ’ &v.

’ DEC. ’

Fig. 4. Average monthly surface temperatures of Atlantic Ocean near Andros Island, and adjusted probable temperatures of bank waters, compared with suggested equilibrium temperature range for oxygen isotope ratios found in sedimentary aragonite needles. (Bank temperatures based on adjustment of FuuLISTER’S 1947 data for ocean water to account for available local data on bank temperatures and seasonal wind changes; oxygen isotope data from LO~EICSTAM and EPSTEIN, 1957).

872

P. E.

CLOUD,

JR.

higher average water temperatures over the bank-proper (Fig. 4), and the probability of water temperatures rather closely following air temperatures (Fig. 3) at the shallow near-shore sites from which the isotopi~ally analysed sedimentary samples were obtained (LOWENSTAM and EPSTEIX, 1957,p. 370,Fig.1; USHO Chart No. 26A). The unadjusted Sol* values of sedimentary needles from various sites, moreover, are internally consistent with regard to expectable local temperature and salinity variations (as are the disequilibrium algal values). For instance, 801s values for the Yellow Cay samples off the tidal pass through North Bight are intermediate between those for needles from the sluggish hypersaline waters off \iVilliams Island and vahres found for the cooler water oiiids at a site of nearnormal oceanic salinity (LOWENSTAM and EPSTEIN, 1957,p. 370).It is to be noted also that the cSOl*values for Halimeda, probably the main source of algal needles, are consistently lower than for other algae or for the sediments themselves (LOWENSTAM and EPSTEIN, 1957, Fig. 5). And it remains to be determined whet,her or not aragonite needles inorganically precipitated from the waters in question actually have equilibrium 60”s values. In view of the preceding I find it less anomalous to suppose that the adjusted siOr* results for the sedimentary aragonite needles fortuitously fall in the middle of the algal needle spread than to conclude that the ooids, on t,he one hand, were inorganically precipitated at sites of relatively low temperature and salinity, while the sedimentary needles, on the other, were formed primarily by organic processes at sites of demonstrably higher salinity and water temperature, and comparable algal floras. As for bacterial effects, the principal identifiable one within the sediments and waters west of Andros Island is the production of CO,, which runs courner to the processes normally thought of as leading to CaCO, precipitation. In addition, only about 20 per cent of the sediment can be identified on other criteria as of skeletal origin (testaceous and algal). This, allowing for trivial recycling and possible bacterial effects, leaves 60-75 per cent of the total sediment unaccounted for. Processes that are demonstrably at work in the area of interest could produce this large unaccounted-for fraction of the sediment through purely chemical reactions. Loss of CO, as a result of evaporation, photosynthesis and rising temperature both across the bank and with advancing season (Fig. s), increases the C0,2- fracLion of the alkalinity complex and induces calcium carbonate precipitation, leading to depression of titration alkalinity and the ratio of alkalinity to chloride (Fig. 2). The systematic chemical changes across the bank are consistent with a primarily open-water chemical mechanisnl, and it is not necessary to invoke other and, according to my data, quantitatively inadequate processes except as adjuncts. I should emphasize here that I do not take the oxygen isotope dat’a lightly, especially in view of their impeccable source. It is disturbing to find that conclusions drawn from such evidence conflict with my own. This means t,hat better controls are needed somewhere. There being no immediate prospect of obtaining the crucial evidence, however, it behooves us to look at all sides of what we have. The results of such an analysis, summarized above, lead me to favour a primarily

873

Behaviour of calcium carbonate in sea water

chemical origin for the aragonite muds; were it not for the contrary interpretations drawn from the isotopic data I should have considered a chemical mechanism proved. This is the point of departure, but by no means the sole basis for t,he discussions that follow. SOLUBILITY RELATIONS Assuming eventual equilibration in water from which CaCO, is precipitating, a value can be computed for the apparent solubility product constant Kbaco,, employing traditional oceanographic methods (e.g. SVERDRUP et al., 1942, p. 193-210) whereby analytical calcium figured as calcium ion is multiplied by Spec;;micdinity cl-%0 0.130 i 0.120 -

