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MAJOR AND MINOR ION GEOCHEMISTRY OF GROUNDWATERS FROM BERMUDA

JAMES ANDRE KENT SIMMONS University of New Hampshire, Durham

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Major and minor ion geochemistry of groundwaters from B erm uda

Simmons, James Andre Kent, Ph.D.

University of New Hampshire, 1987

Copyright ©1987 by Simmons, James Andre Kent. All rights reserved.

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University Microfilms International MAJOR AND MINOR ION GEOCHEMISTRY OF GROUNDWATERS FROM BERMUDA

BY

J.A. KENT SIMMONS B.A., University of Rhode Island, 1981 M.S., University of New Hampshire, 1983

DISSERTATION

Submitted to the University of New Hampshire in Partial Fulfillment of the Requirements for the Degree of

Doctor of Philosophy

in

Earth Sciences

May, 1987 ALL RIGHTS RESERVED

c 1987

J.A. Kent Simoons This dissertation has been examined and approved.

Dissertation directo/, Wm/Berry Lyons Associate Professorrof Earth Sciences P. ~~£UL— Theodore C. Loder Associate Professor of Earth Sciences

fA./e U J j . Francis R. Hall Professor of Hydrology

? * & J . £ ^ Mark E. Hines Researcffftssistant Professor c.f Earth Sciences

Bermuda Bioyogical Station for Research

II, 1197 itV< [j DEDICATION

To Dad, Hon and B.J., whose love, encouragement, and personal sacrifice have made everything worthwhile.

iv ACKNOWLEDGEMENTS

The funds which supported this research were graciously

provided by the Bermuda government through the Department of Public

Wbrks. These funds were provided for hydrogeological and hydro-

geochemical studies of the Pembroke marsh landfill. I wish to

express my sincere thanks to the Bermuda government for selecting me

for this project. Special thanks to the members of the Bermuda

government's Pembroke Dump Committee: Hon. Robert Birritt, Chair­

man; Mr. David Summers; Mr. John Roach and Mr. Geofrey Melotti; and

also to personnel at the Public Works Department: Mr. James Thomson

and Mr. Sandy Taylor.

This research was conducted at the Bermuda Biological Station

(BBS), St. George, Bermuda. The BBS has provided major contribu­

tions in finances and facilities for all of my work in the past five

years, and particularly in the past three years with this project,

and for this I am very grateful. I am also very grateful and highly

indebted to a good friend and advisor Dr. Anthony H. Knap, who has

always taken a keen interest in my work at the Bio-Station. I wish

to thank Dr. Timothy Jickells for his consultation on analytical

techniques and also for many long and involved discussions about groundwater and related geochemical problems. Thanks also goes to

Kandy Binkley, who advised in the collection and analyses of organic

samples. I also wish to thank all resident scientists and research

support staff like Rachael and Sue J., and also Biostation office

v staff like Margaret, Jill, Harry, Nancy, and everyone in accounts, public relations, and development, and also Butch, Sonny, Austin and all the guys in operations and maintenance.

I would like to thank all of the professors and staff (Linda,

Sue and Caroline) in the Earth Sciences Department at the University of New Hampshire, in particular the members of my doctoral committee, they include: Dr. Wto. Berry Lyons, Chairman; Dr.

Francis R. Hall; Dr. Theodore C. Loder; Dr. Mark E. Hines; Dr.

Anthony H. Knap; and also unofficial advisors Dr. Henri Gaudette and

Dr. Mary Jo Spencer, each of whom shared considerable amounts of time and experience with me over the past years. I wish to express my extra special thanks to Berry, not only for his major contribution to my educational aspirations over the past six years, but also for being a fantastic advisor and a truly good friend, a friendship which I know will extend many, many years into the future. Also, I wish to thank Berry's wife, Anne, who sacrificed a great deal of time to share with me some of her computer skills during a particularly busy period of finishing her own thesis.

Lastly, I wish to thank all of my friends and colleagues in both New Hampshire and Bermuda. I especially wish to thank my very good friend Mimi Boxwell, who helped keep things in perspective during my last grueling year at UNH.

I thank you all.

Graduate Committee

Dr. wm. Berry'Lyons Dr. Theodore C. Loder Dr. Francis R. Hall Dr. Mark E. Hines Dr. Anthony H. Knap

vi Research sponsorship was generously provided by;

Bermuda Biological Station for Research, Inc. Bermuda Government Department of Public works Bermuda Junior Service League American Friends of Bermuda

Pembroke Dump Committee

Hon. Robert Barritt Mr. John Roach Mr. David Summers Mr. Geoffrey Melotti

Biostation Scientists and Staff

Dr. Wolfgang E. Sterrer Harry Barnes Dr. Anthony H. Knap Nancy Smythe Dr. Thomas D. Sleeter Phillip Reeves Dr. Clayton B. Cook Patricia Brown Dr. Susan B. Cook Zina Francis Dr. Thomas M. Iliffe Christine Harrell Tim D. Jickells Patricia Little Susan Jickells Mary van Soren Kandy Binkley Gerraldine Pharris Rachael Sherriff-Dow S. Robertson Smith Wilfred "Butch" Stovell Sheila Wyers Chesley "Sonny" Foggo Susan Boyd Austin Smith Jon Moreland Winston "Pene" Foggo Barbara Sheehan Frederick "Teddy" Sewell Warren Smith Margaret Emmott Michael Rhodes Jill Cadwallader Anthony She r ri ff Virginia O'Connor Public Works

James Thomson Patricia Hollis James Conyers Sandy Taylor

University of New Hampshire Faculty and Staff

Dr. Herbert Tischler Dr. James D. Irish Dr. Henri E. Gaudette Dr. Wm. Berry Lyons Dr. Francis R. Hall Dr. Jo Laird Dr. Cecil J. Schneer Dr. William Olszewski Dr. Robert I. Davis Dr. Mary Jo Spencer Dr. Lincoln R. Page Dr. Neil R. Pettigrew Dr. Franz E. Anderson Dr. Mark E. Hines Dr. Theodore C. Loder Dr. David A. Gust Dr. Francis S. Birch Linda Hoadley Dr. Wendell S. Brown Susan Gagnon Dr. Wallace A. Bothner Caroline Miller Dr. S. Lawrence Dingman Susan Lucus Dr. Paul A. Mayewski Frank Smith

vi i T

TABLE OF CONTENTS

DEDICATION 1 v

ACKNOWLEDGEMENTS ...... v

LIST OF TABLES......

LIST OF F I G U R E S xi 1

ABSTRACT...... *v

CHAPTER PAGE

I. INTRODUCTION...... 1

II. BACKGROUND ...... 4

Geology of Bermuda...... 4

Li me stones...... 4

S o i l s ...... 9

Carbonate Diagenesis ...... 11

Ghyben-Herzberg L e n s e s ...... 14

Redox R e a c t i o n s ...... 19

Geochemistry of Bermuda Groundwaters...... 25

Study A r e a ...... 31

III. METHODS...... 35

Sample Col lec ti on...... 35

Cleaning...... 39

Samplers...... 39

Iron and Manganese...... 40

Nu t r i e n t s ...... 41

Blanks...... 41

v i i i Field Methods ...... 42

Analytical Methods ...... 43

Sample Preservation...... 43

pH and Alkalinity ...... 43

O x y g e n ...... 44

Phosphate, Nitrate, Silicate and Aranonia . . . 44

S u l f i d e ...... 44

Fluoride, Nitrate, Sulfate and Bromide .... 44

Chloride...... 45

Calcium, Magnesium, Strontium, Sodium and Po t a s s i u m ...... 45

Iron and Manganese...... 45

Conductivity...... 46

IV. MAJOR ICN GEOCHEMISTRY IN IWO BERMUDA MARSHES...... 49

Pembroke Rest Home (P R H ) ...... 60

Jubilee Road ( J R ) ...... 74

V. ALKALINITY...... 116

Model Development...... 130

Assumptions...... 138

E q u a t i o n s ...... 138

R e s u l t s ...... 145

Saturation...... 150

Aquifer Geochemistry...... 165

VI. THE FATE 07 INORGANIC NITROGEN AND PHOSPHORUS IN BERMUDA GROUNDWATER...... 169

Inorganic Phosphorus...... 169

Inorganic Nitrogen...... 173

VII. CONCLUSIONS...... 195

ix VIII. FUTURE W O R K ...... 199

APPENDIX I. EQUILIBRIUM CALCULATIONS...... 200

APPENDIX II. X-RAY DIFFRACTION ANALYSES...... 215

APPENDIX III. COMPUTERIZED ALKALINITY MOD EL ...... 225

APPENDIX IV. D A T A ...... 229

LIST OF REFERENCES...... 235

X LIST OF TABLES

TABLE PAGE

2.1. Bermuda stratigraphic c o l u m n ...... 6

2.2. Mineralogy of Bermuda s o i l s ...... 10

2.3. Biogeochemical reactions...... 23

2.4. Composition of s e a w a t e r ...... 26

2.5. Composition of Bermuda r a i n ...... 27

3.1. Summary of methods...... 47

3.2. Precision and detection li mi ts...... 48

4.1. Composition of fresh groundwater...... 64

4.2. Major ion speciation...... 97

5.1. Alkalinity m o d e l ...... 146

5.2. Saturation indices for calcite and aragonite...... 153

5.3. Carbonate equilibria...... 154

5.4. pH, alkalinity and calculated PC02 ...... 155

5.5. Composition of limestone...... 166

6.1. Speciation of orthophosphate...... 170

6.2. Nitrogen transfer r e a c t i o n s ...... 175

5.3. Saturation indices for apatite...... 186

6.4. Nutrient box-model d a t a ...... 188

6.5. Nutrient flux r e s u l t s ...... 190

6 .6 . Sewage outfall results...... 194

xi LIST OF FIGURES

Figure 2 .1 . weight percent of carbonate minerals . 7

Figure 2 .2 . Geologic map of Bermuda ...... 8

Figure 2.3. Ghyben-Herzberg hydrostatic model . . . 15

Figure 2.4. Ghyben-Herzberg realistic model .... 17

Figure 2.5. Hydrogeologic cross-section of the Devonshire lense ...... 20

Figure 2 .6 . Map of freshwater nuclei ...... 21

Figure 2.7. Locations of sampling wells ...... 32

Figure 3.1a. Water sampler. Bailer ...... 36

Figure 3.1b. Water sampler. Packer ...... 37

Figure 3.2. Relative salinity profiles ...... 38

Figure 4.1. Pie chart composition of r a i n ...... 51

Figure 4.2. Pie chart composition of seawater . . . 52

Figure 4.3. Pie chart groundwater composition at PRH 53

Figure 4.4. Pie chart groundwater composition at JR surface . . 54

Figure 4.5. Pie chart groundwater composition at JR, 150 cm . . 55

Figure 4.6. Pie chart groundwater composition at JR, 300 cm , . 56

Figure 4.7. Pie chart groundwater composition at JR, 460 cm . . 57

Figure 4.8. Pie chart groundwater composition at JR, 600 cm . . 58

Figure 4.9. Chloride versus depth at P R H ...... 61

Figure 4.10. Sodium versus depth at P R H ...... 66

Figure 4.11. Potassium versus depth at P R H ...... 67

Figure 4.12. Calcium versus depth at P R H ...... 68

Figure 4.13. Magnesium versus depth at P R H ...... 69

xii Figure 4.14. Strontium versus depth at P R H ...... 70

Figure 4.15. Sulfate versus depth at P F H ...... 71

Figure 4.16. Bromide versus depth at P R H ...... 72

Figure 4.17. Fluoride versus depth at P R H ...... 73

Figure 4.18. Chloride versus depth at J R ...... 75

Figure 4.19. Sodium versus depth at J R ...... 79

Figure 4.20. Potassium versus depth at JR ...... 81

Figure 4.21. Calcium versus depth at J R ...... 83

Figure 4.22. Magnesium versus depth at J R ...... 85

Figure 4.23. Strontium versus depth at J R ...... 87

Figure 4.24. Sulfate versus depth at J R ...... 89

Figure 4.25. Bromide versus depth at J R ...... 91

Figure 4.26. Fluoride versus depth at J R ...... 93

Figure 4.27. Ideal mixing c u r v e ...... 96

Figure 4.28. Sodium versus chloride ...... 99

Figure 4.29. Potassium versus chloride ...... 101

Figure 4.30. Silicate versus depth ...... 103

Figure 4.31. Calcium versus chloride ...... 105

Figure 4.32. Strontium versus chloride ...... 108

Figure 4.33. Magnesium versus chloride ...... H O

Figure 4.34. Sulfate versus chloride ...... H I

Figure 4.35. Bromide versus chloride...... 112

Figure 4.36. Fluoride versus chloride ...... 114

Figure 5.1. Eh versus d e p t h ...... H 9

Figure 5.2. Oxygen versus d e p t h ...... 121

Figure 5.3. Nitrate versus d e p t h ...... 123

Figure 5.4. Ammonium versus d e p t h ...... 126

x i l i Figure 5.5. Nitrite versus depth ...... 128

Figure 5.6. Iron and manganese versus de p t h ...... 131

Figure 5.7. Sulfide versus depth ...... 135

Figure 5.8. Graphical representation of weak acid s p e c i a t i o n ...... 142

Figure 5.9. pH, alkalinity and PCO, versus depth at P R H ...... 7 ...... 156

Figure 5.10. pH, alkalinity and PCX)- versus depth at J R ...... 7 ...... 159

Figure 5.11. ^1, alkalinity and PCO~ versus depth at J R ...... 7 ...... 162

Figure 6.1. Graphical representation of orthophosphate s p e c i a t i o n ...... 171

Figure 6.2. Nitrogen cycle for Bermuda groundwater system . . . 174

Figure 6.3. Phosphate versus depth ...... 178

Figure 6.4. Phosphate versus d e p t h ...... 179

Figure 6.5. Phosphate versus chloride ...... 180

Figure 6 .6 . N/P ratios versus depth ...... 182

Figure 6.7. Box model representing fixed nitrogen fluxes . . . 191

Figure 6 .8 . Box model representing fixed nitrogen fluxes . . . 192

xi v ABSTRACT

MAJOR AND MINOR ION GEOCHEMISTRY OF GROUNDWATERS FROM BERMUDA

by

J.A. Kent Simoons University of New Hampshire, May, 1987

Groundwaters near two Bermuda marshes were sampled and studied

for variations in major and minor ion geochemistry. Factors which

influence groundwater geochemistry include Ghyben-Herzberg/

freshwater-seawater mixing, carbonate dissolution and recrystal­

lization, redox and the hydrology. The major ions calcium, magnesium and strontium have been previously studied by Plummer et

al. (1976) near these two marshes and across the Devonshire lens.

This study focuses on two wells and the chemical variations which occur with depth down the groundwater column.

Samples were collected at specific depths from two wells, one near the Pembroke Rest Home (PRH) adjacent to the Pembroke marsh, and the other along Jubilee Road (JR) near the Devonshire marsh.

Collections were performed by bailing water from the borehole with a

thief-type sampler and also from the aquifer using an inflatible packer. Analyses were performed to determine sodium (Na), potassium

(K), calcium (Ca), magnesium (Mg), strontium (Sr), chloride (Cl),

sulfate

(N03~), nitrite (N02~), ammonium (NH4+ ), phosphate (PO^-), iron

xv (Fe), manganese (Mn) and silicate (H4SiO^). With these data, the groundwater could be traced from its rainfall origin through carbonate reactions, biogeochemical reactions and mixing to its discharge to the inshore waters.

Die mixing of fresh water and sea water shows conservative mixing for sodium and bromide, while the other major ions are all influenced by non-conservative input and removal mechanisms. An alkalinity model has been developed, and the contributions to the alkalinity from carbonate dissolution and redox have been quantified. Results suggest that carbon dioxide generation from aerobic respiration and nitrate reduction are probably the dominant reactions influencing the carbonate geochemistry, while iron, manganese and sulfate reduction, although clearly detectible, play only a minor role. Die geochemistry of phosphate and fixed nitrogen has also been considered. Phosphate uptake from the groundwater does not appear to result from the formation of the mineral apatite.

However, evidence does point toward PO^3- removal by calcium- carbonate phosphate surface adsorption. Flux calculations were performed to determine groundwater fixed nitrogen and phosphorus discharge to the Bermuda inshore waters. Die results indicate that groundwaters do contribute significantly to the inshore water nutrient supply. CHAPTER I

INTRODUCTION

Since the initial groundwater studies of Vacher and Plummer in the mid 1970s (Vacher, 1973; 1973; 1978a; 1978b; Plummer, 1976), many advances have been made in both research and development of

Bermuda's groundwater resources, and today groundwater provides a significant potable water supplement to the rainfall catchment and cistern systems. However, along with the increase of groundwater usage has come the recognition that various human-induced contamination problems threaten future development of this vital resource. Itie most comnon anthropogenic activities which create poor water quality conditions of Bermuda groundwater include:

1 ) septic systems and cesspools; 2) landfills; 3) buried product storage tanks and pipelines (e.g. gasoline); 4) agricultural activities; 5) seawater intrusion; 6) brine injections; 7) accidental spills; 8 ) land sludge disposal.

In recent years, much of the groundwater research on the island has focused on several of these problems. Septic system and cesspool contamination has been addressed by Simmons (1983), Simmons et al. (1985), Jickells et al. (1986) and Thomson and Foster (1986), and landfills and petroleum in buried tanks by Simmons (1985)

(Lindorff, 1979). From these investigations the magnitude of these problems are slowly being understood. 2

This study attempts to define some of the fundamental geochemical and hydrological processes occurring in the Bermuda groundwater system. These include the geochemistry of the major and minor ions in the system, their source to the groundwater, and their ultimate fate in the environment. A clear understanding of these elements will provide a good basis for future investigations concerning the processes and activities that might perturb the

"natural" system.

The major ions considered in this study include sodium, potassium, calcium, magnesium, strontium, chloride, sulfate, fluoride, bromide and alkalinity. The definition of "major ion" has been obtained from the literature of seawater chemistry. That is, they occur as dissolved species in open ocean sea water in concen­ trations greater than 1 ppm (mgAg) (Stowe, 1979; Gross, 1982). For

Bermuda, an oceanic island, this comparison of groundwater and also rain water to sea water is quite appropriate since these ground­ waters and rainwaters are to a large extent derivatives of seawater.

The chemistry of open ocean sea water has been studied in great detail since the voyage of the H.M.S. Challenger (1872 to

1876). Mean sea water has a salinity of ”34.9°/oo. The ratio of each of the major ions to one another is constant everywhere, even in areas where salinity may experience minor fluctuations due to evaporation and precipitation (Table 2.4; Broecker and Peng, 1982;

Stumm and Morgan, 1981). For Bermuda groundwaters, which have been shown to be a mixture of fresh water and sea water (Vacher, 1974;

Plummer, 1976), deviations of the major ion ratios from the 3

constancy of composition indicates the presence of input or removal

mechanisms and also the magnitude of these mechanisms.

Minor ions occur in seawater at concentrations less than 1

ppm. Nitrogen and phosphorus and their biological and pH-dependent

species— nitrate, nitrite, ammonia, ammonium, biphosphate, di­

hydrogen phosphate and phosphate— have been considered in this

study, along with hydrogen sulfide, iron and manganese. Nitrogen

and phosphorus compounds are biological nutrients. Due to uptake

and release by organisms, their concentrations in natural waters can

span a large range (ppb to ppm) and do not follow the same

compositional rules as do the major ions. The dissolved nitrogen

species, iron, manganese and hydrogen sulfide are important in

microbial redox reactions where organic matter is oxidized. The

speciation of these elements is strongly dependent upon the amount

of organic matter and dissolved oxygen in the system and the rate of

organic matter decomposition.

This study will consider freshwater-seawater mixing models,

calculate theoretical concentrations of major and minor ions and

attempt to address their input to or removal from the groundwater

system. Particular attention will be focused on the parameters which comprise alkalinity, and implications are made as to the

contribution of the various ions in groundwater to the inshore marine waters of Bermuda. CHAPTER II

BACKGROUND

Geology of Bermuda

Limestones

The limestones and paleosols of Bermuda have been discussed in detail by previous investigators (Verrill, 1907; Sayles, 1931;

Bretz, 1960; Ruhe et al., 1961; Mackenzie, 1964; land, 1967, 1970;

Blackburn and Taylor, 1969; Bricker and Mackenzie, 1970; Vacher,

1973; Harmon et al., 1978). It appears that three tholeiitic basalt pedestals— the Argus or Plantagenet Bank, Challenger Bank and

Bermuda Bank have originated in volcanism at the mid-Atlantic Ridge

95-110 my BP with seafloor spreading accounting for their present position at 32°N latitude and 65°H longitude, 2,000 km west of the active ridge system (Gees, 1969; Morris et al., 1977). The combination of shallow water deposition of marine carbonates together with eolian dune building have been shown to be responsible for the present configuration of the Bermuda Islands (Bretz, 1960;

Land et al., 1967; Mackenzie, 1964; vacher, 1973; Harmon, 1978). 2 Today, 54 km of limestones and interbedded paleosols are exposed above sea level. However, with sea level fluctuations corresponding to glacial and interglacial periods, at times the island's area has 2 been greater than and less than the 54 km of today (Bretz, 1960;

Land et al., 1967; Hannon et al., 1978). 5

The cotnplex sedimentary geology of Bermuda limestones can be attributed to the following factors: 1 ) sea level fluctuations, 2 ) production of carbonate sediments by organisms, 3) eolian processes,

4) soil formation, and 5) destruction of morphological features by chemical action (Bretz, 1960; Ruhe et al., 1961; Land et al., 1967;

Vacher, 1973). The first in-depth descriptions of sedimentological stratigraphy were attempted by Verrill (1907) and Sayles (1931) although previous authors such as Stevenson (1897) and Heilprin

(1889) had also made considerable contributions. The stratigraphic sequence described by Sayles (1931) has not been fully accepted by the more recent investigators (Bretz, 1960; Land et al., 1967 and

Vacher, 1973), although much of Sayles' basic concepts have remained intact. The most recent revision is that of Vacher (1973) (Table

2.1). Vacher's description has significantly simplified previous work in that all limestones younger than the Belmont formation have been redefined as Paget limestone consisting of an upper and lower member (Fig. 2.1; Vacher, 1973).

