Everglades Starfish map of south Florida

South Florida
Ecosystem Research
and Monitoring Program

Research

The Role of Suspended Calcium Carbonate in the Phosphorus Cycle in Florida Bay

Dr. Frank J. Millero
Rosenstiel School of Marine and Atmospheric Sciences
Division of Marine and Atmospheric Chemistry
University of Miami
4600 Rickenbacker Causeway
Miami, FL 33149
Tel: 1 305 361 4707
Fax: 1 305 361 4144
Email: fmillero@rsmas.miami.edu

Dr. Jia-Zhong Zhang
Cooperative Institute for Marine and Atmospheric Studies/AOML,NOAA
Rosenstiel School of Marine and Atmospheric Science
University of Miami,
4301 Rickenbacker Causeway
Miami, FL 33149
Tel: 305 361 4397
Fax: 305 361 439
Email: zhang@aoml.noaa.gov

NOAA GRANT NO. NA67RJ0149
DURATION: 1 April, 1997- 30 April, 2000
August 28, 2000

  1. Project title: THE ROLE OF SUSPENDED CALCIUM CARBONATE IN THE PHOSPHORUS CYCLE IN FLORIDA BAY

  2. Abstract:

    This final report summarizes the project funded by SFERPM to study the adsorption and desorption of phosphate on calcium carbonate particles and its effect on phosphorus cycle in Florida Bay. Kinetic experiments indicate that the exchange of phosphate between particle and seawater is a rapid process. Results of this study provide the essential data for modeling the dynamics of phosphate adsorption and desorption on the calcite and aragonite surface. Several cruises have been completed during 1997-99 in Florida Bay as well as off the Florida Keys (Hawks Channel), and the Shark River. During these cruises, the carbonate system parameters (including total alkalinity (TA), pH, total carbon dioxide (TCO2), partial pressure of carbon dioxide (pCO2), as well as salinity and nutrients were measured by our group. A flowing multi-parameter nutrient system, developed in our lab, has been used on several cruises for continuous monitoring of the nutrient concentrations. Nutrient data (nitrate, nitrite, phosphate, and silicate) have also been combined with the carbonate parameters to examine the relationship of the carbonate system and nutrient elements in the waters of Florida Bay. In addition, Chlorophyll-a data were collected on several cruises to evaluate the primary production of Florida Bay water. We will continue to participate on cruises in Florida Bay over the next year in coordination with other groups working in the program.

  3. Executive Summary:

    The results of 17 parameters investigated in our field study such as carbonate parameters (TA TCO2, pCO2, pH), salinity, temperature nutrients and chlorophyll et al. were all posted on our website. Anyone interested in Florida Bay study can assess and download the contour plots of these parameters. A paper focusing on a time-series study of CO2 parameters over the course of a year in the shallow waters of Florida Bay was submitted to the Bulletin of Marine Science. Our laboratory experiments of adsorption/desorption of phosphate on calcium carbonate indicate that the carbonate phases could be both a scavenger and a source of phosphate in Florida Bay water. Temperature exerts a significant effect on the adsorption of phosphate. The effect of salinity on the adsorption is relatively small and can mostly be accounted for by the bicarbonate concentration. Calcite and aragonite have similar characteristics in the long-term adsorption experiment. A manuscript on the adsorption of PO4 to CaCO3 in seawater was submitted to Aquatic Geochem. last year and is in press.