0.110 -

0.100 -

0.090 -

0.080 -

0.070 -

I

34

35

36

37

38

39

40

41

42

43

44 %. SALINITY

Fig. 5. Seasonal variation in specific alkalinity with salinity in waters over the Great Bahama Bank (unpublished data of SMITH present’ed by him for use in this report).

computed carbonate. This gives the ionic product cCa2+ x cCO,~- = 1.2 x 1OW at, 25°C and 3676, salinity as the minimal value for bank water, identical with the value estimated by SMITH (1941, p. 240) when his figure is adjusted to Karauonitt: 25°C and for differences in pH scale and second apparent dissociation constant of H,CO, (HINDMAN, 1943, p. 131). A value for Khaco, of 0.61 x 1OW at the same tempera,ture and pressure (after similar adjustments by me) was found independently by WATTENBERG and TIMMERMAN (1936, p. 25) and by HINDMAN (1943, p. 132), using calcite as the solid phase. At 30°C the corresponding values are 0.98 x 10e6 for aragonite and O-53 x 10e6 for calcite. The values given are, thus, internally consistent with respect to aragonite and to calcite. Moreover, from the standard relation AF” = - 1364.3 log K in Cal/mole (e.g. LATIMER, 1952, p. S), taking K’ as K, the standard free-energy difference indicated by these results for the transit,ion aragonite to calcite is -360 Cal/mole, which is within the range of the most widely publicized experimental values that have been determined by different methods over the past quarter century. The

874

P.E.CLO~D,JR.

supposed oonflict between estimates for XfCaCO, in sea water is thus in effect eliminated when mineralogy and free-energy difference are taken into consideration. The main steps in estimating the empirical ~ararn~ter~ for selected stations are given in Table f, and the relation to increasing salinity of t,he ionic product cCa2+ x cCO,~- is graphed in Fig. 6. The computations have been discussed by many previous authors (e”g. SVERDRUP et al., 1942, pp. 195-208), and I have summarized and illustrated them at another place (in CLOUD et al., in press).

Field

25

T”C “$30

8.2 I

-b.9

Fk---T,

22

723

24

25

Fig. 6. Relation to changing chloride of empirical cCazi x cCQ2+ and its ratio to chloride: fw Bahaman water samples with accurate c&&m det~~i~ati~ns jionic products are given as apparent vahs nt %S”C, 36?& salinity, at,mospheric pressure and field pH).

The empirical values obtained, however, are more than two orders of magnitude larger than estimates of the activity product constant Xc.co, for tJhe pertinent ~i~~neralogic~~lspecies in pure water solution. Horeover, since the method does not take account of ealcinm activity, and only indirectly of carbonate activity, K’ can never closely approach K at other than infinitely small ionic strengths. The way out of this dilemma, therefore, is not increasing refinement in the determinat’ion of K’ but a method for the estimation of R. When the stoichiometric product mCa+ x wCO,~- is multiplied by the appropriate activity coefficients (+az+ 1 mCs2+) x (p20,2- *~zCO,~-), the activity product aCa2+ x c&O,~- is obtained. Since, also, the data suggest that the nat,uraf system west, of Andros Island approaches something resembling equilibrium (Figs, 6 and 7), the minimal value there found for the activity product should approach

_

Behaviour

of calcium

carbonate

in sea mater

x75

K arajionite. Employing

works published mainly since 1950, but extending ideas that date from the 1920’s, estimation of the activity product was, therefore, attempted. The Debye-Hiickel theory as it applies at different ionic strengths, other methods, the equations and the standard values needed to approximate the activity product aCa2+ x aC0,2- are given in papers by ROBINSON and STOKES (1949, 1955), KLOTZ (1950, p. 300-358), GARRELS and DREYER (1952, p. 332-336), LATIMER (l.952, p. 349-351), HARNED and OWEN (1958, p. 449-490) and GARRELR

Field 1

4 8.0

30%

.X ,. ‘x. . . . . . . . . . . . . . .x”

jA\

29% 1 28°C

i7.0

-6.0 PH

7”” 8.1

120PH

80 - 4.0

lO.O-

80

3.0

c

l------I 21

20

-5.0

i

22

23

rb

i5

26

27

1

9%

cl-

Fig. 7. Relation to changing chloride of activity product aCa2+ x CZCO,~+and it,s radio to chloride for same water samples as in Fig. 6 (temperature 29.25”C, but computations

assume T 25°C).

