Carbonate-seawater interactions

a%oo~~etC~rn~~Acts,lQB8,Vor.

8O,pp.1087 to1048. PergamonPrec*LM. PrinkadinNorthemIrolnnd

Marine ScienceCenter,LehighUniversity,Bethlehem,Pennsylvania

ROBERT F. SCHNALZ Departmentof Geology,The PennsylvaniaStateUniversity UniversityPark, Pennsylvania (&&wed 9 Aecgust 1966; is rev&d fm

11 ilfwch 1906)

Al&r&--Interaotiootionebetween skeletal oarbonatesand seawateram infbnoed by miwral~gy, Experiments using pE-sensing teohniques p&i& sizeand the history of the particle surf-. indicate that the activity of magnesium calcites can be more than four times that of pure uaJeite. Part&lee of calcite IO- cm in diameterheve activitiesmore than 8 times greaterthan 1 cm

particles. Grindingin a mortarproduoosactivitiesas much BB7 timesgreaterthan unground materialof the samecomposition.

SEAWATERand other natural waters come into contact with a wide variety of calcium carbonate miuersls-calcite, aragonite and a spectrum of magnesium calcitess. The individual mineral particles am produced by a variety of biological, chemical aud physical processes, and occur as crystals and crystal a~ga~s of various sizes, and with surfaces, edges and corners possessing differing distortions and stresses. These three factors involving the solid phase-mineralogy, grain size and surface character-control the interaction of the carbonate crystals with the associated waters. Interactions between aqueous solutions and carbonate minerals can be studied experimentally by sensing reactions as changes in the pH of the system. Thus, specific water samples can be tested for over- or ~der-~t~ation with respect to a given carbonate phase by adding the solid and monitoring the pH change. Variations on this technique have been used by CARRELS et al. (1900), WEYL (1961), CHAVEet al. (1962) and SOEMALZand CEAVE(1963). The pH sensing method is based on the principle that in aqueous solutions having a pH near 84the pH of seawater and many other natural waters-the interaction between water, CO, and calcium carbonate minerals can be described by: CaCO, + Hz0 + CO, = C3+ + 2HCO,(1) At pH’s near 8.0, HCO,- is the principal CO,-derived ion in solution. Subtracting from (l), the dissociation of carbonic acid in water (2): H,O + CO, = H+ + HCO,-

(2)

CaCO, + H+ = Ca”+ + HCO,-

(2)

produces equation (3), an expression involving the sensor, H+ or pH. 1037

1038

K. E. CHAVE and R. F. SCA~AZZ

AS CaCO, is added to an undersaturated solution, a portion will dissolve and the pH will rise as reaction (3) shifts to the right. The addition of CaCO, to an oversaturated solution will produce the opposite change by providing nuclei for crystal growth. Thus, it is possible to establish the initial saturation state of a solution by monitoring the pH change upon addition of &CO,. Increase in pH reflects initial unders&turation; decrease in pH reflects oversaturation and uo change implies equilibrium. It is suggested by equation (3) that the pH change, from initial state to equilibrium, should be a function of the degree of original over- or under-saturation with respect to the added carbonate phase. In the oase of pure calcite and aragonite. where accurate thermochemical data are available, the relation between pH change and saturation state should be calculable. Unfortunately in experiments involving natural waters, particularly seawater, three major uncertainties are involved which prevent direct calculation of this relationship. These uncertainties relate to (a) reversibility of equation (3), (b) equilibration of gas phase CO, with dissolved CO, in the reaction vessel, and (c) the change in calcium activity in the solution during the experiment. Natural skeletal carbonates are commonly unstable phases, aragonite and magnesium calcites (MII(XEN, 1903; LOWENSTAM,1954; CRAVE, 1954; and others). The reversibility of (3), particularly in a solution with a high Mg/Ca ratio such as seawater, is diEcult to predict and difficult to measure. In the ease of oversaturation, aragonite nuclei may or may not be seeds for aragonite precipitation. Calcite nuclei may be seeds for caloite, magnesium calcite, or perhaps aragonite precipitation. However, for undersaturation, LAND (1966) has shown that in both distilled and seawater at 2577 and 1 atm CO, pressure, over a period of up to 24 hr, magnesium calcites dissolve congruently. The rate of equilibration of dissolved CO, in the reaction vessel with CO, in the gas phase, as this material is produced or consumed in the reaction, is determined by exactly how the experiment is carried out. ~though CO, appears to go into solution rapidly, expulsion from solution is reasonably slow (KERN, 1960). Recent measurements of the rate of CO, evasion from aqueous solutions (BUBECK,1965) indicate that even in a stirred solution many hours may be neoessary to achieve complete equilibration between the atmosphere and a solution containing an excess of CO%. This observation assumes particular significance under conditions where reaction (2) moves to the left, releasing CO, to the solution. In the YJaturometer Method” (WEYL, 1961; SCHXALZ and CHAVE, 1963), where the pH-sensitive eleotrode is surrounded by the in~rstitial fluid of a static settled sediment, CO,equilibration is by diffusion and is probably quite slow. The equilibration rate can be increased by bubbling air or a CO, mixture through the carbonate-water slurry. This method has the drawback, at least in the seawater system, that, as pointed out by BAYLORet al. (1963) and RILEY (1963), prolonged bubbling of seawater extracts dissolved inorganic and organic substances, aggregating them into discrete particles which changes the composition of the system. A compromise, a stirred slurry with a high water to solid ratio, such as used by GARRELSet al. (1960) and CHAVEet al. (1962), has been used in this study. Calcium activity in seawater is quite high-about 10-2’8aclcording to GARRETS