  4. Purpose:

    It is well known that increased nutrient concentrations and altered nutrient ratios can result in phytoplankton blooms and degradation of the water quality (Smayda, 1989, 1990). It is particularly important in Florida Bay due to the runoff of nutrients from the Everglades. As a result of increased fertilizer and sewage input, Si:N and Si:P ratios have declined in many coastal waters (Smayda, 1989,1990; Turner and Rabalais, 1991, 1994; Sommer, 1994; Justic et al., 1995). This tends to lead to a shift from diatoms (a preferred food source for many marine herbivores) to small flagellates (which are often poor foods) within the phytoplankton community (Schelske and Stoermer, 1971; Officer and Ryther, 1980). Florida Bay sediments are mainly composed of calcium carbonate (c.f. Table 1). Measurements of soluble reactive phosphate in the pore waters of calcium carbonate-rich sediments are much lower than in sediments with little or no calcium carbonate (Morse et al., 1985). Therefore calcium carbonate minerals may exert significant control on phosphate chemistry in marine waters (Ames, 1959; de Kanel and Morse, 1978). Phosphate has been shown to be the limiting nutrient in part of Florida Bay due to adsorption to biogenic carbonate sediments (Fourqueran et al., 1992). The retention of phosphate by carbonate affects the growth of primary producers such as seagrasses (Short, 1987). Some researchers have suggested that adsorption of phosphate on calcium carbonates may lead to the formation of apatite (Ames, 1959; Simpson, 1964, Martens and Harriss, 1970). In contrast, recent work by Louchouarn et al. (1997) indicates that high carbonate ion concentrations and the presence of calcite in marine sediments may inhibit the formation of carbonate fluorapatite.

    Table 1. The mineral composition of the Florida Bay sediments.
    Date Station Latitude Longitude Description CaCO3 [PO4] [Fe] Arag. LMCa HMCb
              wt % mmol/g mM % % %
    May 14/97 10 25.0172 -80.5597 white mud 95.2 0.51 1.0 31.1 9.9 59
    May 13/97 11 25.0788 -80.7477 Fine particle w/shell 85.8 1.80 0.11 31.9 10.5 57.6
    May 13/97 15 25.0463 -80.9167 fine particle 85.8 2.83 0.27 30.3 29.9 39.8
    May 14/97 21 24.9103 -80.661 Mud
    w/debris    		 93.8 1.26 0.71 23.6 16 60.4

    a LMC: Low Magnesium Calcite
    b HMC: High Magnesium Calcite

    The mechanism of the adsorption of PO4 (phosphate) to CaCO3 (calcium carbonate) has been studied previously by Stumm and Leckie (1970) and de Kanel and Morse (1978). Stumm and Leckie (1970) suggested that the initial uptake of phosphate on calcite occurs by chemisorption, followed by a slow transformation of amorphous calcium phosphate to crystalline apatite. The rate of apatite crystal growth depends strongly on the phosphate to carbonate ratio and the authors suggested that this is due to the competition for growth sites. In their experiments, they also found that fluoride increased the rate of phosphate uptake whereas magnesium strongly inhibited the adsorption. de Kanel and Morse (1978) investigated the kinetic behavior of phosphate adsorption on calcite carbonate and their results indicated that fluoride and magnesium have only a minor effect on the rate of phosphate uptake.

    While many studies have focused on phosphate uptake on calcium carbonate, few studies have examined the adsorptive behavior over a wide range of environmental conditions such as temperature and salinity. Little data are available for the desorption of phosphate from calcium carbonates. In this study, we examined the adsorption and desorption of phosphate on calcium carbonates (calcite and aragonite) as a function of temperature (5 to 45°C), salinity (0 to 41), pH (7.2 to 9.0) and solution composition. The adsorption and desorption on natural carbonate sediments from Florida Bay are also studied. The purpose of this work is to model the turnover of the phosphate on suspended CaCO3 in seawater, especially in Florida Bay waters that experience a wide range of salinity and temperature.

  5. Approach:

    Stock solutions of phosphate (1 mM) were prepared with reagent grade Na2HPO4 (Fisher Scientific). Seawater from the Gulf Stream (S = 36.4) (5 meters below the surface) filtered through 0.2 mm filters was used throughout the experiments. The phosphate levels in Gulf Stream waters are below 0.05 mmol kg-1. Since the seawater is supersaturated with respect to calcite and aragonite all of the measurements were made after a pre-equilibration of at least one hour. Reagent grade calcite was obtained from Mallinckrodt. The aragonite used in this study was synthesized at 70°C by the method of Wray and Daniels (1957) as modified by Katz et al. (1972). The equilibrium experiments were carried out in Nalgene HDPE bottles. Phosphate added to these bottles without calcium carbonate showed no noticeable loss due to adsorption.