The diagenetic history of the three primary Bermuda carbonate units— the Paget, the Belmont and the Walsingham— follows pre­ dictable pathways. The younger Paget formation contains a high weight percent of metastable aragonite and high Mg-calclte whereas the older Belmont and Walsingham formations are depleted in these minerals with the more stable low-Mg calcite being dominant (Fig.

2.2; Ristvet, 1971; Plummer et al., 1976). (Carbonate diagenesis is discussed in more detail in a following section.) Perhaps the most impressive features of these limestone units are their difference in karst morphology. As a result of approximately one million years of 6

Table 2.1. Bermuda stratigraphic column, reproduced from Vacher (1973).

Land et al. (1967) Vacher (1973)

Southampton Fm. (eolian) P Upper Paget member A G E St. George's Soil T intra-Paget soil

Spencer's Point Fm. (eolian) F 0 Pembroke Pm. (eolian) R Lower Paget member M Harrington Pm. (soil) A T Devonshire Fm. (marine) 1 O Shore Hills Soil N sub-Paget soil

Belmont Pm. (marine)

Soil

Walsingham Pm. (eolian)

1976). Tooo al.»

} 500

1

120 7 GEO. AGE BP X103yrs 60 \ Si \ 60 30

X function of geologic age (reproduced from e t Plummer Figure 2.1. Relative weight percent of carbonate minerals as a Figure 2.2. Geologic Map of Bermuda (Vacher* 1974). 9

dissolution, the Walsingham formation, an eolian limestone, contains a complex network of caves, many of which are interconnected and display stalactites, stalagmites and flowstones of various forms

(Harmon et &1., 1978). The younger Belmont, a marine deposit

(2 20,000 years), has a less developed karst system characterized by fractures and solution channels and primary features similar to those of the Walsingham. The Paget is composed of essentially unaltered, uncemented marine and eolian units of approximately

125,000 years in age (Bretz, 1960; Land et al., 1967; Vacher, 1973;

Hannon et al., 1978).

Soils

Two varieties of paieosols ace observed as interbedded strati- graphic units of the marine and eolian Bermuda limestones (Ruhe,

1961; Vacher, 1973). The Terra-rosa is a red, clayey soil which owes its color to ferric iron oxide/hydroxide which is a major component. The Regosol is a white or brown soil which can also contain some ferric iron oxide/hydroxide (Ruhe, 1961; Vacher, 1973).

There are strong indications that these soils owe their genesis to deposition of atmospheric dust (Bricker and Mackenzie, 1970), but it has also been suggested that they may be the residual of weathered volcanic debris and impure limestones (Blackburn and Taylor, 1969).

The following minerals have been identified as soil components: anatase, ilmenite, chromite, peroviskite, crandallite, epidote, garnet, sphene, native gold, quartz, gibbsite, iron oxide, kaolin and vermiculite (Table 2.2; Ruhe et al., 1961; Blackburn and Taylor,

1969; Bricker and Mackenzie, 1970). 10

Table 2.2. Minerals identified as components of Bermuda soils (Blackburn and Taylor, 1969; Bricker and Mackenzie, 1970; Ruhe et al., 1961).

anatase Ti02 chromite FeCr2° 4 crandallite CaAl3 (P0 4 )2 (0 H)5 -H20 garnet A3B2 < Si04 > 3 native gold Au gibbsite Al(OH)3 kaolinite Al2 (Si205 )(OT)4 ilmenite FeTiOj perovskite CaTiOj epidote Ca2 (Al, Fe )A120( Si04 ) (Si20?) (OH) sphene CaTi(Si04 ) quartz Si02 iron oxide F e , 0 , 4 J (hematite) vermiculite (Mg,Ca)0 <3 (Mg,Fe,Al)3 0 (Al,Si)4i 11

Carbonate Diaqenesls

The alteration of marine derived carbonate sediments has been addressed by many authors including Berner (1966), Hanshaw and Back

(1979) and Stringfield and Anders (1979). Carbonate deposits have been studied in much detail because of their high degree of chemical reactivity and their potential as highly productive aquifers

(Hanshaw and Back, 1979) and oil reservoir rocks (Chapman, 1976). A carbonate aquifer observed today did not always possess the characteristics of porosity and permeability required of a water bearing formation. These essential properties were acquired over an extended period of time by an alteration process termed diagenesis

(Hanshaw and Back, 1979).

Many carbonate sediments are formed in shallow water marine environments primarily by the accumulation of biologically produced calcium carbonate. Many marine organisms are able to precipitate various polymorphic forms of the mineral from the near surface waters and incorporate it into their skeletons. It is this skeletal material which generally forms very thick calcium carbonate sedimentary deposits.

The surface ocean, in contrast to deep ocean waters, is chemically supersaturated with the carbonate minerals calcite, aragonite, dolomite and magnesite, but rarely if at all does one observe the precipitation of these minerals inorganically (Berner,

1966; 1980; Folk, 1974; Hanshaw and Back, 1979; Cook and Kepkay,

1984). Studies have shown that inorganic precipitation of many of these minerals may be inhibited by kinetic processes. Suess (1970) 12

demonstrated that dissolved organic carbon may react with carbonate mineral surfaces and physically isolate particulates from the

surrounding waters greatly decreasing reaction rates. Likewise, dissolved magnesium and also phosphorus may interact to inhibit calcium carbonate nucleation kinetics (Pytkowicz, 1965; Simpson,

1964; Berner, 1967; Norse, 1974; Kitano et al., 1978).

The two minerals associated with recent carbonate sediments are aragonite and high magnesium calcite. Aragonite (orthorhcmbic crystalline structure) is a metastable polymorph of calcite

(hexagonal rhombohedral crystalline structure) (Fyfe and Bischoff,

1965; Folk, 1974). Because of its structure, aragonite is able to incorporate the elements Sr, Ba, K and Pb into its structure more easily than calcite. High Mg-calcite is also metastable with respect to low Mg-calcite and dolomite. The rhombohedral structure of calcite is commonly associated with higher concentrations of Mg,

Fe, Zn, Cu, Na and Cd compared to aragonite. The prefix "high Mg" is indicative of calcite with 12-30% Mg, while "low Mg" dipicts less than 4% (Folk, 1974; Hanshaw and Back, 1979).

Geochemical evidence suggests that alteration of these metastable minerals may not begin until they are flushed with fresh meteoric waters (Berner, 1966; Folk, 1974; Hanshaw and Back, 1979).

Initially these sedimentary deposits may have high porosity, but low permeability. Through dissolution and recrystallization via fresh water higher in dissolved carbon dioxide than the atmosphere due to its production in bacterially active soils, the limestones are altered and can become highly productive aquifers (Berner, 1966; 13

Pingitore, 1976). The lower ionic composition of freshwater minimizes the kinetic effects discussed above.

In general, diagenetic reactions account for alteration of the initial metastable minerals, aragonite and high-Mg calcite, to the more stable form, low-Mg calcite. This is commonly reflected by increased dissolved species such as Mg and Sr in the associated waters (Dodd, 1966; Langmuir, 1971; Plummer et al., 1976; Hanshaw and Back, 1979). Aquifers are normally divided into two major zones, the unsaturated or vadose zone which periodically shows fluctuations in moisture content, and a saturated or phreatic zone

(Pingitore, 1976; Freeze and Cherry, 1979). In carbonate aquifers the phreatic zone is where diagenetic reactions primarily occur

(Plummer, 1976; Back and Hanshaw, 1970; Pingitore, 1976).

in coastal and island systems, where carbonate rocks occur, freshwater recharge forms a low salinity upper lense (described

2+ — below) with a chemical signature dominated by Ca and HCO^ . These waters are underlain by waters high in Na+ , Ca2+, SOj2- and Cl- derived front mixing with sea water (Hanshaw and Back, 1979; Plummer and Back, 1960). As a result, recrystallization of aragonite and high Mg-calcite dominates in the freshwater zone (Hanshaw and Back,

1979).

At the boundary of the fresh lense and sea water, an intermediate salinity zone of dispersion (transition zone) develops

(Todd, 1959; kohout, 1960). Under certain mixing conditions this zone may display undersaturation or supersaturation with respect to calcium carbonate minerals. This is a result of the non-linear 1 4

relationship between mixing and mineral solubility (Plumner, 1975;

Wigley and Plummer, 1976; Plummer and Back, 1980).

Other reactions important to carbonate mineral solubility are those which effect changes in pH and produce or consume HCO^~ and

COj . The microbial degradation of organic matter which is coninon to all natural water systems is one such reaction and will be discussed later in more detail (Berner, 1966; 1969; Berner et al.,

1970; Friedman, 1972; Lyons and Gaudette, 1979; Bimbaum and

Wireman, 1984; Wood, 1985).

Ghyben-Herzberg Lenses

Fresh groundwater in Bermuda is found in five separate Ghyben-

He r zbe rg lenses. These lenses have been described by vacher (1974) and Plummer (1976), and additional aspects of lense hydrogeology and physiology have been studied to determine the geological and physical processes which affect their hydrology (Vacher, 1978a;

1978b; 1980; Ayers and Vacher, 1980).

The interaction of fresh groundwater with seawater in coastal areas was realized as early as the 1820s by Du Commun (1828); however, it was not until the turn of the century that two European investigators, Ghyben and Herzberg, working independently, described the physical properties which govern freshwater-seawater dynamics

(Todd, 1959; Reilly and Goodman, 1985). The Ghyben-Herzberg relation shows that fresh water, being less dense, will float above sea water in the aquifer and that the elevation of the fresh water table above sea level can be determined by the density difference of the two waters (Fig. 2.3). Fresh water has a density (Pf) of "1.000 15

Wt sea level

seawater freshwater

seawater

Figure 2 3 . Cross-section depicting the Ghyben-Herzberg relationship. This diagram assumes the two eaters to be In hydrostatic equilibrium (R eilly and Goodman, 1985). 16

kg/1, and seawater of 35°/oo has a density (Pg) "1.025 kg/1. In addition, the two investigators assumed the two waters to be at hydrostatic equilibrium and immiscible. Hie Ghyben-Herzberg relation is given as:

pshs - pfhs + pfhf

pshs - pfhs . f

ps - pf hf hs — - h; ^ ■ 40

This equation predicts that the depth of the seawater- freshwater interface is 40 times greater than the elevation of the freshwater table above sea level. Although this relation is a good first-order description, the assumptions of hydrodynamic equilibrium and immiscibility do not account for recharge, vertical and horizontal flow within the fresh groundwater, discharge and the zone of dispersion which are observed under field conditions (Fetter,

1972). Figure 2.4 portrays a somewhat more realistic system.

The fresh water lense is recharged by meteoric waters from above. This water moves vertically and horizontally and eventually mixes with the sea water in the zone of dispersion and is discharged as a brackish water along the coastline (Kohout, 1960; Fetter,

1972). This type of mixing process causes a cyclic influx of sea water from below and as a result the Ghyben-Herzberg interface

(which is now given as the depth of 50:50 seawater-freshwater 17

Wt sea level

mixing zone

seawater freshwater

seawater

Figure 2.4. Cross-section depicting the Ghyben-Herzberg relationship. This diagram depicts a more realistic system of freshwater-seawater mixing, cyclic influx and discharge to the inshore waters (Reilly and Goodman, 1985). 18

mixture) departs from the theoretical and is located further seaward

(Kohout, 1960). This has been observed along several coastal aquifers including the east coast of Florida where discharges have been calculated to be on the order of 45.5 m^/day from a i m wide strip of the aquifer (Kohout, i960).

Oceanic island Ghyben-Herzberg hydrology follows the same basic principles with the addition that sea water may underlie the entire island (Todd, 1959; Fetter, 1972). This results in a distinct freshwater nucleus affected by recharge, discharge, island shape and size, permeability and also tidal and atmospheric fluctuations (Todd, 1959; Vacher, 1978).

In real situations, additional factors are important to understanding coastal aquifers and island freshwater lenses. The previous discussion assumed the existence of a homogeneous, isotropic geologic aquifer. This allowed some symmetry to the thickness of the zone of dispersion. In general, however, this is not the case and Bermuda is an excellent example (Vacher, 1974).

The three dominant geological formations differ considerably in porosity and permeability (hydraulic conductivity) in addition to their carbonate mineralogy (Land, 1970; Vacher, 1974; 1978a; 1978b).

Nearly all of the fresh groundwater occurs in the Paget limestone formation (Vacher, 1974; 1978a; Plummer, 1976). The Paget is permeable enough to be the ideal aquifer material for storage of freshwater, but not permeable enough to allow significant mixing between fresh and sea water. This results in a relatively thin zone of dispersion. Where fresh water overlaps from the Paget into the 19

Belmont a much thicker zone of dispersion exists (Fig. 2.5; Vacher,

1974; 1978a; Plummer, 1976). In addition, groundwater development

in the more permeable Belmont limestone formation is subject to more

rapid upconing of sea water.

The five fresh water lenses in Bermuda are shown in Figure 2 2.6. The Devonshire lense is the largest, covering 10 km , and to

date the most extensively developed. It is believed to contain

7.57xl0®m^ of water (Vacher, 1974). More recent studies have

indicated some potential for the development of the four smaller

lenses and efforts have begun to exploit these waters.

Recharge of the lenses is from rainfall which averages 140

cnyyear (Macky, 1957). It has been estimated that about 87% of this water which reaches the soil zone in Bermuda returns to the

atmosphere via evapotranspiration (Vacher, 1974). More recent

estimates, however, put this figure closer to 75% (Thomson, Public

Works, personal communication).

Redox Reactions

The oxidation of organic matter via the reduction of oxygen, nitrate, iron, manganese and sulfate has been shown to occur in groundwaters in Bermuda (Simmons, 1983). These suboxic and anoxic

reactions generally occur in marshes or with cesspit contamination

(Simmons, 1983). Reactions of this type are common to oxygen- limited natural water systems including stagnant marine basins, lake hypolimnia, marine and lake sediments, marshes, bogs and contaminated groundwaters (Rittenberg et al., 1955; Knull and

Richards, 1969; Berner et al;, 1970; Fanning and Pilson, 1972; 20

PAGET FORMATION BELM 0II1 C0RMA1I0N

10V.

50 V .

3

Figure 2J5. Hydrogeologic cross-section through the Devonshire lense at Prospect (reproduced from Vacher> 1974). IENS

FRESH WAIEI

■RAKISH A NO

s e a w a t e r

DEVONSHIRE I E N S \

SOMERSEI IENS

Figure 2.6 Contour map showing the location of fresh water (Vacher 1974). 22

Brewer and Murray, 1973; Froelich et al., 1979; Kuivila and Murray,

1984; Baedecker and Back, 1979).

These reactions are bacterially mediated and have profound effects on chemical speciation, the dissolution of biogenic and/or detrital minerals as well as deposition of authigenic minerals and have been termed "biogeochemical" reactions (Berner, 1981). The reactions are usually described in terms of the reduction of a particular electron acceptor, for instance, nitrate reduction, iron and manganese reduction, and sulfate reduction. The primary role of the reactions, however, is to oxidize organic matter to carbon dioxide and water (Fenchel and Blackburn, 1979).

Examples of these reactions are given in Table 2.3 in a highly generalized form. Although the electron acceptors have well defined chemical formulas, the composition of organic matter can range from simple structures such as glucose (CgH^Og) to highly complex long chain highly substituted molecules (Fenchel and Blackburn, 1979).

In natural marine systems, algal planktonic matter has an average composition of 106:16:1 C:N:P (Redfield et al., 1963). This is termed the Redfield ratio, and it is described by the formula

(Cl^O^Qg (NHj)^g (POj); however, in freshwaterlakes, marshes, bogs and soils the C:N:P ratios differ in accordance with the native flora and fauna (Fenchel and Blackburn, 1979). In landfills, which may range from domestic garbage to synthetic organic compounds, the

C:N:P ratio may change drastically (Baedecker and Back, 1979).

The order in which the above reactions occur has been rigorously described thermodynamically (Stumm and Morgan, 1981).

Reactions yielding a high degree of free energy are favored over the 23

Table 2.3. Sequence of oxidation-reduction reactions (Champ et al., 1978; Baedecker and Back, 1979). c h 2 o + o2 > co2 + h 2 o

5CH20 + 3N03" + 17H+ ---> C02 + N 2 + NH3 + 8H20

CHzO + 2Mn02 + 4H+ ---> 2Hn2+ + 3H20 + C02

CH20 + 4Fe(OH)3 + 8 H+ ---> 4Fe2+ + H H 20 + C02

2CH20 + S042- + H+ --- > HjS + H20 + COj

2CH20 + C02 ---> CH4 + 2002 2 4

less energetic reactions (Claypool and Kaplan, 1974; Froelich, 1979;

Berner, 1980). It is thought that under normal microbial substrate

conditions each reaction is complete when one of the reactants

becomes limiting, in natural water systems, organic matter (ON)

usually occurs in excess and the electron acceptors become

successively depleted. In most cases, each reaction must go to

completion before the next one begins, but in some instances there

may be some overlap of the zones due to enhanced substrate (ON)

concentration or the development of more highly reducing microniches

within a more oxidized zone (Fenchel and Blackburn, 1979).

For sediments, the depth of each zone is dependent on the rate

of sedimentation, amount of organic matter and the concentration of

electron acceptors, while in groundwaters the rate of groundwater

transport is also important (Baedecker and Back, 1979; Berner,

1980). Because of the high sulfate concentration in seawater,

sulfate reduction is the most important reaction in the consumption

of organic matter in nearshore marine sediments. In freshwater

systems, the sulfate concentration is usually low and becomes depleted very rapidly so that nitrate reduction, fermentation and methanogenesis are usually quantitatively more important. The

reactions involving the reduction of iron and manganese oxides/oxyhydroxides are important to the formation of authigenic minerals such as FeS, FeCO^, Fe^PO^^'SHjO, MnS, and NnCO^. In comparison to' the oxygen, nitrate, sulfate and fermentation

reactions, they have been assumed to be of lesser significance

(Fenchel and Blackburn, 1979). 25

The oxidation of organic matter via microbially mediated electron transfer reactions can cause major changes in water

chemistry. These changes include depletion and removal of oxygen, nitrate and sulfate and the production of carbon dioxide, ammonia, phosphate, nitrogen gas, nitrous oxide, bicarbonate, hydrogen sulfide, reduced iron and manganese, and methane along with the production or removal of hydrogen ions (pH) (Berner, 1968; Berner,

1970; Ben-Yaakov, 1973; Froelich et al., 1979; Friedman and Foner,

1982; Kuivila and Murray, 1984). Changes in pH are usually marked by precipitation or dissolution of pH-dependent minerals such as carbonates.

Geochemistry of Bermuda Groundwaters

The geochemistry of groundwater from an oceanic island such as

Bermuda can be expected to reflect the chemistry of both atmospheric precipitation and of seawater. Open ocean seawater maintains a constant composition of the major ions and atmospheric precipitation shows a strong imprint of sea salt, albeit very highly diluted and subject to minor deviations due to atmospheric dusts and photochemical reactions (Table 2.4; Table 2.5; Stumm and Morgan,

1981; Church et al., 1982).

A first approximation of the chemistry of groundwater from the

Bermuda Ghyben-Herzberg lense may be represented by a mixing model of rain water with seawater resulting in a groundwater which is some mixture of the two end members. However, when calcium carbonate dissolution and the biogeochemical reactions described above are considered, this model becomes too simplistic. 26

Table 2.4. Composition of seawater (Stumm and Morgan, 1981).

moles/kg molar ratio (X/Cl)

Na+ 0.468 0.858

Mg2+ 0.0531 0.0974

Ca2+ 0.01028 0.01886

Cl" 0.545 -

K+ 0.0102 0.0187

Sr2+ 0.00009 0.00016

0.0282 0.0518 V* h c o 3“ 0.00233 0.00428

Br" 0.00084 0.0015

F“ 0.000068 0.000125 27

Table 2.5. Composition of Bermuda rain (Church et al., 1982).

umoles/1 molar ratios (X/Cl)

Na+ 148.0 0.775

Mg2+ 20.0 0.1047

Ca2+ 7.65 0.040

K+ 4.03 0.021

S042* 18.15 0.095

Cl- 191 28

Detailed studies of Bermuda's groundwater geochemistry have

been conducted by Plummer et al. (1976). With the in-depth mapping

of the fresh and brackish groundwater lenses provided by Vacher

(1974), Plummer et al. were able to delineate the pathways of

alkaline earth (calcium, magnesium, strontium) consumption and

production during diagenesis in Bermuda limestone aquifers.

They have determined that the chemistry of Bermuda groundwater

is controlled by three main factors: (1 ) the generation of carbon

dioxide gas in the soils and marshes; (2 ) dissolution of metastable

carbonate minerals; and (3) mixing of freshwater and seawater, in

marshes, high PCC>2, generated via organic matter decomposition

reactions, produces conditions which are undersaturated with respect

to calcium carbonate. While away from the marsh, gaseous carbon

dioxide evasion results in lower PCO2 thus decreasing the

undersaturation of these waters. The dissolution of carbon dioxide

in water causes the reactions

c o 2 + h 2 o < ------> h 2 c o 3

H2C03 <---- > H+ + HC03"

HC03" <---- > H+ + C032-

which result in lower pH values and in turn can cause calcium

carbonate dissolution by the reaction

CaC03 (s) + H+ <--- > Ca2+ + hco3”

Where dissolved carbon dioxide has been lost, supersaturation and

precipitation of carbonate minerals seem to occur within the aquifer 2 9

material (Plummer et al., 1976). The resulting aquifer dissolution, therefore, accounts for the elevated calcium, magnesium and strontium not explained by conservative freshwater-seawater mixing.