1039

Carbonate-seawater interactions

and Trronapsox (19&Z)-and changes in it, due to solution or precipitation of CaCO, phases, are very small, perhaps negligible. Calculated values of calcium activities under the conditions of the experiments are of limited value because of uncertainties in the CO, activities mentioned earlier. Thus the calcium ~ce~~ty remains. On the basis of the above three uncertainties it appears most reasonable to determine an empirical relationship between pH change and carbonate-water interactions, and to compare these results with thermochemical values in order to evaluate the ma~itude of the uncertainties. CALCITE-ARACJONITE AT 25% AND ONE ATMOSPHERH PRESSURE When finely ground carbonates are added to water in a proportion of approximately 1: 100 weight solid to solution, and stirred, the reaction proceeds to an apparent steady state in a few hours. That is to say, in most experiments, no pH change could be observed within the accuracy of the method after 3 to 4 hours. When dealing with seawater it is normally impossible to continue experiments longer than this because of changes in the water, resulting from evaporation or condensation, or biological activity, which are independent of the carbonate minerals. The pR values, presented in the remainder of this paper, are, unless otherwise indicated, these steady state values, and are accurate to fO.02 pH units. Differences in final pH between experiments run simultaneously are probably precise to &O*Ol pH units. Most of the experimental results reported were obtained using a Photovolt 180, battery powered, expanded scale pH meter, standardized with two buffers. In a few cases a Keithley 610B electrometer with Leeds and Northrup pH electrodes was used. Table 1 presents steady state pH values for simultaneous experiments, at 25-26% and 1 atm total pressure, using various calcites and aragonites. Results for distilled water and seawa~r are presented. The solid materials used in these experiments were crushed to pass a 64 ~1seive, then aged and, in most cases washed for several hours in distilled water, in order to minimize anomalous high-solubility effects discussed in a later part of this paper. The ApE values tabulated for each experiment correspond to the logarithm of the ratio of steady-state hydrogen-ion activities: (aHiL***ita

1

@.zH&,~~~~ ) 1 = log

bak.Me 1

[ (aH&,,

1’

>

The average value of these ratios yield the “average” steady-state ApH value shown in the table (As = 0.087). This expe~ment~y determined value for A3 can be compared with a value calculated from thermochemical data by consideration of the equilibrium constant for equation (3). At 25’C and one atmosphere total pressure, the equilibrium constant for calcite in (3), K~afcslc, is: K (3)CdC =

(aCa2+)(uHC0,-) (aCaCOa)(aH+)

where a denotes the activity of the species.