(1960, p. 27-30, 39-40, 43-60). Utilizing these sources, the approach of GARRELS and DREYER and of GARRELS, and computing to standard values at 26’Y’, a minimal activity product of 7.8 x 1O-g was estimated at peak salinity station C7 at the shoreward end of traverse C (Fig. 1). All things considered, this is surprisingly close to the value of 6.9 x 10Wgcited for Kg,ragonite by LATIMER (1962, p. 319). The standard free-energy difference indicated from the aragonite value here found to the figure given in LATIMER for calcite is -300 Cal/mole. This is less than the -360 cal suggested by the empirical data but approaches the experimental value of -273 cal found for the aragonite-+calcite transition by JAMIESOX (1953! p. 1389) and by KELLEY and ANDERSON (1935, p. 15), and the -311 cal found by KOBAYASHI (1952, p. 116). 5

876

P.E.

CLOUD,JR.

The main steps in the activity-product computation, and the relation to salinity of the activity product and mass product for selected stations and groups of are explained sta)tions are given in Table 2 and Figs. 7 and 8. The computations and illustrated in papers cited above, and by me elsewhere (in CLOUD etal., in press), and will not be elaborated here. Comparison between Figs. 6 and 7 (and Tables 1 and 2) shows the similarities and differences between the empirical oceanographic results and the thermodynamic results. From the steps and the values indicated it 0 oca*x oco;

.

* x 109

m&*x

oca* x ace; x ,olo

X

mC0; I lo7

mCa*x mco3 x 109

-34

- I.5

-32

-1.4

- 30

-13

-2.8

-

,.e

- 2.6

-

1.1

- 2.4

- 2.2

8.0.

21 I

22

23

24

4m/Lc-

Fig. 8. Relation to changing chloride of aCa2f x aCOazC, of mCa2+ x VLC’O~~-, and of ratio to chloride of each, averaged for water samples from groups of stations over the banks west of Andros Island and in the Straits of Florida. (Computations assume T 25°C and atmospheric

pressure).

can be seen that the activity product estimates here suggested are actually simpler to make than the empirical quasi-stoichiometric approximations, and they provide closer comparison with thermodynamic data from other fields. It will be recognized, to be sure, that the results indicated derive from an extrapolation of the Debye-Hiickel limiting law far beyond the low ionic strengths at which it begins to break down; so the nice check with experimental data could be purely fortuitous, and the method needs independent and chemically better qualified evaluation. Although detailed discussion of the procedures employed is not repeated here, some features of the method in particular call for elaboration and an attempt at justification. In reaching the values tabulated in Table 2, I/HCO,- and yCO,2were computed from relations expressed by the Debye-Huckel limiting law, while yC’a2+ was estimated by the mean-salt method. This procedure is open to criticism on the grounds that different results are obtained from the use of either DebyeHiickel or mean-salt procedures uniformly, while Debye-Huckel procedures in any