    The natural carbonate sediments used in this work were collected in Florida Bay. Florida Bay is a triangular area on the southern tip of the Florida peninsula. The mainland borders it on the north and both of its southern and eastern portions merge into the reef ridge of the Florida Keys. It opens its western side to the Gulf of Mexico and receives fresh-water runoff from the Everglades. In all but the wettest years, it is a negative estuary, with hypersalinity caused by evaporation exceeding freshwater inputs. The salinity in Florida Bay can be as high as 50 (Fourqueran et al. 1992). The location and mineral compositions of the sediments are given in Table 1. Large debris was removed from the sediments by passing them through a 60-mesh sieve. The sediments were dried at 60°C and stored at 5°C before use. The fraction of CaCO3 in the sediments was determined by subtracting the acid non-dissolvable component from the total weight of the sediment after dissolving a portion of the sample in 10% HCl solution. After filtration through a pre-weighed 0.45 m Nuclepore® filter, the weight of the non-carbonate material retained on the filter was determined after the filter was dried at 60°C overnight. The phosphate and iron concentrations in the sediments were measured from the filtered solution. The iron in the solutions was determined using the ferrozine method (Stookey, 1970). The precision of the Fe measurements was ± 30 nM. All the sediments contained greater than 86% of CaCO3 and 0.5 to 2.8 m mol/g of PO4 and 0.1 to 1.0 m mol/g of Fe. The composition of the calcium carbonates determined by XRD was 24 to 32% aragonite, 10 to 30% low Mg calcite (LMC) and 40 to 60% high Mg calcite (HMC). The cutoff point between low Mg calcite was 10% and for high Mg calcite was 40%.

    Phosphate concentrations (PO4 = H3PO4- + H2PO4- + HPO42- + PO43-) in the solutions were analyzed using the phosphomolybdenum blue method (Murphy and Riley, 1962) at 880 nm on a HP 8453A spectrophotometer. The precision of the phosphate measurements was ± 0.1 mM with 1 cm length cell. Ammonium molybdate, potassium antimonyl tartrate and ascorbic acid were obtained from Sigma Chemical.

    For the adsorption kinetic studies, 5 mmol of phosphate was added to 500 cm3 of seawater in a 1000 cm3-capacity jacketed beaker maintained at 25 ± 0.02° C by a Neslab constant temperature bath. While vigorously stirring, 1 gram of either aragonite or calcite powder (<200 mesh) was added to the seawater.

    For the desorption kinetic studies, phosphate was pre-adsorbed onto calcium carbonate by equilibrating two grams of calcium carbonate (aragonite or calcite) with 225 cm3 of a 110 mM phosphate solution. After being subjected to either a 24-hour (short-term) or 1-week (long-term) equilibrium period, the solids were filtered through a 0.4 mm filter and dried at 60°C overnight. The total amount of phosphate adsorbed was calculated based on the difference in phosphate concentrations in solution before and after the addition of the solids. One gram of calcium carbonate (calcite or aragonite) pretreated in this manner was added to 500 cm3 of seawater.

    For both kinetic experiments, 10 cm3 of slurry were taken with a 10-cm3 syringe at time intervals from 1 minute up to 4 hours. The samples were immediately filtered through 0.2 mm filters (Supor Acrodisc, Gilman). The amount of phosphate adsorbed on calcium carbonate was calculated by subtracting the remaining phosphate concentration in the solution from the initial added phosphate concentration.

    For the adsorption equilibrium studies, 50 cm3 of solutions with different initial phosphate concentrations from 1 mM to 60 mM were combined with 0.1 gram of either calcite or aragonite in 60-cm3 capacity HDPE bottles. As discussed later all the solutions were pre-equilibrated with CaCO3 for at least 1 hr before the addition of PO4. After the pre-equilibrium the solutions were kept at a constant temperature with intermittent shaking for a given period of time. The effect of temperature (5 to 45°C) and salinity (0 to 45) on the adsorption of phosphate on calcium carbonate were studied in a similar manner. Dilution with Milli-Q water or evaporation of the Gulf Stream seawater was used to obtain solutions with salinities other than 36.4. The salinities of the samples were determined on the Practical Salinity Scale using an AutoSal Salinometer to ± 0.003.