Plummer et al. (1976) observed that the diagenetic processes of the Bermuda carbonates do not provide a significant source of excess magnesium to the groundwaters. Magnesium enrichments which do occur, occur in the younger (125,000 yr.) less altered Paget limestones which have a higher composition of the mineral Mg- calcite, but in the older Belmont limestones (220,000 yr.), which have far less Mg-calcite, the seawater-freshwater mixing accounts for all of the observed magnesium. In contrast, strontium levels as much as forty times higher than concentrations predicted by conservative seawater-freshwater mixing have been shown to occur

(Plummer et al., 1976). High strontium indicates that recrystallization of aragonite to calcite is occurring in the limestone and the high strontium/calcium ratio shows that this transformation is primarily incongruent.

Changes in the degree of super saturation or undersaturation which result from changes in ionic strength created by seawater- freshwater mixing have been discussed earlier. Plummer et al.

(1976) have plotted calcite saturation versus percent seawater mixing for Bermuda groundwaters and the results indicate that ionic strength effects do not play any major role in the saturation state of the Bermuda groundwaters (Plummer et al., 1976).

Other geochemical studies of the Bermuda lenses include the work of Simmons (1983) and Simmons et al. (1985). These investigators have studied the organic matter decomposition 30

reactions mentioned above in detail and have documented biogeochemical zonation in depth profiles of the groundwater column.

Oxygen and nitrate both become depleted with depth, while ammonia, iron, manganese and hydrogen sulfide are all produced. Carbon dioxide is also an important product of these reactions and, as discussed by Plummer et al. (1976), accounts for the observed high

PCO2 in these waters. In marshes and other groundwaters in which these reactions occur, different biogeochemical reaction rates may be important to the degree and rate of carbonate diagenesis.

The sequence of events by which recharge waters in Bermuda derive their chemistries can be described as follows. Rainfall in

Bermuda averages approximately 140 cm/year and has the average chemistry given in Table 2.5 (Macky, 1957; Church et al., 1982).

Precipitation in equilibrium with atmospheric carbon dioxide at 25°C yields a pH of "5.6, but more acidic values in Bermuda averaging

4.74 have been measured and are believed to be the result of sulfuric and nitric acids transported eastward from the North

American continent (Church et al., 1982). This precipitation is concentrated by evapotranspiration, neutralized by reactions with carbonates in the soil and vadose zone and may also undergo ion exchange reactions with clay minerals such as vermiculite, montmorillonite and illite in the Bermuda soils (Blackburn and

Taylor, 1961; Lloyd and Heathcote, 1985). Because of the nature of clay minerals-at a natural water pH range (6-9), these reactions are predominantly cation exchange (Lloyd and Heathcote, 1985). Under normal conditions, some Na+ and K+ are replaced in the clay mineral structure by Ca^+, although at high ionic strengths such as with 31

seawater, reverse ion exchange reactions may occur (Lloyd and

Heathcote, 1985). As these recharge waters pass through the soil

zone, they are subject to high PC02 resulting from biological

activity in the root zone (Plummer et al., 1976). Therefore, when

these waters enter the vadose zone they are highly undersaturated with respect to calcium carbonate (CaCO^). Dissolution of calcium

carbonate accounts for increased concentrations of Ca^+ , Mg^+ and

2+ 2- Sr and also an equivalent amount of carbonate alkalinity (2C03 +

HOOj- ) in groundwater in carbonate terrains. In addition to

carbonate dissolution, biologically mediated redox reactio s, discussed above, can be responsible for changes in the sulfate

2- (SOj ) and titration alkalinity as well as other minor species such as N, P, Fe, and Mn. Since many of these components are common to both biogeochemical and carbonate processes, it is likely that the equilibria are interdependent. That is, changes in concentrations of reactants or products of one reaction may affect the equilibria of the others.

Plummer et al. (1976) have concluded that, as a result of the various reactions of carbonate diagensis, 0.32 cm3 of aragonite is recrystallized to calcite per annum and also that 5.2 x 10^ kg of calcium is lost from the aquifer material yearly.

Study Area

Two sites have been chosen to conduct this study, one in the

Devonshire Marsh system and the other in the Pembroke Marsh (Fig.

2.7). The reasons for choosing these sites are as follows. The groundwaters near these two marshes are some of the freshest Figure 2.7. Location of wells PRH and JR. 33

groundwaters on the island. This fact should allow one to detect possible changes in major ion chemistry that are not due to seawater-rainwater mixing. Also, the depth to water table near the marshes is very shallow and thus allows one to use bailing as well as a pumping technique to retrieve samples at specific depths in the groundwater column.

The site of the well in Devonshire Marsh is along Jubilee Road

(JR). The well is located at the southwestern end of the Devonshire

Marsh 300 m south of the Watlington Water Works utility and 100 m to the north and east of the Teceria stables. The area is mostly covered by grasses and trees in which are pastured a few horses and cows. Compared to the rest of Bermuda, this is a sparsely populated region. The only nearby houses are SO m or more to the south on the opposite side of Jubilee Road.

The Watlington Water Works extracts fresh groundwaters from the marsh using a network of horizontal piping called galleries which lie just below the land surface skimming water from the water table. The JR well is the standard type Bermuda groundwater observation well approximately 20 cm in diameter and lined with polyvinyl chloride (PVC) pipe at the top (~1 m) to prevent collapse at the surface. Simmons (1983) was able to collect water samples at depths up to 14 m. However, it appears that some collapse of the well may have occurred since this time, and the maximum depth now is at approximately 6 m.

In the Pembroke Marsh, the well PRH is located along Parsons

Road just opposite the Pembroke Rest Home in a parking lot adjacent to a playground. The Pembroke Marsh landfill is situated 100 m to 34

the north, and it is believed to be hydrologically upgradient from the well (Vacher, 1974). The depth of water at PRH is approximately

5 m and the chloride data suggests that the groundwaters collected from this well are homogeneous from top to bottom. In an initial water quality study, Simmons (1985) determined that this well has not been affected by Pembroke Dump leachate migration. CHAPTER III

METHODS

Sample Collection

Water samples were collected by two techniques: (1) a bailing

sampler and (2) an inflatable packer pumping apparatus. Technique 1

is described in Simmons (1983) and Simmons et al. (1985). Hie bailer (Fig. 3.1a) is a polyvinyl chloride (FVC) tube of diameter

6.5 cm with one-way valves which allow water to enter the sampler and not to escape. Samples have been successfully collected at different depths in open boreholes by lowering the device by nylon

rope to the level of interest in the wells and then flushing the sampler by an up-and-down motion of the rope (Simmons, 1983).

Confirmation that the sampler is working properly has been demonstrated by correlating the observed chemistry in the wells with predicted and observed findings of previous investigations (Fig.

3.2; Vacher, 1974; Plummer, 1976; Simmons, 1983). Salinity, major ions and conductivities increase with depth and varying biogeochemical zones were observed in the vertical profile. Hie main question concerning this technique is: "Do water samples collected within a borehole accurately reflect the chemistry of the aquifer or do the observed chemical reactions result because of the open well?" In order to answer this important question, it was necessary to pump the wells and sample the recharge water. Pumping a well, however, can cause a mixing of the various aquifer zones. 36

OHE HAY VALVE

ONE HAY VALVE

Figure 3.1a. Water sampling apparatus a) Bailer. Simmons. Geotechnlcal Instruments. 37

A

AIR LINE

PACKER

80cm

6Qcm

ONE NAY VALVE

m ///m

\Ocm

Figure 3.1b. Water sampling apparatus b) Inflatable Packer. Simmons Geotechnlcal Instruments. PERCENT RS POSXKT RS I 3 * 61 81 m

• K W I- W t y z H H hI 1 MB II W y 0 o

UB 19

Figure 3.2. a) Percent relative salinity profiles at WPS. RS compares quite well with the b) idealized model where RS is zero in the surface water and 100% at depth (Vacher, 1974; Simoons, 1963). 39

Hence, the inflatable packer was designed to pump the wells at discrete levels while minimizing vertical mixing.

The packer (Fig. 3.1b) is comprised of a 6.5 cm ID length of

PVC pipe (outer sleeve) with 1 cm diameter holes drilled along its length. This outer sleeve is capped at the bottom and fitted with two plexiglass mountings; one 10 cm from the bottom and the other 10 cm from the top. The mountings are designed to hold two inflatable rubber inner-tubes much like the hub of a bicycle wheel. When placed in the well both inner-tubes are connected via polyethylene tubing to an air pump at the surface. A 2.5 cm ID PVC pipe water sample line extends from inside the 6.5 cm PVC sleeve through the top of the packer and is connected to a water pump at the surface.

At the end of the pipe (inside the packer) is attached a 2.5 cm PVC one-way foot valve. The device is placed in a well at the depth of interest and inflated. Water is pumped from the 60 cm interval in the well between the two packers presumably without disturbing the waters above or below. Due to lift limitations of shallow well pumps, as designed, the packer can only be used to depths of 8 m below the ground surface. However, with modifications greater depths are possible.

Cleaning

Samplers

Prior to sample collection, the water samplers described above were washed in hot soapy water, rinsed and soaked overnight in dilute HCl (5% v/v). They were finally rinsed copiously in tap water from the Biological Station to avoid contamination by HCl. 40

Hie pipe through which water samples were pumped was also washed and soaked in dilute HCl just as the samplers. The reason tap water was used over deionized water

Iron and Manganese

Samples for iron and manganese (Fe and Mn) were stored in 125 ml Nalgene LPE bottles. Cleaning of these bottles involved washing in hot soapy tap water and rinsing in tap water. The bottles were then taken into a positive pressure class 100 clean roan where they were again rinsed, first with deionized water (DI) and finally with distilled deionized water {DDI). They were subsequently filled with

50% v/v reagent grade HCl and allowed to leach for at least two weeks before rinsing again in DDI. The bottles were then filled with 1% (v/v) reagent grade HC1-DDI mixture. Precautions taken to minimize contamination include handling the bottles only in the clean room and wearing the appropriate clean room clothing (lab coat, shoes and polyethylene gloves). Samples were filtered through acid leached 0.45 pm Millipore filters. The filter apparatus was cleaned by washing, then storing in 10% reagent grade HCl. The bottles were kept in the clean room and used for sample storage after filtration. 41

Nutrients

Samples collected by technique 1 (bailing) were transferred from the bailer into 250 ml Nalgene HDPE bottles. The bottles were cleaned by washing and soaking in hot soapy tap water, then subsequently rinsing several times in tap water, then several times in DI water. Finally, the bottles were filled with DI water until used. Samples collected by the packer were first vacuum pumped into

4 liter precleaned glass bottles. These bottles had been cleaned to temporarily accommodate both iron and manganese as well as nutrient samples. These glass bottles were washed in hot soapy tap water then rinsed thoroughly in tap water then DI water. Next they were filled with approximately 20% v/v reagent grade HCl and DI water and left to clean for at least one week prior to use. The day before sampling, the bottles were rinsed copiously in DI water, in order to ensure that all the HCl was removed, the bottles were rinsed until the conductivity of the rinse water was the same as the conductivity of the DI water. Upon return to the Bermuda Biological

Station, the samples were filtered and transferred to the 250 ml polyethylene nutrient bottles described above. Nutrient samples were filtered through prerinsed (DDI) 2.4 cm GF/C glass fiber filters; the filter apparatus was prewashed in hot soapy water and rinsed copiously in tap water then DI water.

Blanks

Blank studies were performed by analyzing DI water for nutrients and DDI water for iron and manganese after processing them as groundwater samples. This included sample collection, filtration and storage procedures utilized for "real" samples, in all cases, 4 2

the blanks were determined to be negligible in comparison with the

detection limits (Table 3.2).

Field Methods

The wells were first bailed using technique 1. Water samples were immediately analyzed to determine temperature and redox potential (Eh), and reagents were added to preserve dissolved oxygen and hydrogen sulfide (manganese sulfate and alkaline iodide, and zinc acetate, respectively). Samples for nutrient analyses were stored in 250 ml Nalgene (TM) high density polyethylene (HDPE) bottles, pH and alkalinity samples in 125 ml linear polyethylene

(LPE) bottles and hydrogen sulfide and oxygen samples in 130 ml glass BOD bottles. Samples were generally collected at four to five depths. The packer was then placed in the well at depths corresponding to that of bailing and connected via a 2.5 cm ID PVC pipe to a 1 hp shallow well pump and via polyethylene tube to an air pump. The packer was then inflated by the air pump until the well was sealed and the water pump was started. Water was removed from the well at a rate of approximately 25 liters per minute for 20 minutes with the rate controlled by a PVC ball valve at the outlet, in order to avoid contamination of the samples by the pump, it was necessary to collect samples from the PVC pipe before it reached the pump. The pump was stopped after 20 minutes and the water was trapped inside the pipe by a pair of valves. Samples were vacuum pumped from the pipe directly into two 4 liter glass bottles; one for nutrients, alkalinity and trace metals (Fe and Mn) and the other for temperature, Eh, oxygen and hydrogen sulfide. Again, 43

temperature and Eh were determined immediately, while oxygen and hydrogen sul£ide were fixed using the reagents described above. The water samples were then transported rapidly to the Bermuda

Biological Station (BBS) for sample preservation and storage.

Analytical Methods

Sample Preservation

Sample preservation and storage was started within 1-3 hrs. after collection. The analyses of pH and alkalinity, temperature and conductivity were performed immediately upon return to the BBS, while nutrient samples were filtered and preserved by freezing at

-10°C. Trace metal samples were filtered and acidified to 1% v/v with quartz redistilled HNO^■

The analytical methods along with precision and detection limits are summarized in Tables 3.1 and 3.2. pH and alkalinity

The pH of each sample was determined using an Orion Research ion analyzer after standardizing with both pH-4 and pH-7 buffers.

The samples were allowed to equilibrate with the electrode for several minutes prior to taking readings. Laboratory temperatures were a fairly constant 25 +3°C. Alkalinity analyses were performed using the Gran potentiometric titration procedure described by

Edmond (1970). Forty ml sample volumes were titrated using a 2.00 ml Gilmont microburette with 0.12 N reagent grade HCl. The burette is capable of delivering volumes as small as 0.01 ml. Readings were taken after 0.5 ml additions and readings were made in millivolts.

Each sample was titrated to a pH of approximately 2.5. 4 4

Oxygen

Analyses were conducted using the standard Winkler titration described by Strickland and Parsons (1972). Each sample was fixed in the field by the addition of manganese sulfate and alkaline iodide to 130 ml glass BOD bottles then titrated after acidification

(concentrated reagent grade sulfuric acid) with 0.25 N sodium thiosulfate the day following sample collection.

Phosphate, nitrite, silicate and ammonium

Samples were thawed and analyzed approximately one week after collection by standard procedures described in Strickland and

Parsons (1972). Colorimetric measurements were made using the

Perkin-Elmer Lambda 3 UV-Vis spectrophotometer.

Sulfide

Like oxygen, sanples were preserved in the field in 130 ml glass BOD bottles by the addition of zinc acetate and analyzed using the method of Kloster and King (1977). The analysis was usually performed 1-2 days following sample collection on 50 ml aliquots of sample. The preparation of samples and standards for analysis was conducted in a glove bag under an inert (N2 gas) atmosphere, working standards were made up with analytical grade sodium sulfide and standardized by a sodium thiosulfate titration. The colorimetric determinations were made on the Perkin-Elmer Lambda 3

UV-Vis spectrophotometer.

Fluoride, nitrate, sulfate and bromide

These constituents were determined on the Dionex QIC ion chromatograph following the general analytical procedures outlined by the manufacturer and also Mosko (1984). This involved using the 4 5

HPIC AS3 analytical anion column with a 3 mM/2.4 mH sodium bicarbonate/sodium carbonate eluent and flow rate of "3 ml/min. A series of mixed standards was made up using the appropriate analytical grade salts and dilutions were made as necessary with DI water.

Chloride

Chloride was determined on 5 ml sample volumes using the silver nitrate (Mohr) titration described by Strickland and Parsons

(1972). A “0.028 N silver nitrate titrant was used, along with a

3.5 gm/1 potassium chromate indicator, standards were made by diluting standard seawater (Wormley water Cl - 19.354°/oo).

Calcium, magnesium, strontium, sodium and potassium

The major cations were analyzed by flame atomic absorption spectrophotometry on the Perkin-Elmer model 460 (at BBS) and model

2280 (at UNH) instruments. Standards were made using the appropriate laboratory analytical grade salts and analyzed along with diluted standard seawater (Wormley water) for comparison. The differences were within analytical precision (Table 3.2). Samples were diluted into the necessary concentration range (calcium 1:25, sodium 1:25, potassium 1:25, magnesium 1:250 and strontium 1:1.1) and along with standards they were acidified to 1% with quartz redistilled reagent grade nitric acid to which was also added 2 ga/1

La and 0.5 gn/1 Cs matrix modifier (Smith et al., 1981). iron and manganese

Iron and manganese were determined by flameless atomic absorption spectrophotometry on the Perkin-Elmer 373, HGA 400 furnace (at BBS) and the 2280, HGA 400 furnace (at UNH) using 46

analytical procedures outlined by the manufacturers. All samples were acidified to 1 % with quartz redistilled reagent grade nitric acid and analyzed by auto-sampler injection of 20 ul aliquots into a pyro-coated graphite AA furnace tube with a L'vov platform. The analytical AA furnace programs were as follows.

IRON

STEP DRY CHAR ATOMIZE

TEMPERATURE °C 120 1100 2700

RAMP (secs) 5 10 0

HOLD (secs) 45 20 3

MANGANESE

STEP DRY CHAR ATOMIZE

TEMPERATURE °C 120 1100 2700

RAMP (secs) 10 10 0

HOLD (secs) 45 20 3

miniflow - 100 ml/min

Conductivity

Conductivity was determined using the Markson model 10 portable conductivity meter. 47

Table 3.1. Summary of methods.

Analysis Methods References

Ca2+,Mg2+ ,Sr2+ ,Na+ ,K+ Flame AA Smith et al. (1981) P-E 460

Fe2+ ,Mn2+ Flameless AA P-E 373,HGA400

F_,N03' #SO42~,Br" Ion Chromat Mosko (1984) Dionex

C1 Silver Nitrate Strickland and Parsons (Mohr) titration (1972)

P043",N02 fNH4+ ,Si04 Colorimetry Strickland and Parsons P-E Lambda 3 (1972) UV, Vis spec.

Total Colorimetry Kloster and King (1977) Sulfide P-E Lambda 3 UV, vis spec.

Oxygen Winkler Strickland and Parsons titration (1972)

Alkalinity Gran Edmond (1970) titration pH Orion Res. Ion analyzer

Temperature Thermometer

Conductivity Markson model Kunkle (1984) 10 portable cond. meter 48

Table 3.2. Precision and detection limits.

Detection % Precision limit coefficient at Concentration nmoles/1 variation mmoles/1 2+ Ca 0.01 0.6 2.5 + Mg2 0.001 2 0.3 2+ Sr 0.005 2 0.05

Na"* 0.06 4.0

K+ 0.002 10 0.3

Cl“ 0.01 0.3 1.5

F- 0.01 5 0.1

Br” 0.02 4 0.2

2- SO 0.01 4 0.3 4 2- S 0.001 3 0.01

0.01 8.0

Alk 0.05

|j moles/1

NO, 5 100

NO, 0.01 1 2.5

NH. 0.1 1 1.5 3- PO 0.03 3 3

H4Si04 0.1 5 20

u g / i 2+ Fe' 3.7 8 30 2+ Mn 0 . 4 1 0 2 0 CHAPTER IV

MAJOR ION GEOCHEMISTRY IN TOO BERMUDA MARSHES

Die major studies dealing with carbonate geology and geo­

chemistry in Bermuda have been performed by Bretz (1960), Mackenzie

(1964), Land et al. (1967), Land (1970), vacher (1973), Plummer et

al. (1976), Harmon (1978) and others, but little work other than

with calcium, magnesium and strontium has been performed on the

major ions in the Bermuda groundwaters. The major ions are

important because they can yield a great deal of information about

the source of the waters and hence the chemical processes affecting

their concentrations in the subsurface. In this section, an attempt

is made to quantify the concentration of each of the major ions

sodium, potassium, calcium, magnesium, strontium, chloride, sulfate,

fluoride, bromide and also alkalinity and identify their important

sources and sinks.

Since, in a simplistic sense, groundwater in Bermuda should be a mixture of rain and sea water, the chemical composition of these end members should dictate the composition of any groundwater sample

if the percentage of each end member can be determined. The chemistry of Bermuda rain has been discussed by Church et al. (1982) and Jickells et al. (1982), and it appears that seasalt aerosol contributes over 75% of the major ions present. The sodium/ chloride, magnesium/chloride and potassium/chloride ratios are comparable to that of seawater ratios; however, photodissociation or 50

Figure 4.1 to 4.8. Pie charts comparing the composition of rain water, seawater and groundwater. Figure 4.1

H Na H K 11 Ca + Sr m Mg

13 HC03 fH Cl □ S04 52

□ Na n K m Ca + Sr m Mg

X CATIONS IN SEAWATER

90.27*

X ANIONS IN SEAWATER 53

I0.S7* Figure 4 .3

69.68*

% CATIONS AT PRH

2 98*

HC03 Cl S04

* ANIONS AT PRH 5 4

I7.60R

M a m

H Na 53.31* H K 0 Ca + Sr m Jig

2 64*

% CATIONS AT JR SURFACE

0 HC03 fH Cl m S04

X ANIONS AT JR SURFACE 55

F i g u r e i « 8

X CATIONS AT JR 150cm

10-15*

0 HC03 n Cl 0 S04

63 93*

X ANIONS AT JR )50cm 56

Figure 4.6.

Ca + Sr

7 6 *

X CATIONS AT JR 300 cm

12.5 2 *

[1 HC03 §1 Cl 0 S04

* ANIONS AT JR 300 cm

58

* CATIONS AT JR 600cm

0 HC03 n Cl m S 04

76 02*

* ANIONS AT JR 600cm 5 9

volatilization of chloride causes minor deviations (Tables 2.4 and

2.5; Stumm and Morgan, 1981; Church et al., 1982). The sulfate enrichment observed is believed related to sulfur-rich fossil fuel emissions from the North American continent transported eastward with the prevailing weather patterns (Church et al., 1982).