Behnviour of calcium carbonate in SC&water

877

case are theoretically inapplicable. Whereas mean-salt computations give inconhowever, the same method gives the gruous results for yCO,*- and yHCO,-, apparently most reasonable results for yCa 2+. In contrast, the activity coefficients calculated for carbonate from the Debye-H~cl~el first approximation coincide with those obtained by substitution in the titration curve for H&O, in sea-water according to the method of SVERDRUP et al. (1942, p. 204-205). The so-calculated activity coefficients also check with back-computations based on the assumption of approximate CO, equilibrium with the atmosphere at the particular stations concerned. .At critical station G7, for instance, yCa2+ by the mean-salt method is 0.264, whereas the Debye-Hiickel2d approximation formula (KLOTZ, 1950, p, 329, 21.126) gives the lower value 0,288; an intermediate value of 0.233 would give Karasonit,, precisely the same as that found by LATLMER. Because the “Debye-Hiickel” e&ima&ion of $0 32- checks out so well, and inasmuch as it is also unlikely t.hat the value k’ aragonitc estimated at G7 was either precisely at or below LATT~~~E~‘~value, the mean-salt method is preferred for yCa”* because it give more credible composite results. However, the neat approach of values computed from the Debye-Hiickel Sd approximation also implies support for Gbe mean-salt $a2+, and Debye-Htickel methods could be used throughout the computations with results not very divergent, from those obtained. An explanat.ion for this unexpected concordance is lacking, uniess it be the fact that the principal complexing effect, that. of carbonate with hydrogen, is already accounted for in the computations giving carbonate values. SUPERSATrTRATIOK If t,heminimal computed activity product of 7.8 s 10e9 does approximate J&@@t-: 10-9 computes to a supersaturation of about 100 per cent for aragonite instead of the 200 per cent given by .Karagoniteand ionic product values for the same sites. Comparable figures for calcite are 250 per cent from Ir’ and 450 per cent from fi’. Loss of analytical calcium from bank edge to quasi-equilibrium at island margin, moreover, is in reality only about 5 per cent of the original total in solution. Thus it seems not unreasonable t-o think of the apparent supersaturation observed as a kinetic phenomenon, amplified by some complexing of calcium and carbonate Jr&h ot’her ions. Discussion on the latter point has focused on the fact Ohat both the amount of COS2- in the alkalinity complex, and the product Caz+ x COS2- (SVERDKCP etr&Z., 1942, Big. 42) increase with rising pH. Because one member of the couple (that is COS2-) increases, it seems obvious that the product should also increase if no precipitation takes place. The anomaly is that the calculated ionic product does increase without precipitation, and at constant temperature and salinity. It was recognized by SVERDRUP et al. (1942, p. 207), and has been repeatedly since, that this suggested complexing, GARRELS and DREYER (1952, p. 333-336)*, therefore. worked out a method for computing amounts of calcium carbonate theoret.icalIy present in both ionic and non-ionic states under specified conditions, the non-ionic phase being taken as dissolved CaCO,. Inasmuch, however, as similar computjations

P.E.

878

CLCWD,

JR.

(courtesy of R. M. GARRELS) and experimental data indicate only trivial or no nonionic calcium for Bahaman waters studied, it looks as if complexing of calcium may not be an ~rn~ort~aut factor. Much more signi~cant apparently is the complexing of carbonate with sodium and ma,gnesi~l~, as recently elaborated by G,ARRELS et al. (1959, 1961). In this same connexion it is noteworthy that the few precise calcium values for bank waters nearly coincide with those predicted from the ratio of Wattenberg (SVERDRUP ~?!t nl., 1942, p. 196): C!a,z!+mg atom/l.

= f/2 alkalinity

me/l.

x 04.65 Cl- mgjl.

Such consistency in an environment of active CaCO, withdrawal provides basis for a broad extension of the field survey of calcium anomalies in the s@a, based on simple shipboard analyses for chloride and alkalinity. Inasmuch as WATTENBERG’S ratio using alkalinity and chloride for computation of calcium gives results approximat,ing the best analyt.icnl values in both straits and bank water, it seems to be valid as a means of estimating approximate cafcium present even @er precipitation (or solut’ion) of caIcium carbonate has taken place. Difference between calcium from the simple chloride ratio and that from the WATTENBERG ratio should, then, give an approximate measure of ionic oalcium lost from or added to the sea, without reference to solubil~ty product, real or apparent, and without calcium .%lldyses. CAUSES