(4)

1040

EC. E. CEAVE and R. F. SGIIMALZ

Table 1. Steady state pH for &cite

Solution (1) (2) (3) (4) (5) (6) (7 )

Distilledc Distilk& Seawater Seawater Seawater Seawater Seawater

Temp. (“C)

Salinity (%I)

25 25 25 26 26 25 25

-36 36 36 31 31

and aragonite at about 26°C and 1 atm totat pressure

PCO, 1()-0.01a

K-r; 1O-9.5 10-3.5

IO-0.01 10-0.0’

pH calcite --

pH Aragonite _----

ApHb

6.015 8.43 7.91 8.08 8.10 5.92 5.91

6.09 8.52 x.00 8.18 8.18 6.02 599

0.075 0.09 0.09 0.10 0.08 0.10 0.08

Average vulue, As

= 0.087

(1) Average of 11 values from CECAVE et al. (1962). (2) Calcite-reagent, low in alkali, aragonite-MW;pora. (3) Cal&e-spar, Jacksonburg fm., aragonite-b.GlZe~ra. (4) Cal&e-spar, Ja~nb~ Rn., arago~t~~~o2~. (5) Cak&e--stalagtite, Bermuda cave, aragonite-C~~0~ol~. (6) Calcite--reagent, low in alkali, aragonite-Tellina. (7) Celoite-reagent, low in alkali, aragonite-.4cropor. a 1O-o’o1 ia taken as the CO, partial pressure in H,O-saturated CO, at WC, L atm total pressure: 1O-s’s is the partial pressure of CO, in average air at 1 atm total pressure. b ApH is the difference between steady state pH values, fpH aragonite-pH calcite). It is equivalent to the logarithm of the ratio of hydrogen ion activities

c Because these values were determined in distilled water, the assumption that the ratio of ca;lcium ion activities equal unity may not be strictly valid.

The standard free energy change of this reaction, AG,‘, is: AG,’ = -2-63 kcaljmol

(5)

where the data for the standard free energies of formation (AC,“) are from LANG+ MUM (1964). AC,’ is related to KCstcalc by: AG,’ =

-RT In K

(61

where R is the gas constant and T is the absolute temperature. At 25% and 1 atm total pressure ACT,”= - 1,364 log R (7) Thus : l-cf(l)cs~a= 10*1+3

(81

Similarly, Kt8jm 08x1be oak&t&d with aragonite as the solid phase even though it is m&a&able: AG?“ = -2.86 kcal/mol (9) and, for (6) : K t8)mg = 10f2*“0 so

(191

Carbonate-seawater interactions

1041

To rel&te these v&lues to the equilibrium pII values, consider r&tio ( I 1), denating activities in the equation which involve &r&go~~w&~r with an asterisk *

From the first ionization const&nt for carbonic acid (2): (aHCO,-) = ~ub~itut~g

(~H~o)(~C~~) K,,, (aH+)

(13)

(13) into (12) produces:

K f8mr (~cas+*)(aH~O*)(ocO,*~(~H~~~(~C&~~~ rr,, -=2 K (S)calo (uC&~~)(~H,O)(~CO,)(~H+*)“(~;C&CO,)K(,,

(14)

K&king the convention&l &ssumption th&t the activity of e&oh solid is unity, and that the ratio (~H~O*)~(~~~~~ is &Is0unity, this reduces to:

&n equ&tion which contains, in addition to (aH+), two u~~rt~~ties mentioned earlier, nemely (aC0,) &rid (aCa&+). The rrttio of equilibrium constants m&y now be oalculated either from thermoohemioal d&t& (equation 11) or from experiment&l d&t& using equation (16). The~o~hemi~&l data yield & v&lue, c&loul&tedearlier, of 10+““7. ~batituti~g the &ver&gev&lue of AhpRfrom Table 1 into equrttion (16) gives:

The similarity of these values, 10+0’17from thermochemical d&t&&nd 1O+s’*74 from experiment&l solubility d&t&suggests that those terms inequation (16) rel&ting to the unce~~nties mentioned e&rlier m&y be neglected, at leaet for calcite and &r&gonite,and that & useful empiric&l relationship between ApII &nd solubility is given by: K @f=E_ (f-SH*)” (17) K WC&k (aH+*)” ACTIVITY OF SOLID C&O,