Elevated nitrate (also from fossil fuel burning) and sulfate have been used to explain the relatively low pH of Bermuda rain waters— average pH - 4.74 (Church et al., 1982).

Sodium is the dominant cation in rain water in Bermuda. It accounts for 71% of the cation concentration, with magnesium (19%), calcium (7.4%) and potassium (1.9%) being far less significant (Fig.

4.1). This compares closely to sodium in sea water, where sodium accounts for 77%, magnesium for 17%, calcium for 3.4% and potassium for 1.7%. The anions follow a similar trend, with chloride in rain comprising 84% compared to 90% in sea water, sulfate 15.9% in rain and 9.3% in sea water, and at the observed pH of rain water (4.74) the bicarbonate is negligible, while in sea water (pH ~8.2) bicarbonate comprises 0.38% (Fig. 4.2).

The groundwater data collected in this study show a marked increase in calcium and bicarbonate resulting in these two species becoming the dominant major ions (Fig. 4.1; 4.3); these results are similar to those reported by Plummer et al. (1976). In the low salinity waters at PRH, calcium and bicarbonate (carbonate alkalinity) make up 69% and 80% of the major cations and anions, respectively (Fig. 4.3), while at JR there is a steady decrease in the percent of calcium and percent of bicarbonate with increasing salinity (Fig. 4.4 to 4.8). It appears that the rain waters, which 60

contain C02 from the 2 production in the soil and vadose zone, are highly corrosive to the Bermuda carbonates and quickly become calcium- and bicarbonate-rich waters through soil-aquifer calcium carbonate dissolution. Carbonate dissolution and evapotranspiration are the major geochemical processes occurring in the fresh groundwater (Vacher, 1974); Plummer et al., 1976). Once seawater- freshwater mixing begins, however, sodium and chloride rapidly become more dominant (salinity > l°/oo).

In order to determine the magnitude to which Bermuda geology provides a source or sink for the major ions to the groundwater, the anion chloride was chosen as a reference species. Because chloride minerals are extremely soluble, and hence do not form authigenic minerals other than in highly evaporated waters, its concentrations change only as a result of dilution. Therefore, this study has assumed that chloride mixes conservatively in groundwater systems.

The chloride anion has been used previously in Bermuda groundwaters by Vacher (1974), Plummer et al. (1976) and Simmons (1983) as a conservative component and also to trace groundwater movement in other areas (Konikow, 1977; Lloyd and Heathcote, 1985).

Pembroke Rest Home (PRH)

At the PRH well in the Pembroke marsh the groundwaters are well mixed to at least 4 m. This is the depth of water sampled in the well. The chloride concentration is approximately 1.3 mM (Fig.

4.9). Results from vacher (1974), Plummer et al. (1976) and Simmons

(1983) suggest that the chloride concentration of the upper lens freshwaters is approximately 1.4 mM. This indicates that the

61 CHLORIDE (mM) depth depth of 4 and m do not contain chloride fran a *PRH FEB 11 CHLORIDE VS DEPTH" 1.0 1.2 1.4 1.6 1.8 2.0 0 100 4 0 0 3 0 0 200 Figure 4.9. Chloride versus depth suggests that groundwaters at PRH seawater. are well to mixed

DEPTH (cm ) 62

groundwater at the PRH well does not contain chloride from seawater and represents the upper unmixed freshwater of the Devonshire Lens.

Since it has been established that the groundwaters at PRH do not contain dump contamination in the upper 4 m (Simmons, 1985), it is believed that the observed chemical composition is not due to the landfill leachate. Therefore, it seems likely that the chemical composition of the groundwaters at PRH represents that of rainwater concentrated via evapotranspiration in the soil zone.

with the chloride rain data from Church et al. (1982) (Table

2.5) and the PRH chloride data, it should be possible to place some constraint on the average amount of evapotranspiration necessary to produce groundwater of this chloride concentration. In order to produce the chloride concentration of 1.3 mM observed at PRH, rainwater of chloride concentration of 0.191 mM must undergo 85% evapotranspiration. The same calculation performed using sodium yields an evapotranspiration value of 90%; however, because the extent of ion exchange reactions on the production of sodium in the soil zone is unknown, the chloride derived value is probably the better estimate. It should also be noted that the contribution of dry sea salt deposition has not been taken into account. If the rain dissolves these salts on contact, the rain's actual concentration just prior to evapotranspiration is underestimated by the Church et al. (1982) data. This error would in turn cause the calculated evapotranspiration (85%) to be slightly overestimated.

Using the average annual rainfall rate as 146 cm (Macky, 1957) and the evapotranspiration value above, one can calculate an annual recharge of ~21 cm at PRH. It is of interest that the 63

evapotranspiration has been previously estimated at 67% by Vacher

(1974) using the Pennman technique and 93% by Plummer et al. (1976) with the major ions.

Assuming that 65% is a reasonable estimate of evapo­ transpiration for PRH near the Pembroke marsh, and using the data of

Church et al. (1962), the major ion composition of the upper freshwater lens should have the baseline composition tabulated in

Table 4.1. No bromide, fluoride or strontium data are available for

Bermuda rainwater, but assuming negligible fractionation from seawater ratios, their concentration values are calculated to be

0.002 mM, 0.00016 mM and 0.00021 mM in rain water, respectively. If these groundwaters at PRH do represent the unmixed freshwater end member, then differences in the concentration with respect to the calculated evapotranspiration water must be a result of variability in annual rain chemistry or due to reactions in the soil zone and/or aquifer material (Fig. 4.10 to 4.17).

Calcium, magnesium and strontium show large enrichments over that calculated via evapotranspiration (Appendix IV). The major source of these elements to the groundwater is undoubtedly the diagenetic reactions with the carbonate bedrock as discussed by

Plummer et al. (1976). The sodium concentration is 1.49 mM while potassium concentration of "0.3 mM is one to two orders of magnitude greater than can be explained by evapotranspiration. Both sodium and potassium concentrations can be affected by ion exchange reactions in the soil zone, but due to the presence of large amounts of the doubly charged calcium ion, this mechanism is probably of 64

Table 4.1. Composition of fresh groundwater endmember after 85% evapotranspiration of rainwater.

Sodium 1.00 mM

Potassium 0.027

Calcium 0.052

Magnesium 0.136

Strontium 0.0014

Sulfate 0.122

Bromide 0.013

Fluoride 0.0011 65

Figure 4.10 to 4.17. Major ions calcium, magnesium, strontium, sodium, potassium, sulfate, bromide and fluoride versus depth at PRH. Samples collected by both pumping (P) and bailing (B) techniques. 66 SODIUM (mM) 0.0 0.5 1.0 . 5 1 2.0 0 100 4 0 0 200 300 Flgur* 4.10. "PRH FEB 11 SODIUM VS DEPTH*

DEPTH (cm ) 67 POTASSIUM (mM) POTASSIUM - - 0.0 0.0 0.1 0,2 0.3 0.4 0.5 100 4 0 0 3 0 0' 200 Figure 4.11. -pRH FEB 11 POTASSIUM VS DEPTH* DEPTH (cm ) 6 8 CALCIUM (mM) FEB 11 CALCIUM VS DEPTH* p r h - • • 100 3 0 0 ‘ 400 200 Figure 4.12. DEPTH (cm ) 69 MAGNESIUM (mM) ■ ■ 100 4 0 0 3 0 0■ 200 Flflure 4.13. "PRH FEB 11 MAGNESIUM VS DEPTH* DEPTH (cm ) 70 STRONTIUM (mM) ■ 0.000 0.005 0.010 0.015 0.020 100 4 0 0 3 0 0 ■ 200

DEPTH (cm ) Figure 4.14. *pRH FEB 11 STRONTIUM VS DEPTH’ 71 SULFATE (mM ) ■ ■ 0.00 0.05 0.10 0.15 100 4 0 0 3 0 0 - 200 Flgur# 4.15. “PRH FEB 11 SULFATE VS DEPTH*

DEPTH (cm) 72 BROMIDE (mM) * - 0.000 0.100 0.200 0.300 100 4 0 0 200 3 0 0 ■ Figure 4.16. “PRH FEB 11 BROMIDE VS DEPTH*

DEPTH (cm) FLUORIDE (mM) - ■ 0.00 0.01 0.02 0.03 100 4 0 0 3 0 0 200 Figure 4.17. "PRH FEB i l FLUORIDE VS DEPTH*

DEPTH (cm) 74

limited importance. Mechanisms which may explain the large potassium enrichments will be discussed subsequently.

Sulfate (average 0.122 mM) at PFH follows the evapo­ transpiration model. This suggests that sulfate input from marsh foliage or removal by bacterial reduction processes are negligible at this location. Fluoride and bromide have been shown to be present in the groundwater, fluoride at "0.004 to 0.02 mM and bromide at 0.01 mM. The bromide concentration is a good representation of the seawater ratio/evapotranspiration calculation, but fluoride shows enrichments of up to an order of magnitude over the calculated value. This suggests the presence of an important fluoride source.

Jubilee Road (JR)

Water samples were collected at JR on March 1 and May 16,

1985, by both the pumping and bailing techniques. The profiles show a general increase of salinity with depth in the water column resulting from the mixing of fresh groundwater with the underlying seawater (Fig. 4.18a and b). On March 1, the profiles of samples collected by pumping show a higher salinity water at a depth of 610 cm. This implies a density instability, a condition which can only occur under atypical groundwater conditions. The sample was analyzed on two different occasions ("June 1985 and "June 1986) for the major cations and anions and gave similar results, in addition, the ion balance requirements were not satisfied (37% difference between cations and anions), indicating possible contamination or a problem of mislabeling of bottles during sample handling. Because

CHLORIDE (mM) 75 1 ■ 600 400 - 400 200 Figure 4.18a and b. b. and Figure 4.18a general, that, depth in shows versus Chloride of mixing fresh depth, with a of result salinity increases groundwater withunderlyinggroundwater seawater. the Figure 4.18*. -JR MARCH i CHLORIDE VS DEPTH* DEPTH (cm) CHLORIDE CHLORIDE (m M ) 76 - 6 0 0 ' 4 0 0 ■ 200 Flflur* 4.18b. -J R M AY 1 6 CHLORIDE V S DEPTH’

DEPTH (cm ) 77

of this problem with this sample, it is viewed with much skepticism and not included in the plots or calculations.

As mentioned above, chloride variations with depth are believed to be the result of freshwater/seawater mixing. Variations have also been observed between the two sampling days. These variations can be explained by the fact that the freshwater lense and underlying seawater are dynamic in nature, and the degree of mixing is subject to tidal cycles, time and amount of precipitation, and atmospheric pressure changes. All of these processes have been shown to cause fluctuations in water levels of the Ghyben-Herzberg lense (vacher, 1974; Vacher, 1978) and probably also affecting the degree of mixing at the lense boundary.

Another interesting variation noted in the profiles of all major and minor ion species is the difference in chemistry of samples collected at the same depths, but using the different sanple collecting techniques (Fig. 4.18 to 4.26). At PFH the differences between samples collected by pumping and bailing were small for chloride, sodium, potassium, magnesium and sulfate, but calcium, fluoride and bromide show substantial differences. Jubilee Road pumped samples display distinctly higher concentrations for all the major ions. These results clearly show that groundwater bailed from a well can differ from groundwaters in the nearby aquifer. It appears that the groundwaters from the nearby aquifer are slightly more saline and display a higher degree of calcium carbonate dissolution than the samples taken from the well proper. In order to address whether or not the borehole samples truly are representative of the aquifer waters, it may be necessary to bail 78

Figure 4.19 to 4.26. Major ions calcium, magnesium, strontium, sodium, potassium, sulfate, bromide and fluoride versus depth at JR. Samples collected by both pumping (p) and bailing (B) techniques. SODIUM (mM) SODIUM ■ 600 • 600 200 ■ 400 Flgur* 4.19. "JR MARCH 1 SODIUM VS DEPTH* DEPTH (cm) SODIUM (mM) ’JR MAY 16 SODIUM VS DEPTH* ' 600 ■ 600 200 ‘ 400 Figure 19b. DEPTH (cm) 81 POTASSIUM (mM) POTASSIUM MARCH 1 POTASSIUM VS DEPTH* j r - • 600 ' 600 400 ' 400 200 Flgur* 4 .2 0 . DEPTH (cm) 8 2 POTASSIUM (mM) POTASSIUM *JR MAY 16 POTA331UM V3 DEPTH* 600 ■ 600 400 ■ 400

Figure 20b. DEPTH (cm) 8 3 CALCIUM (mM) ■ 600 - 600 400 ■ 400 200 Figure 4.21. *JR MARCH 1 CALCIUM VS DEPTH*

DEPTH (cm) 8 4 CALCIUM (mM) CALCIUM -JR MAY 16 CALCIUM V3 DEPTH* 2 2 3 4 5 ■ 400 ' 400 600 - 600 200 Figure 21b.

DEPTH (cm) 85 MAGNESIUM (mM) 0 0 1 2 3 4 ■ 600 ■ 600 400 ■ 400 200 Figure 4.22. 'JR MARCH 1 MAGNESIUM VS DEPTH' DEPTH (cm ) 86 MAGNESIUM (mM) 0 0 1 2 3 4 5 6 ■ 600 ■ 600 400 ' 400 200 Figure 22b. *JR MAY 16 MAGNESIUM VS DEPTH* DEPTH (cm ) Figure 4.23. - j r MARCH 1 STRONTIUM VS DEPTH*

STRONTIUM (mM) 0.03 0.02 88 STRONTIUM (mM) VR MAY 16 STRONTIUM VS DEPTH* .01 □ □ 200 400 6 0 0 - Figure 23b. DEPTH (cm) SULFATE (mM) 89 200 400■ 6 0 0- Figure 4.24. “JR MARCH 1 SULFATE VS DEPTH*

DEPTH (cm) SULFATE (mM ) 90 "JR MAY 16 SULFATE VS DEPTH" ■ 200 4 0 0■ 60 0• Figure 24b.

DEPTH (cm) 91 BROMIDE (mM) ■ 0.00 0.01 0,02 0.03 0.04 0.05 4.25. "JR MARCH 1 BROMIDE VS DEPTH" 200 40 0■ 600- F l g u r s

DEPTH (cm) 0.10 0.08 92 0.04 0. 06 BROMIDE (mM ) ■ 0.00 0.02 200 4 0 0' 6 0 0 -| Figure 25b. VJR MAY 16 BROMIDE VS DEPTH*

DEPTH O 3 0.200 0.100 93 FLUORIDE (mM)

- o -f— ? 0.000 200 400 6 0 0 ■ F1gur« 4<26. "JR M A R C H 1 FLUORIDE VS DEPTH' DEPTH (c; FLUORIDE (mM) -jR MAY 16 FLUORIDE VS DEPTH' • 200 40 0■ 6 0 0 { Figure 26b.

DEPTH (cm) 95

and pump several wells in the same area repeatedly and apply various statistical analyses to the chemical data.

All the data collected on March 1 and May 16, 1986, by both pumping and bailing for each ion were plotted together on mixing curves (i.e. vs. chloride) to establish whether or not a conservative relationship exists. Mixing curves assume that any chemical species, when plotted versus the reference species chloride, should fall on a straight line between the concentration of the two end members (Fig. 4.27). If this does occur, the species is said to "mix conservatively." Species which plot below the line indicate removal by processes such as authigenic mineral precipitation or by adsorption. Species which plot above the line infer the presence of a source to the groundwater such as dissolution or desorption.

The data were also entered into the thermodynamic program

MINBQL to determine the chemical speciation of each ion of interest

(Table 4.2a and b; see Appendix I for details on MINEQL). These results are presented in the following discussion.

Sodium. Sodium is a highly soluble species and, according to the thermodynamic calculations, is present completely as the free ion (99-100%). The sodium versus chloride plot exhibits a high 2 degree of correlation (r -99.6%) and a 0.849 ±0.013 slope. The slope of the theoretical mixing curve is 0.859 (Fig. 4.28). A t- test performed on the regression data confirms that the slopes are identical at the 95% confidence level. The fact that the sodium versus chloride plot is linear, coupled with the slope of the mixing curve (0.849) being the same as the theoretical, indicates that Figure 4.27. Ideal conservation mixing curve.

60

CHLORIDE (mM) 97

Table 4.2a. Major ion speciation (PRH) determined using the chemical speciation program M1NEQL.

%Free %S04 2 %C032- %HC03

Na 100 - - -

K 99.9 - - -

Ca 48 - 47.7 3.6

Mg 91.7 1.1 - 6.8

Sr 100 _ __

%Free %Ca2+ %Mg2+ %H+

Cl 10 0 - - -

S04 85.6 10.9 2.8 -

F 97.7 1.1 1.2 -

Br 1 00 - - - co3 2 1 . 1 76.8 98

Table 4.2b. Major ion speciation (JR).

%Free %S04 2 %C032 %HC03

Na 99.5 - --

K 98.7 1.3 - -

Ca 56.3 4.8 - 2.7

Mg 78.0 8.3 - 3.8

Sr 100 - - -

Ca 12.4% as CaMg(C03 h Mg 9.5% as CaMg(C0302

%Free %Ca2+ %Mg2+ %H+ %Na+

Cl 100 - - - so4 74.4 5.7 13.1 - 6.5

F 89.1 1.1 9.8 - -

Br 100 -- - - co3 - 13.0 - 68 -

13.7% as CaMg(C03 )2 SODIUM (m M ) conservatively in conservatively Bermuda groundwater. Figure 4.28. Sodium versus chloride mixing curve. Sodium Sodium mixes versus Sodium chloride curve.mixing 4.28.Figure 0 2 30- 10 50- 60 - - 0 Figure Figure 4 10 . 28 , LP = 0.849 = SLOPE 20 CHLORIDE (mM) 99 1.00 5030 40

100

sodium behaves conservatively in the groundwaters and that processes like ion exchange are o£ little significance in Bermuda soils.

Potassium. Like sodium, potassium is highly soluble, and speciation calculations indicate that it is present 99.9% as the free ion. A regression of the potassium versus chloride mixing 2 curve yields an r - 99.3% correlation, and thus the dominant source of the ionic species to the groundwater is probably seasalt.

Potassium, however, does not mix conservatively (Fig. 4.29). It is enriched to a maximum of 0.3 mM over that predicted by the mixing curve. These enrichments are higher at lower chloride concentrations and decrease with increases in chloride. This results in a potassium/chloride slope (0.0124 +0.00065) significantly less than that of the seawater/freshwater dilution curve (0.018). This potassium/chloride slope clearly indicates potassium enrichment by processes near the surface of the water table or in the overlying soils and not by dissolution of minerals in the aquifer material itself. At higher salinities there is less enrichment due to dilution with higher salinity waters from below.

Three possible sources of potassium to the groundwater exist: 1) ion exchange with clay minerals in the soils producing potassium- enriched waters; 2 ) dissolution of potassium-rich clay minerals in the soil zone and 3) leaching of potassium from plant debris.

Ion exchange is a surface reaction in which a dissolved species, such as potassium or sodium, becomes adsorbed to particle surfaces, replacing other less reactive species which in turn are released to solution. Surface reactions of this type are strongly pH-dependent (Drever, 1982). At very low pHs the surface of most POTASSIUM (mM) per nt omx osraiey Dt sget ta potassium that watertable. suggests Data Potassiumthe curve. mixing mix conservatively. chloride to not versus Potassiumappears 4.29. Figure ece fo pat ise n rsls n h bevd enrichmentobserved near the In and results tissue from plant leaches 0.000 0.2001 0.4001 0.6001 0.800H H 0 0 0 . 1 1.200 0 10 LP = .140 .98 0.0124 = SLOPE 101 20 CHLORIDE (mM) 30 40 50

60 102

clays is positively charged and therefore the dominant exchange

reactions occur between anions and the surface. At the pH of

natural waters (6-9), clay surfaces are negatively charged and the

reactions are commonly cation exchange. Depending on the nature of

the solution (high or low ionic strength) and the clay mineral

(saturated or unsaturated adsorption sites) the exchange reactions

could provide a source or sink for dissolved potassium.

Because the groundwaters in Bermuda have higher concentrations

of reactive divalent species, calcium (Fig. 4.1; 4.3; Appendix IV),

it is likely that a large portion of the adsorption sites are

already saturated with these more reactive species. Therefore, it

is unlikely that ion exchange would provide a significant source of

potassium to the groundwaters.

Various clay minerals are found in the soil zone (Table 2.2).

None of these minerals have significant potassium in their chemical

structures. In addition, the maximum dissolved silicate concen­

trations at JR are on the order of 60 to 70 MM (Fig. 4.30a and b).

Assuming that all of this silica came from the reaction of a

potassium-rich mineral, for example mica

KAl3Si3°10(OH)2 + H+ + 9h2° <----> SAKOH) + 3H4Si04 + K+ which would only account for about 20-25 MM excess potassium in

these groundwaters.

By a process of elimination, it therefore seems likely that

the source of excess potassium is marsh vegetation. Potassium is an

element required for plant growth; it is required in amounts greater

than all other nutrients except nitrogen (Smith et al., 1982).

70

6 0 SO 103 SILICATE (>iM) 40 *JR MARCH 1 SILICATE VS DEPTH* 3 0 ■ 200 4 0 0' 6 0 0 Figure Figure 4.30a. Figure 4.30a and b. b. and Figure 4.30a are JR only at Maximumsilicate concentrations for notthewould account excess probably minerals silicate on the order of 60 to 70 UM. UM. the 60 to 70 of on order potassiumdissolutionof This suggests potassium.

DEPTH (cm) 30 20 SILICATE (pM ) 104 10 ■JR MAY 16 SILICATE VS DEPTH* 1 ■ 200 4 0 0- 6 0 0 Figure Figure 4.30b. DEPTH (cm) 105

Potassium levels of normal plants can range from 15 to 50 mg/g

(Hewitt and Smith, 1975; Robb and Pierpoint, 1983). Potassium can

be leached from both live and dead plant cells, with the process

being more significant in the decomposing material (Smith et al.,

1982). The marsh system is largely covered by vegetation which is

in close contact with the groundwater, and thus much of the potassium losses from vegetation would be leached directly into the

groundwater system.

