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ENVIRONMENTAL LETTERS, 60), 167-174 (1974)

SEDIMENTARY FLUORITE IN TAMPA BAY, FLORIDA

William H. Taft and Dean F. Martin

Department of Geology, Department of Chemistry
University of South Florida
Tampa, Florida 33620

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INTRODUCTION

Florida is a major source of phosphate, and the processing of this resource has undergone a number of changes as environmental degradation was discovered and new technology became available to-resolve those problems. One such problem was the amount of fluorine gas emitted as a waste product through stacks. Complaints by local residents, cattlemen, orange growers, and commercial florists encouraged one such company, Gardinier, Inc. (then owned by US Phosphoric), in approximately 1963, to install scrubbers. The collected fluorine was then discharged into Tampa Bay. Although the amount of fluoride entering the Bay has varied depending on the amount of phosphate produced, in July of 1973 the daily discharge amounted to approximately 24,000 pounds., In addition to fluoride, the effluent contained approximately 27,000 pounds of phosphate and 3,000 pounds of nitrate per day. Gardinier has been warned about such excessive discharge, and through expenditures of approximately $12,000,000 in anti-pollution equipment, hopes to reduce the amount of these pollutants by 97 percent in December, 1973.

EXPERIMENTAL

Physical Observations

Two canals are used for discharging Gardinier's phosphate processing wastes into the Bay. At the mouth of the southernmost canal there is a deltaic deposit of fluorite that, to the best of our knowledge, is the only known deposit of sedimentary fluorite in the world. In cross-section, the deposit is approximately three inches thick at the point of initial discharge and thins rapidly to translucent flakes at the outer edges of the deposit - approximately 1,000 feet into the Bay. The fluorite deposit is composed of alternating layers of fluorite and loose grains except where sticks, limbs and roots of trees have provided a site, free of detrital grains, for fluorite to precipitate.

When standing in relatively shallow water (approximately 0.6 meters), overlying the fluorite deposit, there is usually a noticeable increase in temperature within ten to fifteen centimeters above the sediment-water interface (Table 1). To test the significance of this increased temperature near the base of the water column, bamboo poles sufficiently long to stick out above sea level were pushed into the sediment. After thirty days, there was a coating of fluorite on the poles, but the fluorite was restricted to the lowest ten to fifteen centimeters of the water column - the zone of highest water temperatures which owes its origin to the thermal discharge from the phosphate processing plant. In this relatively shallow water, a great deal of mixing occurs, yet a portion of the thermal discharge is capable of maintaining some of its composition as shown by the physical and chemical differences between the tops and bottoms of measured water columns (Table 1).

Given the quantities of fluoride that are discharged daily, and have been since 1963, we can only account for a small percentage as being incorporated in the fluorite deposit. Clearly, a large quantity of fluoride escapes local precipitation as shown by the fluorine concentrations in Bay water at the most remote station such as 8.

As might have been anticipated by the extremely low pHmeter readings, there were no visible signs of plant or animal life within the area of fluorite deposition. During the time of field study, Tampa Bay had an abundance of living jellyfish. In the study area, however, there were sitings of four to six such organisms, but they were all dead. Presumably, the jellyfish had been killed by the chemical discharge.

Sampling and Analytical Procedures

Water and sediment samples were collected at eight sites in Tampa Bay adjacent to the southernmost discharge canal of Gardinier. Water samples were frozen within four hours of collection and sediment samples were dried and prepared for X-ray diffraction.

Temperature and pH data were obtained in the field at the time of sampling. Temperatures were obtained using a mercury thermometer. Because the temperature increased rapidly during equilibration while water temperature was being read at the surface during a 30-sec period, the temperatures of bottom samples in Table 1 should be thought of as minimum values. The pH values were obtained with a.Beckman Model G pH meter; pH values obtained in the field were checked in the laboratory before analysis and were general to-within 0.1 unit of the field value. Salinity was determined by the Harvey method. 1 Silicate and phosphate analyses were obtained by means of a Technicon Autoanalyzer using standard procedures described in the EPA Methods Manual. 2

Fluoride was determined by using an Orion fluoride electrode (model 94-09A) and a single junction reference electrode (model 90-01) with a Corning Model 10 Expanded pH meter. A procedure described earlier 1 was modified. A 10-ml sample of water was added to 40 ml of stock sea water (salinity 30'/oo) in a plastic beaker, then mixed with 10 ml of total ionic strength buffer (TSIB, pH = 5.5, ionic strength = 1.9). The reading was recorded after 5 minutes of stirring (=E i ). Readings (EX) were also recorded after addition of 0.1, 0.2, 0.3, 0.4, and 0.5 ml of stock fluoride solution (100 PPM F in distilled water). The response, log E x as a function of added.fluoride (in PPM), was linear and the value of apparent initial fluoride (F-)i was obtained by extrapolation knowing E i. The original apparent fluoride concentration (F-)0 was obtained from the relationship

where (F-)s is the apparent fluoride concentration in the stock sea water (a value obtained using 10 ml of distilled water instead of sample), and where 0.2 is the dilution factor. The value of (F-)s was 1.05 PPM, which would correspond to a value of 1.3 PPM for the sample of salinity of 36*/oo (the typical value is 1.28 for salinity 35*/oo). 3 Precision was estimated for samples 6S and 6B.and the mean and standard deviations were 34.210.8 and 35.0 + 0.3, respectively. No significant variation in fluoride values were observed over a 24-hour period after refrigeration.

DISCUSSION

A number of unique features were observed during this study. The most obvious is the existence of the sedimentary fluorite. The second is the remarkably low pH meter readings which indicate the buffering capacity (6,000 milliequivalents per kilogram of sea water) 4 of twelve (12) acre-meters of estuarine water was virtually exhausted. Third, the fluoride concentration is as much as forty times its concentration in normal sea water, undoubtedly because of the complexing by hydrogen ions and silicon. Fourth, the temperature differentials in such relatively shallow water (0.6 meters) are not only significant, but are the inverse of what would normally be expected, i.e., the water temperature is warmer at the base of the water column than at the air water interface.

Although in some instances the solution to pollution may be dilution, in this case, the fluorine problem in stack gasses was translated to water and thermal pollution and precipitation problems in Tampa Bay. When the anti-pollution controls come on line in December, 1973, this particular pollution problem should be sharply curtailed. Meanwhile, the values of pH, fluoride and phosphate concentrations that were obtained in August, 1973 should provide a benchmark to establish the extent of progress in the future.

ACKNOWLEDGEMENTS

We are grateful for the technical information and assistance provided by Mr. Donald Graff, Manager of Gardinier, Inc.; officer Robert Knight of the Florida Marine Patrol; Ms. Val Maynard, USF Department of Marine Science, who performed the analysis of silicate and phosphate; and Mr. James Goetz, who performed the X-ray analyses. One of us (DFM) is grateful for the support provided by a Public Health Service Research Career Development Award (K04-GM 42569-05, National Institute of General Medical Sciences).

REFERENCES

1. D. F. Martin, Marine Chemistry: Analytical Methods, 2nd ed., Dekker, New York, 1973, p. 85, p. 305.

2. Environmental Protection Agency, EPA Methods for Chemical Analysis of Water and Wastes, U. S. Department of the Interior, Washington, 1971.

3. J. P. Riley and R. Chester, Introduction to Marine Chemistry Academic Press, New York, 1971, p. 81.

4. L. G. Sillen, Science, 156, 1189 (1967).

To learn more about fluoride pollution from the phosphate industry, see www.fluoridealert.org/phosphate/overview.htm