    The effect of pH on the phosphate adsorption was examined in 50 cm3 of seawater with 4 mM of phosphate and 0.1 g of aragonite. The pH of the seawater was pre-adjusted to a value between 7.2 and 9.5 by adding HCl or NaOH to the sample. Phosphate was then added to obtain a final concentration of 4 mM. The pH was measured with an Orion 720A pH meter and a Ross combination electrode. The electrode was calibrated with a Tris seawater buffer (Millero, 1986) on the free proton scale. After the equilibration, the solutions were analyzed for the adsorbed phosphate.

  6. Findings:

    In our first series of experiments we examined the effect that the precipitation of CaCO3 can have on the adsorption of phosphate (PO4) on calcium carbonate in seawater. These measurements were made to examine the possible co-precipitation of phosphate with CaCO3. The addition of aragonite to supersaturated seawater results in a decrease in the total alkalinity (TA) of the solution due to the precipitation of CaCO3. This is demonstrated in Figure 1 for the addition of aragonite (2 g L-1) to Gulf Stream seawater (Millero et al., 1999). The TA decreases from 2468 to 1800 m mol kg-1 in the first two hours followed by a slower loss of TA to its equilibrium value (~1100 m mol kg-1 when pCO2 = 365 m atm for aragonite at equilibrium). We have made a series of measurements to see if this precipitation of CaCO3 affects the equilibrium adsorption of PO4 to aragonite.

    Figure 1. The decrease in the total alkalinity of seawater after the addition of aragonite (2 gL-1) to Gulf Stream seawater.
    Figure 1. The decrease in the total alkalinity of seawater after the addition of aragonite (2 gL-1) to Gulf Stream seawater.
    Figure 2. The adsorption of phosphate on aragonite (2 gL -1) following a one hour pre conditioning with the solid.
    Figure 2. The adsorption of phosphate on aragonite (2 gL-1) following a one hour pre conditioning with the solid.

    The adsorption of PO4 to aragonite was measured by first pre-equilibrating the aragonite with seawater for a set period of time (pre-equilibration time), followed by adding phosphate. The solutions were then equilibrated for one day. The results at different levels of phosphate are shown as a function of the pre-equilibration time in Figure 2. The addition of 40 m M of phosphate immediately after addition of the aragonite to the seawater yields a higher adsorption than those additions with pre-equilibration time longer than 60 min. The effect is nearly undetectable at levels below 30 m M of added PO4 or at pre-equilibration times over one hour. These experiments suggested that the apparent co-precipitation of PO4 with CaCO3 is diminished as the seawater approaches equilibrium with aragonite. This is partly due to PO4 retarding the precipitation of CaCO3. To avoid the possible co-precipitation of PO4 with CaCO3, we waited at least one hour before the addition of PO4 in all of our subsequent measurements.

    In our second series of measurements, we examined the adsorption and desorption of phosphate on aragonite and calcite in seawater as a function of time. The results for the adsorption of phosphate to calcite and aragonite as a function of time are shown in Figure 3. The adsorption of both aragonite and calcite appears to be a two steps process in agreement with the results of Stumm and Leckie (1970). In the first 30 minutes, the dissolved phosphate concentration decreased sharply and reached a relative constant concentration within two hours. After this initial fast adsorption period, the uptake rate of phosphate on calcium carbonate was much slower. A comparison of the adsorption of phosphate to aragonite and calcite after one and seven days is shown in Figure 4. The adsorption on aragonite is similar for the one-day and 7-day equilibrations. The adsorption on calcite increases with the longer equilibration. The differences in the extent of the adsorption on the two minerals may be related to the differences in the specific surface areas.

    Figure 3. The percentage of phosphate adsorbed on aragonite and calcite (2gL -1)in seawater at 25°C as a function of time ([PO4]0 = 10 m M).
    Figure 3. The percentage of phosphate adsorbed on aragonite and calcite (2gL-1)in seawater at 25°C as a function of time ([PO4]0 = 10 m M).

    The rates of desorption of phosphate from aragonite and calcite were also studied. These measurements were made on CaCO3 minerals that had been subjected to either a one-day or one-week pre-sorption of phosphate with constant shaking. The initial phosphate adsorbed was 10.5 and 9.9 m mol/g on aragonite and 3.5 and 4.1 m mol/g on calcite, respectively, after one-day equilibration and one-week equilibrations. Figure 5 shows that a fraction of the adsorbed phosphate was quickly released from both calcite and aragonite (20-30 min) that had been pre-equilibrated for one day. The amount of phosphate released from aragonite was the nearly the same for the samples equilibrated for one and seven days. The phosphate released from calcite, however, was lower for the sample equilibrated for seven days. The equilibrium levels of the desorbed phosphate from aragonite and calcite are nearly equal to the values found from adsorption measurements for the one or seven day equilibration.