Calcium. Although the calcium versus chloride plot yields an 2 r » 97% correlation, it does not mix conservatively in the groundwater (Fig. 4.31). Calcium shows 2-3.5 mM enrichments over

that predicted by the mixing curve. This large source of calcium to the groundwaters is due to dissolution of the calcium carbonate in the soil zone and/or in the aquifer material. The results of

Plummer et al. (1976) have shown that the production of large quantities of carbon dioxide (OC^) in the marsh sediments via organic matter degradation is responsible for undersaturated conditions with respect to calcium carbonate and hence dissolution.

The slope of the observed mixing curve is 0.0435 +0.0022. The fact that calcium/chloride shows a strong positive relationship and that the slope of the plot is greater than that of the dilution curve suggests that the degree of dissolution may increase with increasing salinity. Plummer (1975) and Plummer and Wigley (1976) have shown that the mixing of two waters of differing ionic strengths and degrees of saturation can result in waters either supersaturated or undersaturated with respect to calcium carbonate. The results of the calcium/chloride plot presented here seem to suggest the CALCIUM (m M ) iigcre aeudutdy u o islto o h calcium the of dissolution undoubtedly dueto are curves mixing abnt aufr material. aquifer carbonate iue43. ag clim nihet oe ta dslyd y the by displayed that over enrichments Large calcium 4.31. Figure 4 5 3 2 1 0 0 10 a a SLOPE = = SLOPE 106 20 CHLORIDE (mM) .45 = 0.99 = R 0.0435 30 40 50

60 107

development of a higher degree of undersaturation. Plummer et al.

(1976), however, have suggested that seawater-freshwater mixing is not a significant factor contributing to undersaturation in the

Bermuda groundwater and that the high production of C02 is more important to calcium carbonate dissolution in groundwaters associated with marsh areas. Calcium complexes include free Ca2+

48%, Ca2+ C032- 29.6% and Ca2+ HC03“ 3.6%.

Strontium. Strontium concentrations at JR are observed to be

0.01 to 0.02 mM, or approximately 0.01 mM greater than that predicted by the dilution curve (Fig. 4.32). The excess strontium is a product of the diagenetic recrystallization of aragonite to the more stable calcite. High strontium concentrations and high strontium to chloride ratios are indicative of the incongruent dissolution of aragonite (Harris and Mathews, 1968; Baker et al., 2 1982). The strontiuq/chloride regression yields an r - 84.5%, a somewhat less significant relationship than that of other major ion species previously discussed. The strontiun^/chloride observed slope

- 0.0003 +0.000036; the predicted value is 0.00016.

Magnesium. The magnesium content in carbonate rocks can be fairly high (12-30%, Hanshaw and Back, 1979), and dissolution can provide a significant source to the groundwater. However, Plummer et al. (1976) found that Bermuda limestones do not contribute significantly to magnesium enrichments in the groundwaters. Some rock-derived magnesium has been shown to occur in groundwaters from the younger Paget limestones; however, most groundwater magnesium is derived from seawater mixing (Plummer et al., 1976). The results presented here support their findings and indicate that magnesium rocryslaI 1 Izatlon of aragonite to tho more stab le c a lc lte mineral lte lc a diagnostic c of le product stab more a tho Is to strontium aragonite of excess 1 The Izatlon rocryslaI 4.32. Figure polymorph. Sinom iU M (mM) O.CO- .1 H 0.01 0.C2H 0.03 0 10 20 108 CHLCRICE (mM) CHLCRICE 30 0

60

109

mixes conservatively, but theoretical and actual curves are offset by about 0.1 mM. This value probably represents the amount of magnesium input via dissolution of high Mg-calcite in the vadose zone (Fig. 4.33). The theoretical slope of the mixing curve is

0.097, and the observed slope is 0.096 (r2 - 97.8%), identical at the 95% confidence level.

Sulfate. The shape of the sulfate/chloride curve suggests that sulfate also does not mix conservatively in the groundwater

(Fig. 4.34). The depletions in sulfate in the lower chloride waters are probably the result of microbial sulfate reduction to sulfide occurring in the marsh, and enrichment may again be due to the breakdown of plant debris, as was described above for potassium, and/or the oxidation of previously produced sulfide. The average sulfur component of normal plants can range from 1000 to 3000 ppm

(Robb and Pierpoint, 1983).

At chloride concentrations less than 10 mM, sulfate is neither enriched nor depleted. For the chloride concentrations between 10 mM to ~35 mM, sulfate displays a minor but consistent depletion which is undoubtedly due to microbial sulfate reduction. The few data points above 35 mM (35 to 60 mM chloride the deeper samples) show enrichments. The major sulfate species include free sulfate

8 6 % and conplexes with Ca2+ 11% and Mg2+ 3% (Table 4.2).

Brocdde. Bromide appears to mix conservatively with no source or sink term in the Bermuda limestone aquifer (Fig. 4.35). A t-test analysis of the slope (0.0014) is equal to that of the theoretical

(0.0015) to 95% confidence. MAGNESIUM (m M ) M n my ersn mgeim isle I te aoe zone. vadosemM the In magnesium and may represent dissolved osraiey bt hoeia ad cul uvs r ofe by offset are curves actual and theoretical but conservatively, iue .3 Rsls ugs ta magnesium that fairly mixes suggest Results 4.33. Figure 0 2 3 4 5 1 6 0 10 LP = .9 99 0 0.097 = SLOPE 20 CHLORIDE (mM) 110 30 40

50

60 SULFATE (m M ) iue43. uft dpein ae rbby u t irba s lfate su microbial from plant probably to dueizatlon renlneral are may enrichments duedebris. to be while depletions reduction, Sulfate 4.34. Figure 0 1 2 4 3 5 0 10 SLOPE = = SLOPE 20 111 CHLORIDE CHLORIDE .7 R 0 94 0 = R 0.074 30 (mM) 40 50 60

groundwater.

Figure 4.35. Bromide appears to appears mixconservatively Bromide in Bermuda 4.35.Figure BROMIDE (m M ) 0.00 0 0.04- 0.06- 0.08- 0 1 . 0 . 02 0 - 10 SLOPE = = SLOPE 20 112 HOIE mM) M (m CHLORIDE .04 = 1.00 = R 0.0014 30 40 50

60 113

Fluoride. Fluoride displays a very large enrichment relative

to the conservative mixing line (Fig. 4.36). This enrichment is

greatest at the higher chloride concentration and accounts for a

fluoride/chloride slope steeper by an order of magnitude than that

of the dilution curve. The maximum fluoride concentration is 0.18

mM; this is greater than the fluoride concentration observed in

standard seawater (0.068 mM).

There are three possible sources of fluoride to the

groundwaters. They include fluoride contamination from outside

sources like drinking water fluoridation, dissolution of fluoride

minerals in the aquifer such as fluorite and fluorapatite, and

release by decomposing plant debris in the marsh. There are no

known anthropogenic sources of fluoride near the sampling location,

and fluoridation of drinking water is not in general practice in

Bermuda. Therefore, it is unlikely that this is an important

source.

Fluorite is a common mineral associated with limestones and

dolomite (Carpenter, 1969; Minear and Keith, 1982). Many marine

organisms utilize fluorite in their tissues, thus accounting for

fairly large quantities in shelf sediments (Lowenstam, 1981).

Likewise, fluorapatite occurs in marine sediments as a result of

calcium carbonate-phosphate reactions. (The role of apatite is

discussed further in Chapter VI.) Solubility calculations (Appendix

I) indicate that the two minerals fluorite and fluorapatite are

undersaturated in the groundwaters, and this suggests that if these mineral phases are present in the aquifer material, they should be dissolving. But, X-ray diffraction analysis performed on the 114

0.2- SLOPE-0.0037 R - 0,98

0.0 i r 1 i 10 20 30 40 50 60

CHLORIDE (mM)

Figure 4.36. Fluoride displayB very large enrichments relative to the conservative mixing curve. The data suggests that a deep aquifer source may be responsible. 115

carbonate aquifer material show no Indication of either fluorite or apatite. Samples of both limestone and soil were dissolved and analyzed by ion chromatography for fluoride. Ihe material contained less than 0.01 mmoles fluoride per gram. If the limestones were responsible for fluoride of "0.18 mM in the groundwater, it would be necessary to dissolve at least 18 g of solid per liter of water.

The highest calcium concentrations suggest dissolution of at most

~0.4 g calcium carbonate per liter. These results clearly suggest that the aquifer carbonates do not provide a fluoride source to the groundwaters.

The third possible source is fluoride released from decaying plant material. Many species of plants have been shown to accumulate fluorine to levels up to 120 Ug/g (Robb and Pierpoint,

1983). Mun et al. (1966) have shown that peat sediments contain about 0.1% dry weight fluoride, and Windom (1971) has determined average fluoride levels in the rhizomes and leaves of salt marsh vegetation along the Georgia coast to be 1.4 and 2.0 ppm, respectively. Windom (1971) concluded that the decomposition of plant material is responsible for fluoride enrichments observed in

Georgia coastal marsh pore waters. Although these results are similar to the enrichment that has been observed in this study, it seems unlikely that this would explain the fluoride mixing curve.

Hie curve suggests that a deep aquifer source may be responsible for the enrichments. At this point, no satisfactory explanation is available; however, future studies may wish to consider the role of the sub-limestone geology. CHAPTER V

ALKALINITY

The measurement of pH and alkalinity is fundamental to the

chemical analysis of all natural water samples. Along with a

knowledge of the water's buffer capacity, important particularly to

the acidification of these systems, alkalinity is a sensitive

indicator of the dissolution of carbonate minerals and oxidation/

reduction reactions in anoxic systems (Knull and Richards, 1969;

B e r n e r et al., 1970; Einarsson and Stefansson, 1983). Alkalinity is defined as the amount of strong acid required to neutralize a solution to carbon dioxide equivalence end point pH "4.2 (Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981). The principal components that account for alkalinity are the two carbonate species 2— HCO^ and CO^ , as described by the equation;

Alkalinity - -H+ + OH- + HC03“ + 2C032-

(Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981; Morel, 1984).

In natural water systems, weak acids such as borate and silicate may also contribute to the alkalinity.

Alkalinity - -H+ + OH- + HC03“ + 2CX>32_ + B(OH)4~ + H 3Si04~ 117

Under anoxic conditions, organic matter decomposition may produce other weak acids such as sulfide, orthophosphate, ammonium and organic acids, The alkalinity equation then takes on the form:

Alkalinity - -H+ + OH- + HC03“ + 2CC>32- + B(OH)4” + NH3 + HPO^2- +

2P043- + HS“ + 2S2" + H3Si04“ + Organic Acids

Berner (1970), Knull and Richards (1969) and Einarsson and

Stefansson (1983), among others, have investigated the production of excess alkalinity in anoxic marine sediments and anoxic waters and found that the dissolution of metal carbonates along with the production of sulfate reduction by-products such as HjS and NH4+ are of particular importance to titration alkalinity in these systems.

Aerobic respiration, nitrate reduction, iron and manganese reduction and sulfate reduction have all been documented in the

Bermuda groundwater system (Table 2.3; Simmons, 1983; Simmons et al., 1985). The degree to which these reactions occur and their impact on the observed titration alkalinity in these groundwaters have yet to be directly addressed.

The purpose of this chapter is to develop a model similar to that of Berner et al. (1970), Knull and Richards (1969) and Emerson et al. (1983) to determine the major sources of the alkalinity in

Bermuda groundwaters near the Pembroke and Devonshire marsh systems, and to determine whether the observed redox reactions play a significant role in modifying the buffering capacity of the waters and dissolution of the carbonate aquifer. 118

Analysis of water samples collected at JR clearly shows the typical biogeochemical depth profiles for redox potential (Eh), oxygen, nitrate, and ammonium. Although the measurement of Eh in natural water systems is plagued with difficulty, it is used here as a relative measurement of redox conditions in the aquifer. High positive values of redox potential are indicative of oxidizing conditions, and lower, more negative values distinguish the utilization and depletion of available dissolved oxygen. Profiles plotted from water samples collected via bailing show large positive

Eh values (in excess of +300 mv) at the water table and a rapid decrease to less than -80 mv at depth (Fig. 5.1a and b). Pumped profiles are much more consistent, indicating reducing conditions throughout the water column. On May 16 the oxygen profile follows that predicted by Eh, pumped samples being consistently oxygen- depleted, while bailed samples have somewhat higher dissolved oxygen concentration (Pig. 5.2a and b). A decrease in nitrate with depth is also observed at JR for both the March 1 and May 16 (Fig. 5.3a and b). Note that on May 16 the bailed samples display much higher nitrate than do those obtained by pumping. Die lowest value for the bailed is 42 pM, while the highest pump sample is 22 pM.

These data seem to suggest that waters in the borehole may not truly represent water from the aquifer, particularly close to the water table where the water in the borehole is in contact with atmospheric gases. There appears to be an exchange of oxygen between the water standing in the borehole and the atmosphere.

Groundwater pumped from the aquifer is more suboxic and anoxic in character due to limited oxygen flux through the overlying bedrock.

E h 1 1 9 “JR MARCH 1 Eh VS DEPTH* -200 -100 0 100 200 300 400 4 0 0' 200 6 0 0-I via balling large show positive values (In excess of 300 at mv) the profiles Indicate reducing conditions through the water column. Figures 5.1a b. And profiles Eh plotted water samples from col lected watertable and a rapid decrease to less than -80 at mv depth. Pumped DEPTH (cm) Eh Eh (mv) 120 -200 -100 0 100 200 300 400 400- 6 0 0 ■ Figure 5.1b. -JR MAY 16 Eh VS DEPTH’

DEPTH (cm)

OXYGEN (mM) 121 BJR MARCH 1 OXYGEN V3 DEPTH’ ■ 0.00 0.02 0.04 0.06 0.09 200 400 ' 600 ; nave somewhat higher nave dissolved somewhat concentrations. oxygen oxygen The Figures Figures 5.2a and b. profiles Oxygen follow that predicted by Eh* pumped samples being consistentlypumped oxygen depleted while balled samples spike at -450 Is cm unexplained. DEPTH 0 OXYGEN (mM ) • 0.00 0.01 0.02 0.03 0.04 0.05 200 4 0 0' 6 0 0 ; Figure S.2b. “JR MAY 16 OXYGEN VS DEPTH'

DEPTH (cm)

NITRATE (jaM) 123 •JR MARCH 1 NITRATE VS DEPTH* 0 0 50 100 150

0 200 600 400 Figure 5.3a and b. b. and Figure 5.3a nitrate depleted consumed is systems oxygen In by bacterial decomposition of organic matter. organic bydecompositionbacterial of Figure Figure 5.3a. DEPTH (cm) 200 100 NITRATE (yM) 124 •JR MAY 16 NITRATE V3 DEPTH* 200 4 0 0■ 6 0 0 Figure Figure 5.3b.

DEPTH (cm ) 125

in fact, the only direct rapid flux of oxygen to the aquifer groundwaters is probably due to rainfall recharge. An oxygen increase, which at this point is unexplained, is present at a depth of "450 cm below the water table on May 16.

Ammonium concentrations are highest on March 1, with values ranging from 7 to 70 PM bailed and 36 to 56 pM pumped (Fig. 5.4a and b). Ammonium increases with depth due to nitrate reduction and/or nitrogen remineralization from marsh vegetation. On May 16 similar profiles occur, but the range of ammonium values is much narrower and the total concentrations are more than an order of magnitude lower (0.01 to 3.61 pM bailed and 1.8 to 3.5 pM pumped).

The nitrite concentrations are typically low. However, as was observed by Simmons (1983), a characteristic nitrite spike occurs below the surface and signals the presence of nitrogen transfer reactions (Table 6.2). Nitrite is an intermediate of nitrate reduction and also denitrification reactions. Therefore, increased nitrite concentrations may indicate the importance of the nitrogen reactions including its gaseous loss from the system as ^ 0 or Nj

(Fig. 5.5a and b).

Dissolved iron and manganese concentrations range from 4.2 to

32 pg/1 and 2.7 to 33 pg/1, respectively, in pumped and bailed samples collected at JR on March 1 and May 16. Iron concentrations as high as 128 pg/1 and manganese 41 pg/1 are observed at PRH. The higher concentration of iron and manganese (30 to 1500 and 20 to 100 pg/1, respectively) reported by Simmons (1983) at JR were not observed in this study. These lower values may be due to

AMMONIUM (pM) 126 •JR MARCH 1 AMMONIUM V3 DEPTH’ 0 20 40 60 80 400 200 600 H Figures Figures 5.4a concentrations and b. Increase width due to Ammonium DEPTH (cm) nitrate reduction nitrogen and Izatton vegetation. marsh remineral from AMMONIUM (pM) MAY 16 AMMONIUM VS DEPTH" j r - 0 0 1 2 3 4 200 400 • 400 600 • 600 Figure 5.4b. DEPTH (cm)

NITRITE (pM) 128 •JR MARCH 1 NITRITE VS DEPTH’ i 0.0 0.2 0.4 0.6 0.8 1.0 200 400 ■ 600

DEPTH (cm) Figures 5.5a and and denitrification b. Nitrite reactions. Is an Intermediate of nitrate reduction 0.6 0 . 4 NITRITE (pM) 129 0.2 ■ 0.0 4 00 ■ 200 600 ■ Figure 5.5b. -JR MAY 16 NITRITE VS DEPTH*

DEPTH (cm) 130

variability in geochemical or hydrological conditions (Fig. 5.6a, b,

c and d).

Sulfide production at JR is depicted in Figures 5.7a and b.

Sulfide concentrations up to 0.1 mN were observed at the JR well.

Although the sulfide concentrations increase with depth due to

stronger reducing conditions and higher sulfate concentrations, these values are very low and suggest that sulfate reduction is occurring, but is not a major microbially induced process in these groundwaters.

From these data it is clear that microbially induced redox

reactions do occur in groundwaters associated with the marsh similar

to what has been observed in other areas of the Devonshire lense

(Simmons, 1983 and Simoons et al., 1985). The reactions involving oxygen reduction and possibly nitrate reduction appear to be the dominant biogeochemical reactions at least to depths of " 6 m. Iron, manganese and sulfate reduction reactions are of much less significance. These data also indicate that the redox conditions may exhibit a natural variability, both spatially and temporally, in the marsh groundwater, and therefore the concentration of the species ammonia, sulfide and phosphate may also display similar variations in time and space.

Model Development

The modified alkalinity model equation is as follows:

Alkalinity - Ac + -H+ + OH- + 2<*Ca2+ + *Mg2+ + *Sr2+) + HP042- +

2P043- + HS_ + 2S2- + NHj + H3Si04- IRON (ng/1) 'JR MARCH 1 IRON VS DEPTH 600 i 600 400 - 400 200

DEPTH (cm) Figure 5.6a. MANGANESE (pg/l) -JR MARCH 1 MANGANESE VS DEPTH 600 - 600 200 400 Figure Figure 5.6b. DEPTH (cm ) 40 30 (Mg/1) 20 IRON 133 10 OR MAY 16 IRON V3 DEPTH ■ 5 .6 c . 400 • 600 200 Figure DEPTH i'cm) 134 MANGANESE (|ig/l) 10 20 30 40 JR MAY 16 MAN6ANE3E VS DEPTH - 200 400• 600H Figure 5.6d. Figure DEPTH (cm)

135 SULFIDE (mM) SULFIDE "JR MARCH 1 SULFIDE VS DEPTH" 0.00 0.05 0.10 0 200 4 0 0 6 0 0 depth, these values are very low and suggest that sulfate reduction Is Figures 5.7* and b. Although sulfide concentrations Increase with not particularly Important in these groundwaters.

DEPTH (cm) SULFIDE (mM) 136 MAY IB SULFIDE VS DEPTH* aJ R • 200 4 0 0 ■ 6 0 0 ' Figure S.7b.

DEPTH (cm) 137

where H+ and OH- are the concentration of the hydrogen ion and hydroxyl ion. Hie excess calcium, magnesium and strontium (*Ca^+ , 2+ 2+ *Mg and *Sr ) represent the alkalinity derived from the disso­ lution of carbonate minerals, HjSiO^- is introduced via dissolution of silica minerals from the soil zone, the orthophosphate from the degradation of organic matter and/or the dissolution of phosphorus- rich mineral phases, ammonium from nitrate reduction or the degradation of organic matter, and sulfide from sulfate reduction.

It was discussed earlier in this work that Bermuda ground­ waters are a mixture of rain and seawater (Plummer et al., 1976).

Although the alkalinity of the rainwater-derived end member is considered negligible (see Chapter I), that of seawater at 35°/oo is

2.33 meq/1 (Stumm and Morgan, 1981). Alkalinity is a conservative property of oxic surface marine waters, and thus the groundwater in

Bermuda contains a component of alkalinity which has resulted from the conservative mixing of these two end member solutions. This component has been termed conservative alkalinity and is represented in the equation as Ac.

Each of the components of the equation has been either measured directly or determined via a simple mixing model approach

(e.g. conservative alkalinity). The techniques utilized for these measurements are given in Table 3.1 and discussed in Chapter 3. The addition of each of the independent components should give a value equal to that of the analytically determined titration alkalinity.

Deviations from this equality are indicative of species important to the alkalinity which have not been accounted for and/or analytical error. 138

Assumptions

In order to develop the model, several assumptions were made.

They include:

(1) The conservative components of calcium, magnesium, strontium and alkalinity are determined from the freshwater-seawater mixing model (discussed in detail in Chapter IV).

(2) The chemical composition of the freshwater end member is represented by rainwater concentrated 85% by evapotranspiration (Chapter IV; Table 4.1).

(3) The chemical composition of the seawater end member is represented by standard open ocean water published by Stumm and Morgan, 1981 (Table 2.4).

(4) "Excess" borate is considered negligible.

Analyses for borate were not performed. The assumption was made that borate would be conservative in the groundwater, and therefore, from the standard borate seawater concentration of 0.4 mM, the concentration observed in a groundwater of ~10% seawater

(the most saline groundwaters studied) is 0.04 mM. At typical groundwater pHs of ~7, about 0.0004 weq/1 of this total borate would contribute to the titration alkalinity. This accounts for 0.007%, which has been considered to be negligible for modelling purposes.