When de&ling with e&sily ~stinguish~ble &ridphysic&lly sep&r&blesolid forms h&ving different energies, &s in the c&se of c&Me and &~agonite, it is correot and convenient to me&sure the free energy of formation of e&ah phase and work with &~~op~&~ eq~b~um #~st&nts. In many natural systems, suoh ss those desoribed below, & v&riety of st&ble snd unstable solids &redealt with, representing &n energetio oontinuum. IndividusLlphases m&y be iusep&r&bleand their energies m&y be transient and not easily defined. In such ciroumstanoes two or more physically and ohemia&lly distinct phases m&y h&ve the s&me energy oontent. Further, quite similar phsses m&y h&ve measur&bly different energy oontents.

1042

K. E. WAVE and R. F.

fkEiMALZ

For these reasons, it appears most convenient to treat metastable phases in this system as “hyperactive”, and to define their activities relative to that of the standard phase calcite, arbitrarily defined as having unit activity. Calcite is the thermodynamically stable form of CaCO, at 25°C and 1 atm total pressure (equations 5 and 9). Kt3jcaic can then be defined as the equilibrium constant for all reactions involving carbonates of calcium and water under these P-T conditions, with differences in solubility being ascribed to differences in the activity of the solid phase. Thus, (18) can be defined: and from (12) and the equations which follow: (aH+)2 (aCaCO,*) (aCaC0,) =(aH+*jz*

(19)

If (~~aC0~) for calcite is defined as unity, the activity of another solid CaCO, becomes, within the limits of the approximation:

(20) or: log (a~a~O~*) = Z(pH* - pH) = 2ApH

(21)

Using this approach, the average activity of aragonite from Table 1, would be 1.50. FACTORSAFFEGTI~~THE A~TW~TY OF SOLID CALCIOB~ ~~30~~~~s The activity of a solid carbonate, at a given temperature and presmre, according to the above defkition is determined by its mineralogy, partiole size, and the charaoter and history of the particle surfacteexposed to the solution. The experimental data presented above, for calcite and aragonite, are in suffiaiently close agreement with theoretioal values to encourage attempts to evaluate these other effects. (a) Hineralogy LAND (1966) has shown that magnesium calcites, reacting with distilled water or seawater for periods of up to 24 hours, show no measurable departure from congruent solution. In essence, on a short term basis, reaction (3) can be written: XCO, + H+ = X2+ + HCO,-

(3’)

where X is any combination of calcium and magnesium up to about 25 mole per cent magnesium. Thus, magnesium caloites can be treated in the same way as aragonite, and equation (21) oan be used to estimate the aotivity of magnesium calcites relative to the stable phase calcite. The data in Table 2, derived in part from CHAVEet cd. (1962), are s~dy-state pH and ~al~ula~ activity for a variety of organic&y-secreted magnesium calcites at 25”C, in both distilled and sea water. Table 2 shows a olear trend of increase in activity with inctreasingamounts of

1043

Carbonate-seawater interactions

Table 2. Steady-state pH and solid-phase activity, relative to calcite, for magnesium calcites at about 25°C.

Solution (1) (2) (3) (4) (5) (6) (7)

(8) (9) (10) ( 11) (12) (13) (14) (15)

Distilled Distilled Distilled Distilled Distilled Distilled Distilled Distilled Distilled Seawater Seawater Seawater Seawater Seawater Seawater

Sso

3663 33.43 36.48 3663 33.43 3650

Temp (“C)

Calcite mol % Mg

25 25 25 25 25 25 25 25 25 25 26 27 25 26 23

2 7 13 13 14 14 16 20 24 11 11 11 20 20 20

.

pco,

pH talc pH Mg-calc

10-0.01 lo-0.01 lo-0.01 lo-0'01 10-0.01 lo-0.01 10-0.01 10-0’0’