OF" THE

PARTICULAP~

MISERALOCICAL

STATE

pressure, oalcium carbonate may crystallize from so.Lution as calcite, aragonite or vaterite -respectively its stable, metastable and highly unstable forms. These different polymorphs appear in different solutions or in the same solution under varying conditions. The controlling factors, however, remain in dispute despite many thoug~~tf~~l studies of the problem (e.g. JOEWSOK et al. 1916; LOWENSTAM, X954a, 1954b; POBEGUIX, l.954; TOGARI and TOGART, 1965; MOXAGHAN and LYTLE, 1966; ZELLER and 'Vlr~au, 1956; WRAY and DASIELR, 1957). The investigations here summarized support the thesis t.hat the demonstrable facts, and most of the apparent anomalies, can be harmonized by regarding the mineralogical state of the precipitate as primarily a ftrnction of degree of supersaturation of the parent solution and thus of its supersaturation kinetics. This view was anticipated by I)OBEGurN (19Ei& especially pS 95-99, 107; 1.955) from her painstaking work on pure-water solutions, and it was approached by ZELLER and WRAY (1956, p_ 149) and by WEAY and DAXIELS (IB57f. It accepts the now generally recognized facts that, other factors being equal: the formation of aragonite is favoured by increasing temperature and PH. Rut it incorporates these in the broader concept that any and all factors favouring the attainm.ent of high apparent supersaturation favour the initial precipitation of the more soluble, higher energy polymorphs. Among such factors are increasing temperature, increasing couceu~rat~on~ increasing alkalinity at fixed pH, increasing pH, At

25°C and atImosplreric

* GRUELS and DREYERwrite of the ionic product in Fig. 42 of SPERDRCP and others as f(‘, whereas SVEB~BUPand others correctly differentiated between their kind of K’~~II~~~, and the product Ca”+ x COS2-. In fact this does not affoet the argnment,yof CARRELS and I~RETER.

879

Behaviour of calcium carbonate in sea water

accelerated photosynthesis, variations in rate of diffusion, decreasing pressure and removal of carbon dioxide for other reasons-any or all of which operate individually or in combination under given circumstances. Decreasing activity of the combining ions with increasing ionic strength of the solution up to the range of sea water and the body fluids of organisms may play a governing role (Fig. 9; Table %; GARRELS et al., 1961, Figs. 3 and 4). Drawing now mostly on evidence detailed elsewhere (CLOUD et al., in press), t’he route by which this interpretation was reached may be summarized as follows:

3.28 0.27

-

x

--‘__X_

---

q

----x

0.26 0016-

3.25 0.0150.014 0.0130.012 -

0.011 -

Q p Groups of stations, number in parentheses + x lndwrdual stations

I I I I 0.0101 0.80 0.72 Fig. 9. Change in compbed

1

I

I

I

I

I

I

,

‘+I

i

I( )C)JJ 0.90 activity coefficients of CO,*- and Ca2f wit,11change in ionic strength of sea water.

1. Supposedly conflicting values for the apparent solubility product constant when crystallography and free energy K’caco, in sea water become compatible relations are considered. The computed ionic concentration products show a parallel relationship to computed activity products, and the apparent equilibrium values of both indicate standard free-energy differences for the transition from aragonite to calcite that compare well with those from experimental work. 2. R#atios to chloride of the concentration products and activity products of calcium carbonate across the banks west of Andros Island (Figs. 6, 7) curve regularly downward with increasing salinity from a minimum of perhaps 100 per cent supersaturation toward the apparent level of metastable equilibrium for aragonite. The system is thus regular, predictable and measureable, and it says that the sea is supersaturated with respect to both aragonite and calcite in the areas where aragonite is precipitating,

880

I?. E. CLOUD,JR.