Equations

In order to determine the components' excess alkalinity, excess Ca2+, excess Mg2+ and excess Sr2+, it is first necessary to determine the amount of alkalinity, Ca2+, Mg2+ and Sr2+ derived from the freshwater-seawater mixing. This is accomplished with the following equation, where X represents the species of interest.

XSW " XFW Conservative X - ( } Clsal[lple + X ^ sw SW - seawater value FW - freshwater value 139

This is similar to the technique used by Berner et al. <1970) and

Einarsson and Stefansson (1983).

The waters at the PBH well do not appear to contain a

measurable seawater component (Chapter IV). If this is the case,

then there would also be no conservative Ca2+, Mg2+, Sr2+ and

alkalinity component from seawater. Therefore, in the alkalinity

calculations, it was assumed that all Ca2+, Mg2+ and Sr2+ over that

present in rainwater has resulted from the dissolution of carbonates

and is determined by:

excess X - xfc -

t - total X

The relative degree of protonation of the weak acids H^SiO^",

HjPO^, H2 S and NH4+ is highly pH-dependent, with tenqperature and

ionic strength also being important (Appendix III). Natural waters usually range in pH between 6.5 and 9.5, and therefore the weak acid 2- 3— species most important to alkalinity include HP0 4 , 2 P0 4 , H3Si0 4,

NH4+ , HS~ and S2~ at various degrees of protonation according to the pH. The analytical methods used in this study have measured total

concentrations of these species. For example, when reference is made to phosphate, we are usually referring to all species _ 5- 3- comprising orthophosphate

(Strickland and Parsons, 1972). In order to calculate the

speciation of the various weak acids, the total concentration (Cfc) of each weak acid is multiplied by its pH-dependent ionization

fraction a so that: 1 4 0

concentration of species - Cfc * ct

For a monoprotic acid:

[H+ 1 a0 - [H+ ] + Ka

al - [H+ ] + Ka

(Snoeyink and Jenkins, 1980) where a0 is the fraction of weak acid undissociated, and al is the fraction of dissociated. Ammonium

(NH4+ ) is considered a monoprotic acid because it will dissociate to

form NHj. For a diprotic acid:

[H+ ]2 a 0 - E

[H+ ]K i a l ------“ E

Kai Ka2 a2 - 31 ^

E - [H+ ]2 + [H+ ]Kal + KalKa 2

(Snoeyink and Jenkins, 1980). H2 COj, H2S and H^SiO^ are examples of diprotic acids. For a triprotic acid, e.g. phosphoric acid (HjPO^) As indicated by the above equations, the concentrations of each weak acid species is dependent solely on the hydrogen ion concentration and respective acidity constant. Thus, given pH, pKs, and total concentration of weak acids, the chemical speciation can be determined. Graphical representations of pH dependence of the four acids of interest here are given in Figure 5.8a, b, c, d, e.

Ionic strength corrections were made with the use of the

Davies equation:

2 1 1 / 2 -LOG a- -Az ( ------0.21) 1 + I1/z

a - activity coefficient I - Ionic Strength activity « o * [X] [X] - concentration of X

(Stumm and Morgan, 1981). The temperatures of all the samples collected in this study were approximately 21°C; therefore. Fraction of Cy pq 4 , Figure 5.8a. 0.5 dependent solely on the pH and pKa. This is described by the is described by This the solely on pKa. pH and dependent iue58 oe Theis speciation the of each of acids weak 5.8a e.to Figure following graphical'following representations(Snoeyink 1980).and Jenkins, 0 “ pK., 6 . 8 0 2 14 j . K p 12 10 8 ,.j K p 6 4 .,, K p 2 21 . =12.3 2 1 = 7.2 “ u 2.1 = PH 142 “i “a =

Fraction ofCj. a Fraction of CT • q P o in o o (71 O iue 5.8c Figure

T? z

N

n r» Fraction of CT a Fraction ofC t,s »v P - o — O U1 o o w o Figure 5.8e. <£) C "I 0 VI • CD

X ■f*

0 145

temperature effects on the acidity constants Ka (25°C) were ignored

(Appendix ill).

The above equations were subsequently programmed in Applesoft

BASIC for the Apple lie computer (Appendix I I I ) . The model alkalinities were then calculated for both the JR and PRH wells, and the results are given in Table 5.1.

Results

Deviations of the model alkalinities from that of the analytical alkalinities range from 9.6% overestimation to 22.2% underestimation, the largest deviations being observed at JR on May

16. The reasons for the observed deviations may result from one or all of the following possibilities. The model is dependent on the summation of eight different components, each of which were independently measured by several different techniques. Each of the analyses was subject to analytical errors (Table 3.1) which would have been compounded in the model calculations. This combined error, which may account for either overestimated or underestimated alkalinity deviations, has been determined by the equation:

X - (( A )2 + ( B )2 + ... )1 / 2 where A, B, ... are the standard deviations of each of the analyses, and X is the combined standard deviation error. For the model alkalinity equation, this combined error accounts for only

“+0.02 meq/1, which is ~0.5% precision.

The possibility that dissolved humic substances may contribute to the titration alkalinity was not considered by the model. 146

Table 5.1a. Alkalinity model.

PRH - February 11, 1985 Units in mM

Excess Excess Excess Depth Ac 2+ 2+ 2+ ,2- (cm) Ca Mg Sr HS

Surface 0 2.732 0.149 0.014 2 xlO”U 1 .2 2 x1 0r'4 -4 152.4 0 2.750 0.149 0.014 3xl0- 1 1 1.28x10'

304.8 0 2.882 0.162 0.014 2 xl0 - 1 1 1 .2x1 0 ”'

396.2 0 2.750 0.157 0.015 3xl0— 11 1.3x10”'

Depth (cm) h p o 4” nh3 H3 Si04 pH P°43"

Surface 6 .8 xl0 ” 4 9xl0” 9 3.9xl0" 4 7.120

152.4 8 .3xl0~ 4 1 .3xl0- 8 3.4xl0“ 4 7.171

304.8 6.5xl0” 4 9x10” 9 3.4xl0“ 4 7.116

396.2 5.3xl0~ 4 8 x1 O-9 3.1xl0” 4 7.179

Ana­ Contri­ Depth lytical Model Differ­ % Dif­ bution (cm) Aik Aik ence ference S N P !

Surface 6.31 5.79 0.52 8 . 2 lxlO” 3

152.4 6.28 5.83 0.45 7.1 1.3x10'

304.8 6.25 6.12 0.13 2.1 lxlO” 3

396.2 6.31 5.85 0.46 7.4 lxlO” 9 147

Table 5.1b. Alkalinity model.

JR - March 1, 1985 Units in mM

Excess Excess Excess Depth Ac (cm) Ca2+ Mg2+ Sr2+ S2- HS

Surface 0.037 2.424 0.439 0.012 l.lxlO- 8 0.024

152.4 0.117 2.813 0.319 0.017 1.23xl0-8 0.030

304.8 0.133 2.962 0.373 0.019 1.73xl0- 8 0.042

457.2 0.151 3.063 0.545 0.0188 2.56xl0-8 0.048

Depth (cm) hpo4 NH 3 H3 Si0 4 P°43" PH

Surface 4.18x10" 4 1.47xl0-8 4.48xl0-4 7.392

152.4 6.38x10~ 4 2.78xl0~8 4.72xl0-4 7.271

304.8 4.69x10“ 4 2.15xl0-8 4.97xl0' 4 7.257

457.2 5.80x10~ 4 3.6 6 xl0 -8 5.91xl0-4 7.357

Ana­ Contri­ Depth lytical Model Differ­ % Dif­ bution (cm) Aik Aik ence ference S N P Si

Surface 5.51 0.93 13.6 0.025

152.4 7.18 6.45 0.73 1 0 . 2 0.031

304.8 7.31 6 . 8 8 0.43 5.8 0.044

457.2 7.48 7.45 0 .02 0.2 0.049 148

Table 5.1c. Alkalinity model.

JR - May 16, 1985 Units in mM

Excess Excess Excess Depth Ac (cm) ca2+ Mg2+ Sr2+ s2- HS-

Surface 0.048 2.139 0.506 0 . 0 1 0 3.3xl0- 9 0.013

152.4 0.207 3.304 0.777 0.016 l.lxlO- 8 0.042

304.8 0.075 2.671 0.369 0 . 0 1 1 8xl0- 9 0.036

457.2 0.209 3.280 0.794 0.014 1 .2xl0- 8 0.043

609.6 0.244 3.425 0.126 0.016 lxlO- 8 0.041

Depth (cm) h p o 4" po43- nh3 H3Si0 4 pH

Surface 9.47xl0- 4 1.9xl0- 8 1.137xl0-5 2.9xl0- 5 7.108

152.4 9.05xl0-4 3x10- 8 1.64xl0-5 1.57xl0- 5 7.028

304.8 9.1xl0- 4 1.9xl0- 8 1.26xl0-5 2 .6 xl0 - 5 7.040

457.2 1.5xl0~ 3 6.1xl0- 8 1.76xl0-5 7.5xl0- 6 7.064

609.6 2.09xl0-3 7.5xl0- 8 1.5xl0- 5 1 .6xl0 - 5 6.974

Ana­ Contri­ Depth lytical Model Differ- % Dif- bution (cm) Aik Alk ence ference! S N P Si

Surface 6.91 5.37 1.54 22.2 0.014

152.4 7.87 8.44 -0.57 -7.3 0.043

304.8 7.10 6.21 0.88 12.4 0.037

457.2 7.69 8.43 -0.74 -9.6 0.044

609.6 7.76 7.42 0.34 4.3 0.043 149

However, since the JR and PRH groundwaters are very much a part of a marsh/bog type system, a humic-organic acid contribution to the alkalinity cannot be considered negligible. Host groundwater samples, from JR in particular, have a yellow-brown color; this is a clear indication of the presence of humic substances (Eshleman and

Hemond, 1985; McKnight et al., 1985). Organic acids in soil leachates, and marsh and bog waters have been well documented

(Eshleman and Hemond, 1985; McKnight et al., 1985; Cronan et al.,

1978) and have been shown to account in some cases for 13 to 21% of the total anions and strong acids. Typical values for anion deficits can range between 0.1 and 0.5 meg/1 and are believed to represent organic acids (McKnight et al., 1985; Cronan et al.,

1978). Not including the organic acids would cause an under­ estimation in the model alkalinity.

in the calculation of the conservative components of Ca2+,

Mg2+, Sr2+ and alkalinity, the assumption was made that the freshwater end member has the composition calculated in Chapter IV

(Table 4.1), the seawater end member has the average composition of standard open ocean seawater reported by Stumm and Morgan (1981).

If the actual composition of either of these waters varies even slightly from these data that have been used, a positive or negative error would be reflected in the calculation of the conservative calcium, magnesium, strontium and alkalinity in the model alkalinity. In all probability, the data for the seawater end member is a good approximation of Bermuda coastal waters, and there is confidence that this potential error contribution is minimal.

The freshwater data, however, are expected to show much more 150

variability due to a) seasonal fluctuation in evapotranspiration, b)

coastal sea spray effects, c) ion exchange in the soil, and d)

anthropogenic effects. In any case, the values that have been used

here represent the best available estimates.

In the well mixed waters of PRH, the average alkalinity over

the 4m depth interval is 6.29 +0.03 meq/1. Hie contribution to

alkalinity from sulfide, ammonium, phosphate, and silicate appears minute, the combined effects of these species in each case being

less than 0.02% of the total titration alkalinity. This suggests

that the dissolution of metal carbonate minerals is responsible for virtually all of the observed excess alkalinity.

In none of the waters studied did the combined contribution of

sulfide, ammonia, phosphate and silicate account for greater than

0.7% of the observed alkalinity. Of this minor contribution,

sulfide is the dominant influence (Table 5.1). It is concluded from

this that the dissolution of carbonates, represented primarily by

the excess Ca^+ , is responsible for almost all of the observed excess alkalinity at the JR and PRH wells.

Saturation

Thermodynamic equilibrium calculations were conducted on the

chemical data to determine the chemical speciation and the degree of

saturation with respect to the carbonate minerals calcite and aragonite in the groundwater. The speciation calculations were performed by the thermodynamic model, MINEQL (Westall et al., 1976;

Appendix I).

The pH and other chemical data representing water samples

collected in both the Pembroke and Devonshire marshes (PRH and JR) 151

2+ 2- were entered into the program, and free Ion Ca and C03

concentrations were determined (Table 5.2). With this information

the saturation indexes (SI) for each water sample considered was

calculated for calcite and aragonite. The saturation index is described by the equation:

SI - log

where IAP is the product of activity of the free calcium Ca2+ and 2_ carbonate C0 3 ions, and the Kgo is the solubility product of the —A 4 8 mineral phase of interest (in this case, K - 1 0 ' for calcite SO a 336 and 10~ ’ for aragonite (Plummer and Busenberg, 1961)). For each water sample considered, the solubility constants (KfiQ) were corrected for the ionic strength effect using the Davies equation given earlier in this chapter (Stumm and Morgan, 1981).

CaC03 <---- > Ca2+ + C032-

{Ca2+} {C032“ } - Kso

o[Ca2+] a[C032-] - Kgo

[Ca2+] [ C O 2-] - ^ J era

(X) - activity of component X

A positive value for SI indicates that the water sample is supersaturated with a particular solid, a negative value indicates undersaturation, and SI equal to zero represents equilibrium between 152

solid and solution phase. One limitation to this lies in that the

solubility products (K ) represent the pure calcite and aragonite SO mineral phases when in actual fact the limestones are a mixture of

high and low magnesium calcite and aragonite.

Carbonate saturation in Bermuda marshes has been discussed by

Plummer et al. (1976) and reviewed in Chapter II of this work.

These investigations found calcite and aragonite to be under­

saturated in the groundwaters' surrounding marsh. They showed that

high PCO2 generated by reactions in the marsh are responsible for

decreasing the pH and creating the under saturated conditions.

The results in this study indicate the groundwaters to be at

approximate equilibrium with calcite and slightly undersaturated

with respect to aragonite (Table 5.2). The PC02 has been calculated

from pH and alkalinity via carbonate equilibria (Table 5.3) for each

sample and are comparable to values reported by Plummer et al.

(1976). The pH, alkalinity and calculated PC02 are given in Table

5.4 and have been plotted in Figures 5.9 to 5.11.

In the well at PRH, it was established that the groundwaters were mixed down to a 4m depth, and therefore no distinct pH and

alkalinity variations are observed. What this seems to suggest is

that much of the carbonate dissolution in this well may have

occurred in the vadose zone (Table 5.1, Fig. 5.9a, b and c). JR,

however, displays distinct variations. In general, for both pumped and bailed samples, pH tends to decrease with depth and alkalinity

tends to increase. According to the model discussed previously,

this alkalinity increase is not due to the production of the minor

ions sulfide, phosphate, ammonium, etc., but rather due to 153

Table 5.2. Saturation indices for calcite and aragonite in marsh groundwaters. I - ionic strength.

Free Free Corrected Corrected , 2_ log Kso SI log Kso SI well I Ca C03 calcite calcite aragonite arag.

O £ Surface 0.02 1.37x10 J 7.31x10 0 -7.999 -0.0005 -7.855 -0.14

_£ 457 cm 0.07 2.39x10 J 8.6x10 0 -7.698 0.01 -7.554 -0.13

n _£ 609 cm 0.08 2.76xl0-J 8.14x10 0 -7.661 0.01 -7.516 -0.13

PRH 0.008 1.34xl0- 3 5.19xl0-6 -8.156 -0.002 -8.012 -0.14

TCD 0.01 1.76xl0-3 4.31xl0-6 -8.122 -0.002 -7.978 -0.14

1WL 0.62 l.lOxlO’ 2 3.74xl0- 6 -7.212 0.17 -7.068 -0.32 154

Table 5.3. Carbonate equilibria. Equilibrium constants have again been corrected for ionic strength using the Davies equation.

C02 + h20 < > w FKh “ 1.5

<----> H+ + h c o 3“ 6.3 H2°°3 PRa l - h c o 3" <----> H+ + C032- PRa2 - 10.3 caco3 <----> Ca2+ + C032- PRso - 8.34 155

Table 5.4.

PRH, February 11, pump

Depth (cm) pH Aik p c o 2

0 7.120 6.312 -1.48

152.4 7.171 6.276 -1.53

304.8 7.116 6.251 -1.48

396.2 7.179 6.311 -1.54

JR, May 16, bail

0 7.150 5.378 -1.57

152.4 7.017 7.313 -1.29

304.8 6.882 7.944 - 1.12

457.2 6.927 7.874

609.6 6.938 7.356 -1.51

JR, May 16, pump

0 7.108 6.906 -1.41

152.4 7.040 7.098 -1.32

304.8 7.028 7.871 -1.23

457.2 7.064 7.686 -1.28

609.6 6.974 7.764 -1.18

JR, March 1, pump

0 7.392 6.748 -1.71

152.4 7.271 7.178 -1.53

304.8 7.257 7.314 -1.51

457.2 7.357 7.475 -1.59

609.6 7.317 7.482 -1.55

0 0 9 pH 156 - 6.8 6.9 7.0 7.1 7.2 7.3 7.4 - jre 5.9a. "PRH FEB 11 PH VS DEPTH" 1 0 0 200 500* 60 0 H 300 - 4 0 0 - Figure 5.9 to 5.11. The al kal Inlty Increase with depth appears not to be be due to sulfide* phosphate or but ammonia* rather due to dissolution. accumulation* decrease pH and the subsequent calclun carbonate DEPTH (cm) ALK (meq/l) 5.9b. "PR H FEB 11 ALKALINITY VS DEPTH" ■ - 100 200 3 0 0 * 5 0 0 - 4 0 0 - 6 0 0 H Figure

DEPTH (cm) 0.08 P C 0 2 (aim) 158 0.02 0.03 0.04 0.05 0.06 0.07 0 100 200 3 0 0 500 600 4 0 0

Figure 5.9c. "PRH FEB 11 CALCULATED P C 02 VS DEPTH" DEPTH (cm) pH MARCH 1 pH VS DEPTH" "JR - - 6.9 7.0 7.1 7.2 7.3 7.4 7.5 gure 5.10a. 1 0 0 2 0 0 3 0 0 - 5 0 0 - 600 H 4 0 0 - F1 F1

DEPTH (cm) ALK (meq/l) ALK - - 1 0 0 2 0 0 3 0 0 - 4 0 0 - 5 0 0 - 6 0 0 H Figure 5.10b. "JR MARCH 1 ALKALINITY VS DEPTH"

DEPTH (cm) 0.07

& ___ 0 I . 0

« 0.05 - - P C 0 2(atm) 161 - ■ 0.01 0.02 0.03 0.04 100 600 H 500- 300- 2 0 0 400- 5.10c. .JR MARCH 1 CALCULATED PC02 VS DEPTH"

Figure DEPTH (on) pH 162 MAY MAY 16 pH VS DEPTH” "JR 6.9 7.0 7.1 7.2 7.3 7.4 6 .8 - Figure 5.11a. 200 4 0 0 -

DEPTH (cm) ALK (meq/l) ALK 163 16 ALKALINITY VS DEPTH” m a y

j r ” * 200 4 0 0 - 6 0 0 H Figure 5.11b.

DEPTH (cm) PC02 (atm) PC02 1 6 4 - 0.02 0.03 0.04 0.05 0.06 0.07 0.08 200 60 0 H 4 0 0 - Figure 5.11c. "JR 16MAY CALCULATED PCQ2 VS DEPTH" DEPTH (cm) 165

dissolution of the limestone aquifer material in both the vadose and phreatic zones.

The results discussed here support the findings of Plummer et al. (1976) and show that PCO2 concentrations in the groundwater column vary vertically in response to biogeochemical oxygen reduction and probably nitrate reduction reactions.

Aquifer Geochemistry

Rock samples from both wells were found to contain between 11 and 20 mmoles/gm (400 and 800 mg/gm) calcium, 0.008 to 0.8 mooles/gm

(0.19 to 190 mg/gm) magnesium, and 0.02 to 0.06 mmoles/gm (2 to 5 mg/gm) strontium. The magnesium/calcium ratios were found to be lower in the rock material (0.005 to 0.007 molar ratio) than in the water in the wells (0.11 PRH and 0.33 to 1 JR). Similar results were observed with the strontium/calcium ratios which are fairly constant at 0.005 in the waters at PRH and 0.005 to 0.007 at JR.

The strontium/calcium in the rock ranged from 0.002 to 0.003 and

0.003 to 0.005 at PRH and JR, respectively, with slightly higher values at JR.

In addition to the chemical analyses, X-ray diffraction analyses were conducted on the pulverized rock samples. The plots are shown in Appendix II. They clearly indicate that calcite is the dominant carbonate crystalline structure present at both locations.

Aragonite is present in JR limestone, but no aragonite can be identified at PRH. The X-ray diffraction technique used in this study was not quantitative, and therefore the actual concentrations of the respective minerals are not known. Semiquantitative calculations using the relative peak heights of calcite versus 166

Table 5.5. Composition of limestone.

mmoles per 2 ?

PRH

Depth (cm) Calcium Magnesium Strontium

0 11.76 0.0819 0.024

200 21.71 0.312 0.054

400 17.50 0.197 0.047

JR

0 13.42 0.111 0.068

300 15.46 0.185 0.043

600 20.16 0.226 0.064

Ratios

PRH

Depth (cm) Mg2 +/Ca2+ Sr2 +/Ca2+

0 0.007 0.0020

200 0.014 0.0025

400 0.011 0.0027

JR

0 0.008 0.005

300 0.012 0.003

600 0.011 0.003 167

aragonite suggest that aragonite accounts £or approximately 1% of

the JR limestone samples. This agrees with the slightly higher

strontium in the aquifer material at this well location.