10-0'01 10-3.5 1@-*5 10-3.5 10-3~5 10-3.5 10-3~6

6.02 6.02 6.02 6.02 6.02 6.02 6.02 6.02 6.02 7.91 7.89 7.86 7.91 769 8.07

6.02 6.08 6.12 6.20 6.21 6.14 6.24 6.26 6.35 8.08 8.06 8.02 8.01 8.09 8.29

ApHa

cCaCO,*

0.00 0.06 0.10 0.18 0.19 0.12 0.22 0.24 0.33 0.17 0.17 0.16 0.10 0.20 0.22

l-00 1.32 1.58 2.29 2.40 1.74 2.75 3.02 4.57 2.19 2.19 2.09 1.58 2.51 2.75

(l)-(9) CHAVEet al. (1962), (lo)-(12) Lytcchinu.9,(13)-(15) Amphiroa. a ApH is the differencein steady state pH between calcite and the magnesium calcite.

magnesium in solid solution. With the exception of sample number 13, there is reasonably close agreement, as would be expected, between experimental runs in distilled water and seawater at different CO, partial pressures. (b) Particle size The surface energy of a crystalline solid is part of the total energy of the partiole and is thus a contribution to its activity and its solubility. The significanae of this surface energy increases with the increase of the surface:volume ratio of the particle. The total surface energy of a single crystalline particle may be represented by an equation of the form: E, = aAr,x2

+ bBr,x2

.. .

(22)

where a is the number of bounding faces whose surface area is defined by Ax2 and exhibiting surface energy TA,, b is the number of bounding faces whose surface area is defined by Bx2 and exhibiting surface energy TV, etc. In this equation z is a linear dimension of the partiole (minimum edge length, diameter etc.). Although the bounding faces of most crystalline solids are not energetically equivalent and equation (22) must therefore be used to calculate surface .energy, certain crystal forms are bounded by equivalent faces (e.g. cube, cleavage rhomb of calcite) and for others an average value for the surface energies of exposed faces may be obtained. In such cases equation (22) may be simplified to: E, = Cx2i:

(23)

where C is a geometric parameter relating the total exposed surface to the linear

1044

IL E. CHAVE and R. F. Scmz

dimension x, and ? is the average surface energy of the exposed faces. The surface energy contribution to the molar free energy will then be nE s = nCx%

(24)

where n is the number of partictles per mole. If all particles are assumed to have the same size and shape, the volume of eaoh partiole will be given by: V, = KX3 (25) where K is the approp~a~ geometrio parame~r relating particle volume to linear dimension, and the number of particles per mole will be given by dividing the molar volume by the particle volume: P n==3

Thus nE,=-

f;rci: Kx

(27)

where C/K ia dependent on particle shape. The molar free energy, CJ,,,, is related to particle size by: c mol

-“G”

;

‘r

WI

Where a0 is the standard &ate free energy of formation, r is the shape factor (C/K), i7 is the molar volume, F is the mean surface energy of the exposed faoes, and x is some linear dimension of the particles. Previous me~~ements of the surface energy of calcite (GILMAX,1360) suggest that the free energy ohange ocoasioned by a reduction in pa&Me size might have a signifioant effect upon the solubility of calcium oarbonate. In an experiment designed to test this hypothesis, it was necessary to measure the size of very small particles with high aoouracy, and to make precise measurements of ApH. The size of large particles oould be measured with the necessary aocuraey by means of a vernier ctaliper. The sizes of smaller particles were Gal~~a~d from their settling velocities in carbonate-saturated solution at the end of the experiment observed with the aid of 250x magnification and a micrometer ocular. Differential pH measurements (ApH) of the requisite precision were obtained with the aid of a Keithley Differential Electrometer (Model 603) with matched Beokman eleotrodes. The electrode pairs were reversed for replicate rne~u~~nts of each de~~~~o~ to eliminate spurious voltages developed by symmetry of the electrodes. Absolute accuracy of the pR values reported in -Table 3 is determined by the aoouraey of the bufFemused to standardize the electrodes ( &O*OZpH), but the acouraoy of the reported pH merenoes is fO401 pH or better. Although the absolute free energies reported may therefore be in error by a small amount, the free energy Merenoes are sign&ant. In the following experiments, a distilled water solution in eq~b~um with 1 cm calcite partioles was used as a reference solution, and the steady state differential pH was reoorded in solutions allowed to equPibr&e with o&&e partiobs