3. Upon precipitating CaCO, from Bahaman sea water (Station Al’) by slow washing-out of CO, during 9 months at laboratory temperature, first aragonite (Sept.-Feb.) and then calcite with subordinate aragonite (Feb.-June) was obtained (data by courtesy of P. BLAC~MON). 4. Experimental precipitation with simple solutions showed correlation of mineralogy with temperature, but not with barium, strontium or other impurities. 5. Artificial stalactites made by POBEGUI~ (1955) from simple pure water solutions, however, showed no correlation between mineralogy and temperature. She obtained either calcite or aragonite at 9” to 13”C, depending on rate of evapIf the solution was allowed to drip slowly, oration or diffusion of the solutions. if it was retarded to a slow oozing, aragonite formed. calcite was precipitated; ELLIS (1959) also found diffusion or desorption to be a rate-determining (saturationcontrolling) process. 6. Determinations of Dhe mineralogy of calcium-carbonate sediments precipitated in open water bodies of low ionic strengt’h seem to be primarily calcite Aragonite reported in sediments of fresh water wit,hout respect t’o temperature. lakes (TWENHOFEL and MCKELVEY, 1941, p. 840) is perhaps mainly shell detritus. Sediments from a fresh-water lake on Andros Island itself proved to be calcite (NEWELL and RIGBY, 1957, p. 61, pl. 16, Fig. 2). X-ray analysis shows a calcite mineralogy for deposits collected from Unter Lunzer See in Austria (see also GoT~I~GER, 1912). 7. Naturally precipitated aragonite in the range of atmospheric pressure and temperature appears either to be restricted to solutions of relatively high ionic strength such as the sea, some salt lakes and the internal fluids of organisms, or to form as cave and other dripstone deposits where rates of diffusion become a controlling factor (POBEGUIN, 1955). 5. The irregularities observed diminish, and the regularities become part of a consistent whole if viewed as an expression of OSTWALD’S rule of successive reactions (OSTWALD, ‘1900, p. 447-445; FINDLAY and CA~IPBELL, 1938, p. -49-50; EWEL, 1954, p. 589-590, 628*). OSTWALD held that the release of solid polymorphs from an unstable solution takes place stepwise, from the least to the most stable solid form that can precipitate from a given initial concentration. Stated differently, this is to say that equilibration of a solution that is supersaturated with respect to polymorphs of the same chemical composition tends to take place (assuming no extraneous effects) by the smallest possible successive energy losses. The existence of entropy barriers (GLASSTONE et al., 1941, p. 99-100; GOLDSMITH, 1953) can create gaps in the sequence and retard the reaction at metastable levels. In nature, the kinetic factors that permit attainment of apparent supersat,uration for aragonite seem to be strongly influenced by increasing salinity of the solution, up to an ionic strength of about one, and by variations in rate of diffusion. Reduced to its simplest terms the generalization here proposed-one that could be deduced, to be sure, from free energy relations alone-is that aragonite is the expectable precipitate from solutions that are supersaturated with respect to bot,h calcite and aragonite, whereas only calcite should form between the saturation levels for the two mineral species. Trace elements of suitable ionic radius are