The high magnesiums/calcium and strontium/calcium ratios

indicate that incongruent dissolution accounts for the diagenetic alteration of aragonite and high magnesium calcite to low magnesium calcite and calcite at these well sites. Because high magnesium calcite and aragonite contain significant quantities of magnesium and strontium, respectively, these elements are put into solution to a larger proportion than calcium. Therefore, the solution chemistry differs in that the magnesiums/calcium and strontium/calcium ratios are higher than the minerals from which they were derived.

There is a distinctly higher magnesium/calcium ratio at JR compared to that found at PRH, but this is not believed to be the

result of the magnesium released during dissolution. As indicated earlier, the JR well is subject to mixing with seawater, while the

PRH well appears not to be. The higher magnesium from seawater accounts for higher magnesium/calcium observed at JR over PRH. The fact that the limestones at JR contain aragonite and those at PRH do not shows that less carbonate alteration has occurred at JR and that these limestones are probably younger than the limestones at

PRH. Thus, the JR limestones are more "Paget" in character, while the PRH limestones appear more "Belmont." CHAPTER V I

THE FATE OF INORGANIC NITROGEN AND PHOSPHORUS IN BERMUDA GROUNDWATER

Recently, much attention has been given to anomalously high

inorganic nitrogen concentrations in Bermuda groundwaters. Simmons

(1983) and Thomson and Foster (1986) have shown that the practice of sewage disposal via cesspits on the island has lead to nitrate levels in excess of 2 mM in densely populated regions of the island.

On the other hand, phosphate concentrations are very low with maximum values of approximately 4.0 pM and most values less than 1

UM (Simmons, 1983).

Cesspits are unlined subsurface sewage containments specifically designed for maximum dispersion of the sewage to the subsurface. They are the primary means of sewage disposal on the island, and there are concerns that the discharge of the resulting nitrogen- and phosphorus-contaminated groundwaters into the inshore waters may be responsible for the periodic eutrophic processes observed in many harbors and bays of Bermuda in the past decade

(Morris et al., 1977; Barnes and Bodungen, 1978).

lhe nitrogen-to-phosphorus ratio (N/P) of domestic sewage is approximately 9:1 (Murray et al., 1981), yet the values reported by

Simmons (1983) in cesspit-contaminated groundwaters of Bermuda are as great as 1500:1. Although these findings clearly indicate selective phosphorus removal relative to nitrogen, fixed nitrogen itself is subject to losses resulting from denitrification. 169

The geochemical processes which control phosphorus concen­

trations in Bermuda groundwaters are of great interest because

phosphorus is believed to be the limiting nutrient for algal growth

in the inshore waters (Morris et al., 1977; Barnes and Bodungen,

1978). Convincing evidence has yet to be provided that groundwaters

in Bermuda provide a major source of the nutrients nitrogen and

phosphorus to the nearshore waters. Studies of Ghyben-Herzberg type

systems along the east coast of Florida have clearly demonstrated

brackish groundwater fluxes into the ocean on the order of 45 m 3/day

from a strip one meter wide (Kohout, 1960). It can be assumed that

groundwaters are discharging in a similar fashion along the Bermuda

coastline and that these waters contain dissolved nitrogen and phosphorus compounds geochemically similar to the actual groundwaters (D'Elia et al., 1981). Therefore, an understanding of

the mechanisms which are responsible for controlling nitrogen and phosphorus levels may be of great inportance to understanding the potential for eutrophication in Bermuda's inshore waters.

Inorganic Phosphorus

In most natural water systems, inorganic phosphorus is present as orthophosphate. Orthophosphate is a weak acid, and its speciation is therefore highly pH-dependent (Table 6.1; Fig. 6.1).

Phosphate is known for its adsorption onto oxides of iron

(II), iron (III) and aluminum (III). It also forms authigenic minerals with calcium, iron and aluminum under low temperature and pressure conditions (Stumm and Morgan, 1981; Berner, 1981; Drever,

1982). In carbonate rocks and sediments, phosphate.is known to 170

Table 6.1. Speciation of orthophosphate (Stum and Morgan, 1981; Snoeyink and Jenkins, 1980).

H3P04 <--- > H+ + H2 P04“ pKal - 2.15

H2P04_ <-- > H+ + HP042- pKa2 - 7.20

HP042_ <-- > H+ + P043- pKa3 - 12.35 171

1.0 o i u*- O 0.5 -

12.3

Figure 6.1. Orthophosphate speciation. At pHs < "2 undissociated phosphoric acid (H.PO.) is the dominant species. Most natural_ waters range in pH between 6 and 9, and thus the species H^PO,- and HPO. are dominant. At pHs > ~9, orthophosphate occurs primarily as PO. ” (Snoeyink and Jenkins, 1980). 172

ceact and co-precipitate with calcium carbonate to form calcium

car bona te-phosphate surface complexes and/or the mineral apatite.

In addition, phosphate can also be adsorbed by various organic

compounds (Kitano et al., 1977; Berner, 1981; Avnimich, 1980;

Gaudette and Lyons, 1980).

The two primary apatite minerals are fluorapatite (Ca5 (P04 )3F) and hydroxyapatite (Cag(P0 4 )30H), with the fluoride-rich form being most predominant in the marine environment. Phosphate produced from

the biochemical decomposition of organic matter can react with

carbonate to form the apatite minerals. This is thought to be a

replacement reaction (Ames, 1959; Simpson, 1966; Ishikawa and

Ichikuni, 1981; Jahnke and Emerson, 1982). in both oxic and anoxic sulfidic sediments, apatite is thought to be the dominant mineral

sink of phosphate. However, in anoxic sediment systems where sulfate is not present in any significant degree and the sulfide has been completely removed as metal sulfide minerals, iron (II) concentration may become high enough that eventually vivianite

(Fe-jtPO^j'SHjO) becomes supersaturated and precipitates (Berner,

1981).

Previous studies in Bermuda have suggested that phosphate

removal may be due to adsorption by the carbonate aquifer material, apatite formation (Kessaram, 1979; Simmons, 1983) or the formation of iron (III) and aluminum (III) phosphate minerals (FeP04 strengite, AlP04 variscite) (Simmons 1983; Jickells et al., in preparation). 173

Inorganic Nitrogen

Nitrogen is the most abundant atmospheric gas. It is present as the molecular dinitrogen (Nj) and accounts for approximately 79% of the composition of the atmosphere (Raiswell et al., 1980).

Dinitrogen is not an inert gas, but due to a strong triple bond between the two atoms it behaves as an inert gas, being involved to a very limited extent in chemical reactions (Haite, 1984). It is highly soluble in water and comprises 48% of the dissolved gases in the surface ocean (Gross, 1982).

Nitrogen compounds such as nitrate and ammonium along with orthophosphate are extremely important in the production of proteins and amino acids. When supplies of nutrient nitrogen are limited, biological processes, are severely affected (Fenchel and Blackburn,

1979). Hie forms of nitrogen available to most organisms for biological functions are ammonium (NHj+ ), nitrite (NO^- ) and nitrate

(NOj- ), with ammonium being the most readily used (Fenchel and

Blackburn, 1979). In order for nitrite and nitrate to be used in biological processes they must first undergo assimilatory reduction to ammonium (Fenchel and Blackburn, 1979).

Few organisms are able to reduce and use molecular dinitrogen for their needs. Organisms which are able to perform this include some species of blue-green algae, and they provide a critical link between gaseous dinitrogen and biologically available fixed nitrogen. This dinitrogen reduction to ammonium is called nitrogen fixation, and it initiates the nitrogen cycle (Raiswell et al.,

1980; Waite, 1984; Fig. 6.2; Table 6.2). 174

Figure 6.2. Nitrogen Cycle for the Bermuda Groundwater System

Recharge

Nitrogen Fixation 'ON

Cesspit

water table

Inorqanic Organic Denitrification Nitrogen Nitrogen / ON IN

Inshore Waters 175

Table 6.2. Nitrogen transfer reactions (Raiswell et al., 1980),

Nj + 3H2 > 2NHj nitrogen fixation

NH4+ ---> NH2OT > N02" > N03- nitrification

NOj- ---> N0 2~ > NH4+ nitrate reduction

N0 3 > N0 2 > N20 > N2 denitrification 176

Groundwater fixed nitrogen data are available for many of the highly populated areas of the Devonshire lense, and for comparison, samples have been collected near the sparsely populated Devonshire marsh. The marsh was expected to yield waters uncontaminted by cesspit effluent, but because of nutrient cycling between the marsh waters and vegetation, probably containing the maximum fixed nitrogen and phosphorus that would occur naturally in Bermuda groundwaters.

Groundwater samples were collected at JR and analyzed by techniques described in Chapter III. The mixing chemistry of the waters has been discussed previously, and was shown to be a mixture of rainwater concentrated by evapotranspiration and seawater. The maximum chloride concentration at a depth of about 6 m has been determined to be 60 mM or approximately 1 0 % seawater.

The average phosphate concentration at JR is 1.5 PM (range 0.5 to 3.8 PH), and the source of phosphate is most probably the decomposing marsh vegetation. In general, phosphate concentrations associated with the marshes are somewhat higher than that observed at most non-marsh locations as noted by Simmons (1983). Hie profiles of phosphate versus depth suggest that phosphate concentrations may increase with depth, with the highest values being associated with the highest chlorides concentration (Fig. 6.3,

6.4, 6.5). This phosphate is undoubtedly produced by suboxic and anoxic bacterially mediated organic matter degradation discussed above.

For the samples collected by bailing, the highest EN concen­ trations (where EN - NOj” + NC^- + NH4+ ) were observed at the water 177

Figures 6.3 to 6.5. Data suggest that phosphate concentrations increase with depth, highest phosphates being associated with highest chlorides. PHOSPHATE (>iM) 17 8 0.0 0.5 1.0 1.5 • 200 4 0 0• 6 0 0 1

DEPTH Figure 6.3. -JR MARCH i PHOSPHATE VS DEPTH* PHOSPHATE (pM) 179 - 200 4 0 0- 6 0 0 Figure 6.4. -jR MAY 16 PHOSPHATE VS DEPTH' DEPTH (c m ) PHOSPHATE (pM ) A 3 2 0 0 iue 5. .5 6 Figure 10 HOIE (mM) CHLORIDE 180 4030 50 60 181

surface on both March 1 (136 yM) and May 16 (123 yM). The IN from

the pumped samples were lower than those collected by balling on

both dates. As discussed earlier, bailed samples represent

groundwater in the borehole while pumped samples are more

representative of groundwater from the aquifer. The IN samples

collected by pumping also seem to show less variation with depth;

however, overall variations between March 1 and May 16 are more

pronounced. This suggests that temporal variations may be of some

significance.

The IN/P ratios range between 4 and 27, with one surface value

at JR on May 16 yielding a ratio of 260 (Fig. 6 .6 a and b). The IN/?

ratios of many areas covered by vegetation are on the order of 15:1

(Likens et al., 1977), and one may expect the marsh IN to display

similar values. Unlike the high IN/P values observed by Simmons

(1983) for more widespread sampling in non-marsh areas, these observed values are comparable to the 15:1 ratios.

The results of Kessaram (1979) and Simmons (1983) have

suggested that phosphate removal is related to carbonate geochemistry through the formation of calcium carbonate-phosphate

complexes and/or the mineral apatite. Although phosphate removal in these marsh-influenced wells is not as evident from thelN/P ratios, this does not mean that phosphate is not affected by similar processes since the same basic carbonate geology exist island-wide.

The higher phosphate and lower IN/P can be explained as follows.

Many of the wells sampled by Simmons (1983) were at higher elevations, and the initially phosphate-rich cesspit effluent had to seep through several meters of vadose carbonate, allowing greater 300 182 S U M N / P ■JR MARCH 1 SUM NITROGEN/PHOSPHATE' 600 200 400' Figure Figure 6.6a. n DEPTH (■ 3 300 200 SU M N /P 183 100 ■JR MAY 16 SUM NITROGEN/PHOSPHATE' 600 200 400 Figure 6.6b. Figure DEPTH (cm ) 184

exposure to carbonate surface before reaching the water table. This

separation of the groundwater from the source allows for more

efficient phosphate removal by the carbonates. In contrast, much of

the degradation of vegetation in the marsh is occurring essentially

in the groundwater environs. If the rate of phosphate removal is

less than its rate of release by the decay of vegetation, the

observed phosphate concentration in the marsh would be higher than phosphate concentrations observed in non-marsh areas. Another possible explanation for the observed E N/P ratios may involve rapid

fixed nitrogen losses via denitrification. This would also effectively lower the ratio. Denitrification will be addressed

subsequently.

The regulation of phosphate concentrations in the groundwaters by the precipitation of the authigenic minerals hydroxylapatite and fluorapatite has been investigated using thermodynamic/equilibrium computer calculations. The program MINEQL (Appendix I; Westall et al., 1976) was used to determine calcium, phosphate and fluoride 2+ 3- — speciation and concentrations of the free Ca , PO^ and F . With these, the appropriate chemical formulas and the most reliable thermodynamic solubility constants, the saturation state of the phosphate minerals hydroxylapiatite and fluorap>atite were determined.

Ca5(P0 4 )3(ffl 57.8 (Krauskopf, 1979)

Ca5 (P04)3F 60.4 (Krauslopf, 1979)

The saturation indices for these minerals were computed indepen­ dently of MINEQL, and for calculation purposes the chemical formulas 185

were used in the form:

Ca2.5

Ca2.5(PO4>1.5F0.S pKso " 3 0 *2

The Davies equation (Chapter V) was used to correct for ionic

strength effects on the solubility constants. The ion activity

products (IAP), the corrected solubility constants and the

saturation indices (SI), all discussed in Chapter V, are given in

Table 6.3. The saturation indices (SI) show that in all cases the

two apatites are under saturated and therefore should not be

responsible for phosphate removal. This result has been confirmed by X-ray diffraction (XRD) analyses of limestone samples from inside

the wells at JR and PRH (Appendix II). These analyses show no

evidence that either hydroxylapatite or fluorapatite are present in

the aquifer material.

The possibility of vivianite (Fe3 (P04 )2 '8H20) was also considered.

(Fe3 (P04 )2 ’8H20) pKso - 33.30

Dissolved iron measured in this study is taken to be predominantly

Fe2+. According to the calculations, vivianite was also found to be under saturated, and thus not an important sink for phosphate. The concentration of dissolved free Fe3+ species is believed to be very low in the carbonate material, although this may not be the case in the soils. If the removal of phosphate as Fe3+ (strengite FePO^) or 186

Table 6.3. Saturation indices £or fluorapatite and hydroxylapatite (free ion concentrations in moles/liter).

2+ - 3- Well I Free Ca Free F Free FO^ Free OH

JR 0 cm 0.02 1.37xl0-3 1 .20xl0-5 1.14xl0-16 1.29xl0”7

JR 457 cm 0.07 2.39xl0-3 1.44xl0” 4 1 .86xl0-16 1.15xl0-7

JR 609 cm 0.08 2.76xl0'3 1.61xl0-4 1 .86xl0-16 9.33x10”®

PRH 0.008 1.34xl0-3 4.llxl0” 6 5.84xl0-17 1.32xl0-7

TCD 0.01 1.76xl0"3 1.24xl0-5 4.83xl0-17 9.77x10”®

TWL 0.62 l.lxlO-2 1.42xl0”4 1.16xl0-16 1.15x10”7

LAP Corrected IAP Corrected well OH apatite F apatite lo9 koh SIOH log k f SIF

JR 0 cm 3.03xl0”35 -27.5 -7.03 2.93xl0-34 -28.8 -4.77

JR 457 cm 2.40xl0-34 -26.6 -7.01 8.50xl0”33 27.8 -4.21

JR 609 cm 3.10x10” 34 -26.5 -7.01 1.39xl0” 32 -27.7 -4.14

PRH 1.06xl0-35 -27.9 -7.02 5.95xl0-35 -29.2 -4.99

TCD 1.36xl0-35 -27.8 -7.01 1.54xl0-34 -29.1 -4.68

1WL 5.37xl0-33 -25.2 -7.09 1.89xlO”31 -26.4 -4.32 187

Al3+ (variscite AlPO^) or the scavenging processes by clay minerals were of any significance, they would occur in Bermuda's red, iron- and aluminum-rich soils (Stumm and Morgan, 1981). This has been considered recently by Jickells et al. (in press) for Bermuda groundwaters.

In light of these results, the observed phosphate removal must be due to adsorption onto limestone surfaces and/or the formation of calcium carbonate-phosphate surface complexes (Kitano et al., 1978;

DeKanel, 1978; Kessaram, 1979; Avnimelech, 1980; 1983; Ishikawa and

Ichikuni, 1981).

Inorganic nitrogen compounds are far more mobile than phosphate; however, the species ammonium is prone to adsorption and ion exchange, in order to address the fate of fixed nitrogen in

Bermuda groundwaters, two steady state assumptions have been made to define and simplify the system; they are as follows:

1) The annual rainfall recharge is equal to the annual groundwater discharge to the inshore waters (~2 i cn/yr). Plummer et al. (1976) have indicated, however, that recharge via the non-marsh soil zone can be up to 20 times that of marshes due to cracks in the limestone, particularly in the Belmont and Walsingham formations.

2) The total inorganic nitrogen (IN) in groundwater does not change significantly on an annual basis.

The total inorganic nitrogen (IN) is the sum of nitrate + nitrite + ammonium. The data used in the calculations of IN were determined by analytical procedures described above. Average recharge was calculated earlier to be 21 cnyyr from major ion data, annual rainfall (140 cm) and rainfall IN chemistry were obtained from Macky (1957) and Church et al. (1982) respectively (Table 6.4). 188

Table 6.4. Groundwater nutrient box-model data.

ANNUAL RAINFALL 140 cm Mackey, 1957

EVAPOTRANS. 85% This report

RECHARGE 21 cm This report

IN IN RAIN 11.11 M Church etal., 1982 1 8 9

With this, the average amount of inorganic nitrogen (Zn ) reaching the groundwater annually via rainfall is calculated to be 2 on the order of 15 nmoles/m . This yields an initial IN concentra­ tion at the water table of approximately 74 um (0.074 raM) after 85% evapotranspiration. The average N concentration for this upper 6 m of groundwater is 65 UM (0.065 mM), with values ranging from 10 to

130 pM. The fact that evapotranspi ration of rainwater yields a groundwater concentration of ~74 pM suggests that this average of 65 pM is a good representation of unpolluted groundwater in Bermuda.

If the annual discharge of groundwaters into the inshore 2 waters is 21 cm/tat , then the annual nitrogen loss is 13.6 2 mmoles/m *yr. By subtraction, the average annual nitrogen loss to 2 denitrification is 1.4 mmoles/m *yr (Fig. 6.7; Table 6.5). It appears that about 90% of the groundwater inorganic nitrogen is lost through discharge into the inshore waters as opposed to being lost through denitrification.

Let us now consider an area of the island where cesspits contribute significantly to the inorganic nitrogen composition of the groundwaters. Simmons (1983) has reported ZN concentrations averaging greater than 2 mM in the heavily populated areas to the north of the city of Hamilton, and recently these values have been corroborated by Thomson and Foster (1986). Maintaining the same assumptions discussed above, the calculated annual discharge of fixed nitrogen from these nitrate-rich waters was also determined

(Fig. 6 .8 ; Table 6.5). The discharge into the inshore waters is approximately 420 mmoles/m 'yr, some 30 times the amount calculated for the groundwater unaffected by cesspit-derived fixed nitrogen. 190

Table 6.5. Nutrient fluxes.

Natural Cesspit groundwaters groundwaters

2 2 IN RAINFALL 15 nmoles/lm 'yr 15 mmoles/ln ‘yr

AVE. IN CONCENT. 0.065 mM 2 mM 2 2 IN DISCHARGE 13.6 nmoles/to ‘yr 420 mmoles/lm ‘yr 2 DENITRIFICATION 1.4 mmoles/ta *yr 2 P043- DISCHARGE 0.21 mmoles/ta *yr 191

Rainfall Recharge 21 cm/yr 1.5 x lO^^moles/m^yr

'•I'

1.4 x 10 3 <

Organic Nitrogen * Denitrification — * no 3-

no 2"

nh 4+

Discharge to Inshore Waters A o 1.36 x 10 xvnoles/nryr

Figure 6.7. Box model representing natural fixed nitrogen groundwater fluxes. 192

Rainfall Recharge 21 cm/yr 1.5 x 10^ JUfnoles/m^■ yr

CessDit Effluent

Inorganic Nitrogen Pool 2000 Organic Nitrogen NO Denitrification NO

NH

Discharge to Inshore Waters 5 o 4.2x10 Atfnoles/m -yr

Figure 6 .8 . Box model representing cesspit contaminated fixed nitrogen groundwater fluxes. 1 9 3

For phosphate, typical concentrations in these polluted groundwaters

range from < 1 pM to 4 pM, with most values less than 1 pH even where very high nitrates occur. This suggests that the groundwater- 2 to-ocean flux could range from 0.21 nmoles/m 'yr up to 0.84 2 mmoles/m ’yr (Table 6.5).

Other sources of fixed nitrogen and phosphate to the marine environment include rainfall, runoff via storm drains and inputs

from desalinization and coolant (Morris et al., 1977; Bodungen et al., 1982), these latter sources probably being minor island-wide.

Other direct inputs considered were sewage outfalls which account g for 2.5x10 1/yr to the marine environment, 99% of which is directed offshore (Morris et al., 1977). If sewage contains 0.434 mmoles N/l 2 (Waite, 1984), the outfalls are responsible for ~20 nmoles N/m ’yr 2 (0.2 mmoles N/m ‘yr to the nearshore waters) (Table 6 .6 ). Thus, in comparison, the groundwater flux of nitrogen to the nearshore marine 2 environment, which may range from 14 to 420 mmoles N/m ‘yr, must be considered a significant contribution to the total input. Similar sewage outfall calculations for the nutrient phosphate indicate a 2 2 total yearly discharge of 2.2 mmoles/m , 0.02 mmoles/m of which

reach the nearshore as opposed to being directed offshore (Table

6.6).