    Figure 4. Comparison of adsorption of phosphate on aragonite and calcite (2gL -1) in seawater at 25°C at different equilibration times.
    Figure 4. Comparison of adsorption of phosphate on aragonite and calcite
    (2gL-1) in seawater at 25°C at different equilibration times.
    The percentage of phosphate remaining on aragonite and calcite (2gL -1) in seawater at 25°C after a one day and one week pre-sorption equilibrium.
    Figure 5. The percentage of phosphate remaining on aragonite and calcite (2gL-1) in seawater at 25°C after a one day and one week pre-sorption equilibrium. The initial phosphate adsorbed was 10.5 and 9.9 m mol/g on aragonite and 3.5 and 4.1 m mol/g on calcite, respectively, after one-day equilibration and one-week equilibration.

    To examine the effect of composition on the adsorption we carried out a number of measurements of the adsorption of PO4 to aragonite in seawater and the major seasalts at 25°C. Measurements made in 0.7 M NaCl (with 0.002 M NaHCO3) and seawater after one day equilibration are compared in Figure 6. Adsorption of phosphate from seawater was higher than from the 0.7 M NaCl solution. Measurements carried out in artificial seawater composed of the major components (Na+ = 0.47 M, Mg2+ = 0.053 M, Ca2+ = 0.010 M, K+ = 0.0109 M, Cl- = 0.54 M, HCO3- = 0.002 M, SO42- = 0.028 M; Millero, 1996) are identical to those obtained in natural seawater within the experimental error of the measurements (0.1 mM). To further investigate the cause of the differences we have made adsorption measurements in mixtures of NaCl (0.002 M NaHCO3) and some of the major seasalts.

    Figure 6. Comparisons of the adsorption of phosphate to aragonite in 0.7 M NaCl with 0.002 M NaHCO 3 and seawater at 25°C.
    Figure 6. Comparisons of the adsorption of phosphate to aragonite in 0.7 M NaCl with 0.002 M NaHCO3 and seawater at 25°C.

    Figure 7 shows the comparisons of the effects of different seasalts on the adsorption of phosphate using 0.7 M NaCl as reference. Adsorption of PO4 onto aragonite in NaCl (0.505 M) with Mg2+ (0.053 M) and Ca2+ (0.010 M) (as chlorides) increased the adsorption above the values for seawater. The effects of Mg2+ and Ca2+ at their seawater concentrations are similar. This means that the effect of Ca2+ is about five times greater than Mg2+. The addition of both Mg2+ and Ca2+ to NaCl increases the sorption further, but not as much as the sum of the individual effects. The addition of Na2SO4 (0.0028M) decreases the effects of both Mg2+ and Ca2+. These results suggest that the formation of MgSO4 and CaSO4 ion pairs decrease the adsorption effect of Mg2+

    Figure 7. The difference in the adsorption of phosphate to aragonite
    Figure 7. The difference in the adsorption of phosphate to aragonite in 0.7 M NaCl, NaCl + MgCl2, NaCl + CaCl2, NaCl + MgCl2 + CaCl2, NaCl + MgCl2 + Na2SO4, NaCl + CaCl2 + Na2SO4 and seawater at 25°C. The concentrations of all the components are the same as in average seawater (Millero, 1996). NaHCO3 (0.002M) was added to all the solutions except seawater.

    and Ca2+ or SO42- competes with PO4 for the adsorption sites on the carbonate minerals. The additions of some of the minor components at their seawater concentrations (K+ = 0.009M, Br- = 0.0008M, B(OH)3 = 0.0004M and F- = 0.0007M) do not affect the adsorption. Our results and those of de Kanel and Morse (1978) in seawater indicate that F- does not affect the adsorption in contrast to the dilute solution results of Stumm and Leckie (1970). The formation of MgF+ and CaF+ complexes in seawater (Millero and Schreiber, 1982) may explain this discrepancy. Measurements made in solutions of the major seawater cations and anions (Na+, Mg2+, Ca2+, Cl- and SO42-) yield adsorptions that are similar to that of natural seawater within the experimental error of the measurements (Figure 7). The effect of Mg2+ on the equilibrium adsorption of PO4 to CaCO3 agrees with the earlier studies of Stumm and Leckie (1970) and de Kanel and Morse (1978).