Behaviour of calcium carbonate in sea water

851

presumably trapped in the crystal structure of the host mineral as traditiormlly supposed and do not exercise a control over primary mineralogy. It remains to be explained why ZELLER and WRAY (1956) and WRAP and DASIELS (1957) found the sequence of formation in some of their experiments to be vaterite-calcite-aragonite, why aragonite and calcite form successive layers of shell in the same molluscs, and why some aragonite persists unaltered for long periods of time. It may be germane that apparently neither aragonite nor calcite have ever appeared before vaterite in any published experiments. In those tabulated by WRAY and DANIELS either vaterite or aragonite appears with or before calcite, with one exception where calcite preceded aragonite. Thereafter calcite increases at the expense of the accompanying species with aging up to 18 hr. These results could be reconciled with the preceding conclusions if the calcite accompanying vaterite in the absence of aragonite, or preceding aragonite in the apparent absence of vaterite, were being formed by rapid direct changeover from vaterite. The great instability of vaterite and its crystallographic relations to calcit’e might cause bypassing of t,he aragonite step if the yet unknown but presumably small free-energy loss from vaterite to aragonite were outweighed by crystal kinetics. As for the rest of this experirnental system, if the instability of aragonite were increased by the high temperatures and simple solutions, much of the calcite might be due to changeover from aragonite itself. Variant diffusion rates and crystal seeding are other possibly complicating factors. The influence of biologic factors on crystallography especially needs more study. The work of POBEGUIN (1964, p. 90-91, 97, 101) indicates that it is primarily a metabolic, and, above all an enzymatic function, but there might also be some relation between degree of saturation of the body fluids and calcium carbonate crystallography. Studies by ROCHE et al. (1951) of the shells of three pelecypod species that secrete both aragonite and calcite layers suggest that even this takes place in contrasting chemical environments, for they found higher concentrations of glycine and tyrosine in the calcite layers than in the aragonite layers of the same shells. TOGARI and TOGARI (1959) found a relation between shell mineralogy and chemistry of the ambient waters. The fact that Bahaman aragonite needles thousands of years old are still pure aragonite, whereas in other circumstances alteration of aragonite to calcite takes place in a few hours, days or months indicates marked differences in degree of stabiliby of aragonite under different circumstances. Geological evidence strongly implies that aragonite at low temperature and pressure will remain aragonite indefinitely (even in minute particles) in contact with solutions similar to those from which precipitated (or dry sealed), and that alteration to calcite of submerged marine sediments indicates former exposure to fresh water or the moist atmosphere. Included trace elements may have a retarding effect on the inversion of aragonite to calcite in a dry or nearly dry environment (at atmospheric temperature and pressure). Replicate X-ray determinations by BLACKMON (in CLOLTDet al., in mess) * This was originally suggested to me by Dr. S. 8.

GOLDICH.

582

P. E. %XJD,

JR.

on pure, ground, silt-sized, vesuvian araganite used for instrumental calibration showed 10-20 per cent inversion to calcite over a four-year interval. In contrast, washed and similarly stored aragonitic sediments from the Bahama Banks showed The slightly higher strontium and lead no inversion over the same interval. content and relat.ively high magnesium content of the Bahaman aragonite particles (and possibly the presence in them of sulphate ions) may bear on their slowness to It is contradictory but not necessarily fatal to such an explanation for invert. long-term retarding effects that M~GDONALD (1956) found no stabilizat.ion of aragonite from inchxded isamorphous casbonates below 30 mole per cent at high ~em~erat~~res and pressures.

R. N. Gr~sm.~~a: There seams to be no corrolatiion between the various types of non-skektal precipitates, ooliths, hardened fecal pellets and fibrous aragonite cement and 13:ater eondit‘ions. These fatt.er ooeur in areas of normal salinity and temperature for t,he banks. Author’s response: Inasmuch as calcium-carbonate precipitation may take place ahore or below the depositional interface, and in the open water or in cantact with or within organisms, we must recognise a variety of Huid conditions at any given geographical site. Then we can see that t,here is a oorrelation between t,he kind of calcium carbonate and fluid conditions. Formation off? brous asagonite cement and haxdening of fecal pellets presumably t,akes place below t*he &positional interface and is a function of ~te~~,i~~l or even nl~~roen~ironn~ental wat,er conditions, The chemistry of the interstital waters is related to a variet,y of factors other than the overlying waters, especially to bacterial populations which, in turn, are closcl~ relat’ed to grain size and texture of the sediments. The o(ilitic sands, as you know, are most provrxlent in areas of st#rong current flow, especially near the bank edge. If enough is yet known to warrant a general theory for their origin, holvever, it hasn’t come to my attention. J: don’t find it unreasonable to relste oiiids to o~asion of carbon dioxide from waters that are characterized by greater turbulence than others. On the other hand, I do not myself reject the possibility of biologic aggregation ar precipitation far some, nor can I cite a reason for discounting the prospect &hat some oalitia se&men& may be concentrated from originally dispersed o6id8 in finer sediments by winnowing processes. In any event, t,he relatively large nuclei at the cent)er of most of the Bahaman oijids suggest that t#hey too were formed mainly at or beneath the depositional interface and thus could in some degree be ciependent, on the into&&al water. The skel63tal precipitates, of course, arc related to their own particular biologic milieu. The needle-muds have been reported so far as I know only from areas of restricted cirrulation and relatively high salinity. As you know, it is my opinion that t,he aragonite needles which characterize them are due mainly to chemical processes that ta,ke place within Lhe wat,er mass above the depositional interface, and that t#here is a correlation between the needle-muds and the environment where they now occur. I believe thnL my dat,a substa,ntiate and require this interpretation. G. Y. CHZKNGAR: The eal~ium-magnesi~~~~ rat.ios of ealeareorrs sediments were determined by the writer along two traverses. The calcium--rHagnesirrm r&o seems to increase on going away from the shores of Andras Island. The contributor is attempting to explain this general t,rend. Author’s response: My colleague, PAUL BLACKMON, found that magnesium was contained in t.be 5-12 per cent of the total sediment represented as calcite and t.hat, a given calcite was either relatively high-m~esian calcite or relatively loafs-magnes~an eatrite+ The low-magnesian calcite included bedrock detritus and some skelete,l material. The high-magnesian calcite was all skeletal. %Ve found that skeletal material and therefore magnesium generally increased toward the bank edge, away from Andros Island, and southward toward more actively ciroulmt,ing wat,ers.