Carbonatiawater

1046

interactions

of lOA a-nd2 x 10” centimeters. All solutions were kept at 26% and were exposed to one rttm oarbon dioxide pressure. Because in distilled water solutions significant changes in the aotivity of calcium ion make equation (17) invalid, free energy differences were c&lcul&~d dire&y from the ApH values using the method of GARRELS et al. (1960), and activity ratios were then calculated from the relation:

AQfl -

AGfX”= 1.364 log

(29)

The resulting values are presented in Table 3. Table 3. Effects of particle size on calcite activity Size (om)

FinalpH ( fO+w

1

6.016

IO-4

6.02

2 x 10-s 10-a

6.04

APH

(&O.OOl)

a”

(kcal/mol)

0,005

0.01

0.025

0.08

Q0f

(kcal/mol)

(aC&O,‘)

-269+37

1

-269M

1.02

-269.79 -26E+58”

1.16 8~92~

@Calculated from the average V&.W of i.

From these experimental data it ia possible to calculete the value of ?. The results are as follows: for lOA cm particles, ?: = 188 ergs/cm* for 2 x 10” cm particles, d = 301 ergs/cm*. The increase in apparent surf~8 energy with decreasing size probably r8flects in part the increasingly important edge and corner contributions in the smaller particles. The average value of the surface energy, 244 ergs/cm2 is in very satisfactory .agreement with that determined for the caloite cleavage plane (10*4) by others. GILM&S (1960) reports & value of 230 ergs~om~. Using this average vslue of ? in equation (28) we may calculate it minimum value for the activity of csloite particles in the size range 1O-6 cm (Table 3). The results of the calculation suggest the 8normous increase in calcite solubility which may be expected as particles approach ,OsOlp in size. Although particles smrtllerthan 0.1 p( lo9 cm) will be quite unstable in aqueous systems, particles between 10m6and 1O-6 cm may be produced in environments (1) where CaCO, is precipitating, (2) where carbon& particles are subject to abrasion, .as in the surf zone, or (3) where organic integuments of fine grained skeletal parts are being broken down microbiologicrtily. In such environments the resulting highly active grains may dissolve rapidly.

Mechanical strain and strain-induced defects in the crystal lattioe may represent a substantial contribution to the total energy of a particle. In an extreme osse, the energy introduced by grinding, emph&sized perhaps by abr~ion~l hesting, mrty be

1046

B. E.

CRAVE

and

R.

F.

SCHMALZ

sufficient to effect a phase transition. Calcite can be inverted to aragonite, for instance, with a mortar and pestle (JAMIESONand GOLDSMITH,1960). The energy added to a crystal during grinding increases its activity and its solubility. The effect of abrasional energy on carbonate activity can be evaluated by comparing the equilibrium pH of freshly ground carbonate particles with that of particles of a similar size which have been annealed at SO--100°C for 30 days The data, in Table 4 compare steady state pH values. at 25°C 10-3’5 atm PoO, in artificial seawater of salinity 35 per mil, for freshly ground and annealed optical calcite. In each case the annealed value is used as the standard state for the calculation of (aCaCO,*). Table 4, Effect of grinding on calcite activity: 25”G, ~CI-~‘~atm CO,, artificial seawater Grinding time (min)

Size (P) (1) (2) (3)

177 > x > 125 125 > x > 88 88 > x > 63

pH ~~t,e~dy*s~ate) Freshly Annealed ground *PH 8.37 8.20 8.32

10 20 10

8~30 8.45 8.49

0.13 0.25 0.17

(aCaCO,*) 1.82 3.17 2-19

It is evident from Table 4 that considerable energy can be added to carbonate crystals by simple grinding, and that in just 10 min the activity of aragonite, 1.50, may be exceeded. The effects of particle size and mechanical strain cannot be easily ~stinguished in many environments, for example in surf zones where size reduction and strain are accomplished simultaneously. By grinding carbonates, wet, in a mortar, and monitoring pH, an activity value for the summation of the effects of size reduction and strain can be determined. Such values, presented in Table 5, may be Iow because the high degree of supersaturation of these slurries, with respect to less active carbonate phases, may have caused some precipitation and lowering of pH . Table 5. Activity of calcite and aragonite during grinclingin seawater at 1O-s’5 atm PC*%