    Figure 8. The effect of temperature on the adsorption of phosphate on aragonite in seawater (S = 36).
    Figure 8. The effect of temperature on the adsorption of phosphate on aragonite in seawater (S = 36).
    Figure 9. The effect of salinity on the adsorption of phosphate on aragonite in seawater at 35°C.
    Figure 9. The effect of salinity on the adsorption of phosphate on aragonite in seawater at 35°C.

    The effect of temperature (5 to 45°C) on the adsorption of phosphate on aragonite was examined in seawater as a function of salinity (S = 0, 15 and 36). The results at S = 36 are shown in Figure 8. The adsorption increases with increasing temperature for all of the salinities studied. The effect of salinity (S = 0 to 41) on the adsorption of phosphate on aragonite was also examined at different temperatures. The results at 35°C are shown in Figure 9. The adsorption increases as the salinity decreases. A decrease in the salinity from 41 to 36 results in a 9% increase in the observed phosphate adsorption. The effect is much larger at the lower salinities. For example, the adsorption increased by 29% when the salinity was lowered from 5 to 0.

    To elucidate the effects of salinity on the adsorption of phosphate, we examined the effect of composition on the adsorption. Measurements in pure water with added HCO3- serving as a buffer showed that the adsorption was strongly affected by the amount of HCO3- added to the solution (see Figure 10a). Measurements made at the same HCO3- concentration or carbonate alkalinity (Figure 10b) give an adsorption that is nearly independent of ionic strength or salinity.

    The effect of pH (7.4 to 9) on the adsorption of phosphate to aragonite was examined in seawater (S = 36) at 25°C. The results given in Figure 11 show that the adsorption increases to a maximum at a pH of 8.6. The effect of pH on the adsorption can be related to changes in the surface sites or changes in the speciation of phosphate in the solution. At a pH near 8.6 the fractions of HPO42- and PO43- are nearly equal (Millero, 1996) and neither species go through a maximum near this pH. This leads us to believe that the effect of pH on the adsorption is not related to the speciation of PO4 in solution, but to changes in the surface speciation.

    Figure 10. The effect of bicarbonate on the adsorption of phosphate on aragonite at 25°C.
    Figure 10. The effect of bicarbonate on the adsorption of phosphate on aragonite at 25°C.
    Figure 11. The effect of pH on the adsorption of phosphate on aragonite at 25°C.
    Figure 11. The effect of pH on the adsorption of phosphate on aragonite at 25°C.

    Distribution coefficients for phosphorus partitioning between sediment/seawater in Florida Bay have been experimentally determined. The partitioning of any elements between water and sediment is usually quantified by the distribution coefficient Kd. Distribution coefficient Kd of phosphorus is defined as Kd = Cs/Cw where Cs is the concentration of phosphorus on particle surface and Cw is the concentration of phosphorus in seawater. Kd is a key parameter that governs phosphorus partitioning between seawater and particle surface. To estimate the distribution coefficients for phosphorus partitioning between suspended sediments and seawater, surface sediments have been collected from Florida Bay at several locations with typical environmental conditions.

    Figure 12. The station locations of sediment samples collected in Florida Bay
    Figure 12. The station locations of sediment samples collected in Florida Bay