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of calcium

carbonate

in sea water

853

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POI~EGUIN T. (1954) Contribution a 1’6tude des carbonates du calcium, precipitation du calcaire par les &g&aux, comparison avec le monde animal. Ann. Sci. iVat. Bot. (11th s8r) 15,29-109. POHEGUIN T. (1955) Sur les concretions calcaires observees dans la grotte de Moulis (Ariege): C’.R. Acad. Sci. Paris 241, 1791-1793. ROBINSON R. A. and STOKES R. H. (1949) The role of hydration in the Debye-Hiickel theory. S.Y. Acad. Sci. Ann. 51, 593-604. ROBINSON R. A. and STOKES R. H. (1955) Electrolyte Solutions. Butterworths, London ROCHE J., RANSON G. and EYSSERIC-LAFON M. (1951) Sur la composition des s&roprot&nes des coquilles des Mollusques (conchiolines): C.R. Sot. Biol. 45, 1474-1477. S~IILTH C. L. (1940) The Great Bahama Bank-I. General hydrographical and chemical features. II. Calcium carbonate precipitation. Yale Univ., Bingham Oceanogr. Lab., Sears Found. Marine Res. J. Marine Res. 3, 147-189. SMJTHC. L. (1941) The solubility of calcium carbonate in tropical sea water. iWarine Biol. Assoc. U.K. J. 25, 235-242. S:VERDRUPH. U., JOHNSON M. W. and FLEMING R. H. (1942) The Oceans; Their Physics, C’hem,istry and General Biology. Prentice-Hall, New Hork. TOI:~~RIK. and TOGARI S. (1955) Conditions controlling the crystal form of calcium carbonate minerals (1); On the influence of the temperature and the presence of the magnesium ion. Hokkaido Univ., Fat. Sci. J. (ser. 4) 9, 55-65. TO(XRI K. and TOGARI S. (1959) Conditions controlling the crystal form of calcium carbonate minerals (2); Mineralogical study of Molluska. Hokkaido Univ. Fat. Sci. J. (Ser. 4) 10, 447-456. TWENHOFEL TIT.H. and MCKELVEY V. E. (1941) Sediments of fresh-water lakes. Amer. Assoc. Petrol. Geol. Bull. 25, 826-849. I~ATTENBERG H. and TIMMERMAN E. (1936) uber die SBttigung des Seewassers an CaCO, und die anorganogene Bildung von Kalksedimenten. Ann. Hydrograph. LWaritimen Meteorol. Z. fiir Seefahrt-Meereskunde 64, 23-31. \vB.AY J. L. and D~NIELS F. (1957) Precipitation of calcite and aragonite. A,mer. Chem. Sot. J. 79, 2031-1034. ZELLER E. J. and WRAY J. L. (1956) Factors influencing precipitation of calcium carbonate. Amer. Assoc. Petrol. Geol. Bull. 40, 140-152.