Mineral (1) Calcite (2) Ax-ago&e (3) Ar&gonite

Source reagent ~~ M&%jwm

Temp (“C)

Salinity (%)

PH Steady state

12 12 15

29.79 29.67 36.63

7.80 7.99 8.19

Maxhmm 8,22 8.30 8.27

prz 0.42 0.31 O-08

(aCaCO,*) 6.92 6.31Q 2.19a

0 Relative to calcite and therefore includes (o Aragonite = l-50).

The large activities observed in these experiments--aCaCO,* = 6.92 is equivalent to 1145 cal-suggest that such particles are short-lived in aqueous media However in the sea, in the surf zone, where such particles are continuousIy produced, solution will take place and the saturation state of the seawater will be affected. We have observed that on Bermuda beaohes waters are reacting with a very energetic carbonate species and are oversat~a~d with respect to aE unstrained skeletal carbonate particles in contact with the water.

Carbonat+seawater interactions

1047

SUMMARY A?XD CONCLUSIONS

with chemical interactions between calcium carbonate minerals and aqueous solutions, including seawater indicate that a wide range of solubilities exist in this system. These variations result from differences in mineralogy, particle size and abrasion-induced strain. Because of the continuum of energies and equilibrium constants involved, it appears most convenient to ascribe the hypersolubility of the many phases to activity of the solid carbonate (aCaCO,*) relative to a standard phase, coarse unstrained calcite. It is well known that tropical and subtropical surface seawater is oversaturated with respect to many carbonate phases, and remains so for long periods of time. It has been further observed that shallow water carbonate sediments are undergoing selective solution of hyperactive mineral phases (CHAVE,1962). Both of these phenomena may be explained as reflecting interaction between the solution and hyperactive metastable phases with which it is in contact. CELAVE(1965) has shown, however, that in certain circumstances carbonate particles suspended in surface seawater do not interact freely with the water, possibly because of organic coatings on the carbonate grains. A further complication is introduced by LAND (1966) who has shown that on a long term basis, magnesium calcites do not dissolve congruently, but rather react with solutions to produce lower magnesium calcites having lower activities. When dealing with natural seawater and carbonates, both potential and real interactions must be considered. The degree to which the full potential of the interactions is realized is influenced by the environmental characteristics. The surf zone and other high-energy near-shore environments are probably natural environments where maximum interactions occur. Interactions in such areas may exert substantial control over the characteristics of surface waters over large distances. Experiments

Acknowledgenzenthis work was supported by the American Chemical Society P.R.F. Grants 651-A2 and 909.A2, and N.S.F. Grant GP3606. Contribution 66-2 from the Marine Science Center, Lehigh University; 66-63 Mineral Industries Experiment Station, The Pennsylvania State University, University Park, Pennsylvania; 388 Bermuda Biological Station. The writers are indebted to J. DAEN, D. R. SIXPSON,R. M. GARRELSand L. S. Lm for valuable discussionsduring the preparation of this paper. REFERENCES E. R., SUTCLIFFE W. H. and HLRSCHFELD D. S. (1963) Adsorption of phosphates into bubbles. Deep Sets Res. 9, 126-124. BUBEC~R. C. (1966) A kinetic study of the evolution of carbon dioxide from an aqueoussolution. Thesis, The Pennsylvania State University, 62 pp. COVE K. E. (1964) Aspects of the biogeochemistryof magnesium. I and II. J. Geol.62,266283; 587-699. CJXAVEK. E. (1962) Factors influencing the mineralogy of carbonate sediments. Limnol. Oceanogr.7, 213-223. CHAVEK. E. (1966) Cal&un carbonate: Association with organic matter in surface seawater. Science148,1723-1724. Cn~vx K. E., DEF~IZXESK. S., WEYL P. K., GA~~ELS R. M. and THOXPSONM. E. (1962) Observationson the solubility of skeletal carbonatesin aqueous solutions. Science187,33-34. BAYLOR

5

1048

Ii. E. CHAVE and R. P. SCKMALZ

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