    The sediments were equilibrated with low nutrient seawater at a constant temperature. A 0.5 g of sediment sample were mixed with 50 ml phosphate-free seawater in a plastic bottle and placed on a shaker for continuous shaking. These interaction conditions simulate the resuspension of sediment into water column. Phosphate in sediment were released to seawater while desorption takes place. After 16 hours, the particles will then be separated from seawater by filtration. The phosphate concentrations in equilibrated seawater were determined by a spectrophotometric method using an autoanalyzer. The sediments will then be equilibrated with a 50 ml MgCl2 solution (1 M) at pH of 8 for 4 hours. The sorbed and desorbable phosphorus were extracted into the solution by a complexing reaction with MgCl2. The extracted phosphate in MgCl2 solution were determined after separation from the suspended sediments. Our preliminarily results with limited sediment samples showed a linear correlation between phosphate concentrations in seawater and exchangeable phosphate on the sediment surface from Florida Bay (see Figure 12 for station locations). It is interesting to note that samples taken from Rankin Bight showed the highest phosphate content in both P released to seawater and exchangeable P in sediment (Figure 13). It has been shown that blue green alga Synechococcus sp. blooms are frequently occurring in Rankin Bight (Philips et al., 1999). Sediments in Rankin Bight are anoxic with evidence of sulfate reduction and denitrification as evidenced by the maximum in nitrite concentration.

    Figure 13. Kd estimated based on the partitioning of phosphate between seawater and Florida Bay sediment.
    Figure 13. Kd estimated based on the partitioning of phosphate between seawater and Florida Bay sediment.

    Preliminary estimated Kd is in the order of 0.1 L/g. Since Florida Bay is subdivided by mud banks into partially-isolated basins, spatial variation in sediment characteristics is expected due to difference in environmental conditions. We proposed to collect surface sediment samples from the all sub-basins in Florida Bay and conduct similar experiments to verify any spatial variation of Kd in Florida Bay. We also proposed to study the effect of salinity and temperature on the Kd. With such a systematic study, a quantitative relationship between Kd and temperature and salinity can be established. A fitted equation of Kd as a function of salinity and temperature can be used in a water quality model to predict the fate of input phosphorus in Florida Bay.

  7. Evaluation:

    The main goals of this project have been achieved during the project period. These studies provide equations that can be used to model the adsorption of phosphate on CaCO3 sediments in tropical waters. Our kinetic measurements show that the rates of adsorption and desorption are quite fast (30 min). Up to 80% of the adsorbed phosphate is released from calcium carbonate over one day. The amount of PO4 left on the CaCO3 is close to the equilibrium adsorption. For shallow waters like Florida Bay, the resuspension of the sediments may provide a short-term source of phosphate needed by plants, and represent a sink for excess amounts of phosphate from land sources. The effect of CaCO3 precipitation on the adsorption of PO4 to CaCO3 was found to be unimportant except at high concentrations of PO4. Measurements made after pre-equilibrating the added CaCO3 with seawater for an hour were not affected by the co-precipitation of phosphate with CaCO3. The solution composition affects the adsorption of phosphate on calcium carbonate. The effect of salinity on the adsorption is largely related to the dilution of the concentration of HCO3- in the waters. This may be important in estuaries where the alkalinity of the rivers may be high. Measurements made at the same level of HCO3- are nearly independent of salinity. The adsorption of PO4 to CaCO3 in NaCl solutions was less than in seawater at the same ionic strength and level of HCO3-. The higher values in seawater were due to Mg2+ and Ca2+ ions. The addition of Mg2+ and Ca2+ to NaCl increases the adsorption whereas the addition of SO42- decreases the adsorption. The effects of Ca2+ and Mg2+ are diminished with the addition of SO42- apparently due to the formation of MgSO4 and CaSO4 ion pairs and/or the adsorption of SO42- on the surface of CaCO3. Adsorption of PO4 in solutions with the major seasalts (Na+, Mg2+, Ca2+, Cl-, HCO3- and SO42-) was in good agreement with the values in seawater.

    We have taken full advantage of internet to disseminate our project results. The results of our project from both field and laboratory studies have been posted on our/ website since the day the project started. Anyone interested the Florida Bay area that needs to know the parameters we investigated can access them through our website. The nutrient data from the bimonthly survey have been compiled in the SFERPM database. Results of this study were presented in 1998 and 1999 at the Florida Bay Science Conferences. Two papers resulting from this project have been submitted, one to Bulletin of Marine Science (in press) and another to Aquatic Geochem.

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    Fourqueran J. W., Zieman J. C. and Powell G. V. N. (1992) Phosphorus limitation of primary production in Florida Bay: Evidence from C:N:P ratios of the dominant seagrass Thalassiatestudinum, Limnol. Oceanogr. 37, 162-171.

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