If groundwaters are, therefore, responsible for much of the fixed nitrogen and phosphate supplied to the inshore waters, then the resulting EN/P ratio should be on the order of 70:1 to 500:1 in enclosed bays and harbors. Ratios of 100:1 have been reported in

Harrington Sound (Morris et al., 1977) and therefore support the data presented here. 1 9 4

Table 6 .6 . Sewage outfalls.

TOTAL SEWAGE DISCHARGED 2.5xl09 1/yc

NITROGEN PHOSPHATE 2 TOTAL 20 nmoles/in2 ‘yr 0.2 nmoles/ta "yr 2 INSHORE 0.2 mmoles/in2 'yr 0 . 0 2 nmoles/to yr CHAPTER VII

CONCLUSIONS

Groundwaters at PRH have some of the lowest chloride concentrations observed in Bermuda groundwaters. These waters have chloride concentrations of "1.3 mM, and previous investigations have shown that this value corresponds with what is believed to be the freshest groundwaters derived from rainfall recharge (Vacher, 1974;

Simmons, 1963). Studies of chloride, ammonium, trace metals (Fe,

Mn, Cd, Cu), alkalinity and petroleum hydrocarbons from landfill leachate and groundwater samples collected in the top 5 m at PRH and at seven other wells around the landfill show no discernible sign of contamination (Simmons, 1985). It should be noted, however, that groundwater density may play an important role in this finding. The density of the landfill leachate may result in a plume which occurs deeper in the aquifer than the standard observations wells allowed us to sample.

It has been possible to approximate the amount of evapo­ transpi ration at 85% of the rainfall, recharge at 21 cm and an initial composition of the freshwater end member associated with the

Pembroke marsh (Table 4.1). These parameters were subsequently extended to a similar wetland, the Devonshire marsh. With detailed chemical data of the two end members, fresh groundwater and seawater of 35% (Stumm and Morgan, 1981) and major element ratios of seawater, it was possible to compare plots of all the major ions 196

(calcium, magnesium, strontium, sodium, potassium, sulfate, fluoride and bromide) versus chloride to theoretical mixing curves to study the geochemical parameters that control the chemistry of Bermuda groundwaters.

The elements calcium, magnesium and strontium have been treated by similar techniques previously by Plummer et al. (1976) for a wide range of wells across the Devonshire lens. The results derived here for depth profiles at PRH and JR wells in the Pembroke and Devonshire marshes yield very similar conclusions to that of

Plummer's work indicating calcium carbonate dissolution and diagenetic processes to be the prime sources of these elements, especially in the vadose zone.

Both sodium and bromide mix conservatively showing seawater- freshwater mixing to be the only process affecting their concentrations while potassium and fluoride do not mix conservatively. In the marshes particularly, and probably in other groundwaters on the island, it appears that potassium may leach from plant cells and result in an enrichment in the upper low salinity groundwaters. The relative enrichment decreases at higher salinities due to mixing with the underlying seawater. Fluoride, on the other hand, evolves from a deep aquifer source. At this time, however, this source has not been identified.

Biogeochemical redox reactions which have been observed in other Bermuda groundwater wells (Simmons, 1983) also occur in the marsh. At the wells studied, PRH and JR, it appears that aerobic respiration and nitrate reduction are the important reactions.

Iron, manganese and sulfate reduction are observed to a very limited 197

extent displayed by the relatively low dissolved concentrations of iron, manganese and hydrogen sulfide. The reason redox reactions do not proceed much beyond nitrate reduction is due to a fairly large and continuous influx of dissolved oxygen and nitrate to the groundwater via precipitation.

An alkalinity model was developed for groundwaters at the PRH and JR sites. According to the alkalinity model, the production of excess alkalinity comes predominantly through the dissolution of the calcium carbonate with the other components sulfide, ammonia, phosphate and silicate contributing negligibly. High PCOj, which evolves from the aerobic respiration and nitrate reduction, causes pH to decrease down the groundwater column and thus calcium carbonate to dissolve. This is clearly depicted at JR, where decreases with depth due to C0 2 production and alkalinity correspondingly increases. However, according to the amount of excess alkalinity present in groundwater at the water table, it appears that the majority of the excess alkalinity may be produced in the unsaturated (vadose) zone.

The flux of fixed nitrogen into and out of the groundwater system seems to be governed by the amount of nitrogen input to the subsurface via rainfall, remineralization of vegetation, cesspits and the amount of annual recharge. For phosphate, reaction with the calcium carbonate aquifer material, and hence retention by the solid phase, is also of great importance. Calculations suggest that in highly populated areas of Bermuda, cesspit effluent dominates the fixed nitrogen chemistry over all other sources. Conservative estimates suggest nutrient nitrogen fluxes to the inshore waters can 1 9 8

be 30 times as great as fluxes from an unpopulated marsh system, and even greater if the Belmont and Walsingham formations, as Plummer et al. (1976) has suggested, have recharges up to 20 times that of the marsh areas. Conversely, phosphate concentrations are controlled probably through phosphate-calcium carbonate surface absorption.

This maintains low and relatively constant groundwater phosphate concentrations. Although, in comparison, marsh phosphate concentrations are statistically somewhat higher than the non-marsh values, annual phosphate fluxes to the inshore waters probably show less variation between densely populated and low populated areas such as the marsh. As a result, groundwater nutrient fluxes, other than direct cesspit input through cracks and fissures, into the nearshore waters are phosphate limited. Therefore, it is unlikely that groundwater seepage alone would be responsible for rapid algal blooms in nearshore coastal areas even when elevated nitrogen from cesspits is significant. CHAPTER VIII

FUTURE WORK

Ihe conclusions of this study have highlighted several interesting trends in the major and minor ions of Bermuda groundwater. Three areas in particular deserve further study.

The first would be identifying the source of fluoride enrichment. Data suggests that a deep aquifer source may be responsible. An obvious next step would include sampling the groundwaters island-wide to confirm whether or not the trends in fluoride enrichment hold, sampling the nearshore sea waters for fluoride enrichments and considering the possibility that the sub­ limestone geology may be in some way involved.

Secondly, the observed excess potassium is believed to have resulted from the breakdown of marsh vegetation. However, questions concerning the vegetation's initial potassium source still need to be addressed. And finally, the importance of groundwater nutrient fluxes to the nearshore waters has been supported by budget calculations. A logical and interesting follow-up would involve making actual groundwater nutrient flux measurements in selected nearshore enclosed bays in order to confirm the calculated fluxes. APPENDIX I

EQUILIBRIUM CALCULATIONS

The equilibrium reactions between chemical species in solution or between solid and solution can be described by the reversible

reaction

aA + bB <----> dD + cC

(Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981) where a, b, c and d are the stochiometric coefficients of the components A, B , C and D. If the rate of the forward reaction is equal to the rate of the reverse reaction, the concentration of each of the components remains constant, and the reaction is said to be at equilibrium. At equilibrium the ratio of the products to the reactants is a constant

K, the equilibrium constant.

{C)C {D}d / {A}* {B}b - K

The brackets {} represent the activity of each of the species.

Activity or effective concentration is a function of the ionic strength of the solution.

{?} - oC

logo - -A2 ^(I"0 .5/(1+(I'S0 .5 ) )-0.2I (Davies equation) 2 0 1

{?} - activity a - activity coefficient C - free ion concentration I - ionic strength z - charge on the ion

The equilibrium constant is a function of both temperature and pressure. For most of the systems that we consider, the effect of pressure on K may be considered negligible, but many waters display wide temperature ranges. Unless otherwise stated, equilibrium constants are usually reported for 25°C. For temperatures other than 25°C corrections for K are made using the Van't Hoff equation

In K - (-AH/R) * ( (1/T2) - (1/Tl))

AH - change in enthalpy Tl - 25 C (298 Kelvin) T2 - new 'temperature (Kelvin) Kl - equilibrium constant at Tl K2 - equilibrium constant at T2

For the groundwater samples in this study the temperatures were approximately 21 to 22°C; therefore, no temperature corrections were deemed necessary, and the calculations were made using the standard equilibrium constants for 25°C (198 Kelvin).

For dissolved solids (e.g. calcite) the equilibrium constant

(K) is termed the solubility constant (Kso), and it defines conditions under which solid dissolution or precipitation takes place.

CaCO, <---- > Ca2+ + CO,- K - 10-8 -48 3 z SO 202

With the appropriate Kgo values and the appropriate free ion concentrations (activity, when ionic strengths are important), the saturation index SI can be determined. The saturation index is described by the equation:

SI - log (IAP/KSQ)

where IAP is the product of the free ions, and the KgQ is the equilibrium solubility constant of the mineral phase of interest. A positive value for SI indicates that the water sample is supersaturated with a particular solid, a negative value indicates under saturation, and SI equal to zero represents equilibrium between the solid and solution phase.

The solids of particular interest in this study were calcite, aragonite, fluorapatite, hydroxyapatite and fluorite.

Solid pK Reference

Calcite 8.48 Plummer and Busenberg (1981) Aragonite 8.34 Plummer and Busenberg (1981) Fluorite 10.4 Stunxn and Morgan (1981) Fluorapatite 60.4 Krauskopf (1979) Hydroxyapatite 57.8 Krauskopf (1979)

Thermodynamic/equilibrium calculations were performed for several water samples from both the JR and PRH wells and also two other wells to determine the aqueous chemical speciation and saturation state. The speciation calculations were performed on the

IBM-compatible Zenith 150 computer using a FORTRAN version of chemical speciation program MINEQL. MINEQL was developed by J.C.

Westall, J.L. Zachary and F.M.M. Morel of the Water Quality 203

Laboratory at the Department of Civil Engineering of the

Massachusetts Institute of Technology (Westall et al., 1976).

MINEQL is designed to handle over 1400 dissolved chemical complexes.

The saturation indices for each of the minerals given above were calculated independently of MINEQL using the equation given above and the Davies equation to correct for ionic strengh effects.

Hie input data was comprised of the species Na+ , K+ , Ca2+,

Mg2+, Sr2+, Cl", S042+, F~, Br- , COj2-, pH, P043+, Fe, Mn, NH4+ , S2_ and Si, each of which were determined by the analyses of the water samples collected from the wells of interest (see Chapter 3 for sample collection and analytic methods). The free ion and composition results of MINEQL and the SI have been tabulated in each chapter.

MINEQL INPUT/OUTPUT

Abbreviations

ID Identification number T Total concentration X Free ion concentration

ID Species

01 Calcium 02 Magnesium 03 Strontium 04 Potassium 05 Sodium 07 Iron 08 Manganese 50 Hydrogen ion 101 Carbonate 102 Sulfate 103 Chloride 104 Fluoride 105 Bromide 107 Ammonium 108 Sulfide 204

109 Phosphate 112 Silicate

ID Solid

20000 calcite 20005 Aragonite 20008 Dolomite 20030 Hydroxyapatite 20035 Fluorapatite 205

MINEQL: Sample printout.

INPUT DATA

01 -2.4 4.S4E-3 02 -2.3 S.69E-3 03 -4.7 2.98E-9 04 -3.01 9.73E-4 OS -1.31 4.91E-2 07 - S . B B.9SE-6 06 -0.9 1.B2E-6 90 -6.97 1.07E-7 101 -S. 1 7.76E-3 102 - 2 . S 9 4.72E-3 103 -1.3 S.70E-2 104 - 3 . 8 1.B2E-4 109 -4. 13 8.29E-S 107 -6.9 3.93E-6 108 -S. 1 7.74E-S 109 - 4 . 2 3.B1E-6 112 -6.0 9.6SE-6 OO 03 SO 6 . 9 7 oo- 04 20000 B. 48 2000S B. 34 1 1101 1 20039 60.4 1 S 104 1 109 3 20030 9 7 . 8 ‘ oo 09 20000 8 . 4 8 20009 8 . 3 4 1 1 lOl 1 20033 60.4 1 9 104 1 109 20030 9 7 . 8 20008 16.7 1 1 2 1 101 20320 10.24 20330 18. 1 20340 20370 10.98 20380 13.9 00 OO

OUTPUT DATA

O IONIC STRENGTH - B.40E-02 LOO F <2— 1) - 10 O ITERATIONS- 3i SOLID 20030 PRECIPITATES O ITERATIONS- 4i SOLID 20330 PRECIPITATES O ITERATIONS- Bl SOLID 20008 PRECIPITATES O ITERATIONS- tOl SOLID 20000 PRECIPITATES O ITERATIONS- 13l SOLID 20000 DISSOLVES 1 INPUT DATA O ID ’ XLOOK T COMPONENTS 0 1 2 . 76E—03 -2.S6 4.S4E-03 CA 0 108 8.12E-11 -10.09 7.74E-0S 8 0 3 2.9BE-0S -4.S9 2.9SE-0S S R 0 4 9.S7E-04 -3.02 9.73E-04 K 0 S 4 . BSE-02 -1.31 4.9IE-02 NA 0 103 8.29E—OS -4.08 S.29E-05 BR 0 a 1.17E-06 -5.93 1.S2E-06 M N 2 0 112 7.05E-14 — 13.15 9.6BE-06 S I Q 3 2 0 6

NH3 .107 3.26E-0B -7 . 4 9 3.5TE-06 o 4.72E-03 804 o 102 3.4BE-03 -2 . 4 6 S.70E-02 -1.24 5.70E-02 CL 0 103 1.B2E-04 F 104 1.61E-04 - 3 . 7 9 © -2 . 3 2 S.69E-03 no 0 ■ 2 4.80E-03 1.07E-07 -6 . 9 7 1.07E-07 H o SO 3.S1E-06 F04 0 109 1.86E-16 - I S . 73 6.64E-08 -7. IB B.95E-06 FE2 0 7 7.76E-03 COO- o lOl S.14E-06 -S. 09 o TYPE I - COMPONENTS o ID LOOK 8PEC 1 E S 0 1 .00 CA 0 2 .00 MB 0 3 .00 8R 0 4 .00 K 0 S .00 NA 0 .7 .00 F E 2 0 8 .00 (4N2 0 112 .00 8 1 0 3 0 101 .00 C03- 0 102 .00 804 0 103 .00 CL 0 104 .00 F 0 105 .00 PR o 107 .00 NH3 .00 8 0 10B P04 o 109. .00 CA o V 0 ID LOOK SPECIES TYPE II 013595 -13.79 H CA C03- 0 1000 2.17 C03- 0 1010 10.77 CA H 0 1020 1.47 CA 804 0 1030 . 6B CA F 0 1040 -. 10 CA NH3 0 1050 -.70 CA NH3 0 1060 -l.SO CA NH3 CA NH3 0 1070 -2.60 P04 0 1080 13.14 CA H 0 1330 -12.41 CA H MS C03- 0 1360 2.37 C03- 0 1370 10.77 MG H 0 13B0 1.57 804 HG 0 1390 1.3B F MG 0 1400 .10 N H 3 MG 0 1410 .00 NH3 2 MG 0 1420 -.30 N H 3 3 MG 4 MG 0 1430 -1.00 N H 3 P04 1 0 1440 13.64 CA 0 MG 1 M 0 1740 -11.41 MG P04 1 12.64 CA 0 BR H 1 O 1730 P04 O 0 1760 4.75 CA 0 8R 0 1950 -12.81 8 R 1 H 0 1960 .68 K 1 804 0 2000 .78 NA 1 C03- 0 2010 .28 NA 1 804 0 2720 1.47 8 0 4 1 FE2 0 2730 .48 CL 1 FE2 0 2740 1.30 N H 3 1 FE2 0 27SO 2. 10 N H 3 FE2 0 2760 3.60 N H 3 4 FE2 0 3280 -8.81 M -1 FE2 0 3290 11.27 M N 2 1 H C 0 3 - 0 3300 1.47 MN2 1 804 0 3310 . 68 M N2 1 CL 0 3320 .4* M N 2 1 CL 2 07

O 3330 -.02 U N 2 1 CL 3 0 3340 .70 MN2 1 NH3 1 0 3350 1.20 MN2 1 NH3 2 O 3360 14.74 CA 0 MN2 1 H 1 P04 0 3400 -10.31 MN2 1 H -1 0 3 * JO -34.20 MN2 1 H -3 012530 4.78 H 1 COS— 1 012540 15.88 H 2 C03- 1 012550 1.7B 804 1 H 1 012560 2. 79 F 1 H 1 012570 9.00 NH3 1 H 1 012580 12.5B S 1 H 1 012590 19.SB 8 1 H 2 012600 11.88 CA 0 H 1 P04 1 012610 1B.B6 CA 0 H 2 P04 0 012620 20.65 CA 0 H 3 P04 1 012710 12.68 8103 1 H 1 012720 22.08 8103 1 H 2 0 0 0 ID LOOK SPECXESi TYPE III - FIXED SOLIDS 0 50 6.97 H 1 0 0 ID LOGK SPECIESt TYPE IV - PRECIPITATED BOLIDS 020030 53.02 CA 5 H -1 P04 3 020330 17.27 B 1 FE2 1 020000 7.65 CA 1 COS- . 1 O 10 LOGK SPECIESI TYPE V - DISSOLVED BOLIDS 020008 15.04 CA 1 MS 1 C03- 2 020040 43.63 CA 4 H 1 P04 3 020050 17.84 CA 1 H 1 P04 1 020070 7.87 CA 1 BIOS 1 020130 -22.11 CA 1 H -2 020140 4.57 MB 1 C O S — 1 020150 7.38 F 2 MB 1 020160 25. 2B MB ’ 3 P04 2 020200 -16.61 MB 1 H -2 020210 B* 17 SR 1 COS— 1 020220 6.47 8R 1 804 1 020230 7.48 ' BR 1 F 2 020240 1.57 BR 1 BIOS 1 020320 9.41 FE2 1 COS— 1 020035 55.41 CA 5 F 1 P04 3 020340 30. IB P04 2 FE2 3 020350 15.57 8103 1 FE2 1 020360 -12.31 H -2 FE2 1 020370 4.75 MN2 1 COS— 1 020380 12.67 -8 1 MN2 1 020400 9.89 MN2 1 BIOS 1 020430 -14.71 MN2 1 H -2 021440 24.78 BIOS 1 H 2 020010 3.87 CA 1 804 1 020020 9.98 CA 1 F 2 020005 7.51 CA 1 COS- 1 AV 0 ID LOGK BPECIEBl TYPE VI - SPECIES NOT CONSIDERED 025000 17.38 H 2 C O S — 1 1 OUTPUT DATAi ITERATIONS ■ IS 0 ID • X LOGX TY SPECIES 0 1 2.76E—03 -2.56 4.54E-03 -1. 04E-O9 CA 0 108 8.12E- 11 -10.09 7.74E-OS -8. OOE-11 8 0 3 2.38E- 05 -4.54 2.5BE-0S -2. 26E-11 BR 0 4 4.57E- 04 -3.02 4.73E—04 8. 19E-11 K 0 5 4.BBE-02 -1.31 4.41E-02 4, 07E-10 NA 0 105 B.24E-05 -4.08 8.29E-05 5. 04E-11 BR 0 8 1. RE­06 -5.93 1.826-06 1. 22E-1S MN2 2 0 8

0 11Z 7.03E-14 -13.13 9.6BE--06 1.09E-11 SI 03 0 107 3.26E-OB -7.49 3.33E—06 2.27E-13 NH3 0 102 3. 4BE-03 -2.46 4.72E-03 -3.B6E-IO 804 0 103 3.70E-02 -1.24 3.70E--02 B.64E-10 CL 0 104 1.61E-04 -3.79 1.82E-04 -1.39E-1I F o 2 4.B0E-O3 -2.32 5.69E— 03 -1.6BE-09 MB 0 3 0 1.07E-07 -6.97 1.07E-07 .OOE+OO H 0 109 1.B6E-16 -15.73 3.B1E--06 .OOE+OO P04 0 7 6.64E-CB -7. IB B. 93 E —06 .0OE*OO FE2 o lOl B. 14E-06 -3.09 7.76E-03 .00E*00 COS- 0 0 ID C LOBC LOOK BPEClEBi TYPE I - COMPONENTS o 1 2.76E-03 -2.56 . OO CA 1 0 2 4.B0E-03 -2.32 .00 MG 1 0 3 2.3BE-OS -4.39 .OO SR 1 0 4 9.37E-04 -3.02 .00 K 1 o 4.SBE-02 -1.31 • OO NA 1 o 7,3 6.64E-0B -7. IB .OO FE2 1 o e 1.17E-06 -3.93 .OO MN2 1 o 112 7.03E-14 -13.15 .OO 6103 1 o lOl B.14E-06 -3.09 • OO C03- 1 o 102 3.4BE-03 -2.46 .00 B04 1 0 103 3.70E-O2 -1.24 .00 CL 1 0 104 1.61E-04 -3.79 • OO F 1 0 105 B.29E-03 -4. OB ■ OO BR 1 0 107 3.26E-06 -7.49 • OO NH3 1 0 108 B.12E-11 -10.09 .OO S 1 Q 109 1.B6E-16 -15.73 .OO CA 0 P04 0 V/s 0 ID C LOGC LOOK BPEClEBi TYPE 11 - COMPLEXES 0 1 3 5 9 9 1.51E-07 -6.82 -13.79 H -1 0 1O00 3.31E-06 -5.48 2.17 CA 1 COS- 1 0 i o i o 1.41E-04 -3.B3 10.77 CA 1 H 1 COS— 0 1020 2.B2E-04 -3.33 1.47 CA 1 604 1 o 1030 2.13E—06 -5.67 . 68 CA 1 F 1 0 1040 7.16E-11 -10.15 -.10 CA 1 NH3 1 0 1050 5.B7E-19 - I B . 23 -.70 CA 1 NH3 2 0 1060 3.04E-27 -26.32 -1.30 CA 1 N H 3 3 0 1070 7.B7E-36 -33.10 -2.60 CA 1 N H 3 4 0 lOBO 7.66E-13 -12.12 13.14 CA 1 H 1 P04 0 1350 I.QIE-OB —B.00 -12.41 CA 1 H -1 0 1360 9.12E-06 -5.04 2.37 MG 1 COS- 1 o 1370 2.4SE-04 -3.61 10.77 MG 1 H 1 C O S — 0 1380 6.1BE-04 -3.21 1.37 804 1 M B t o 1390 1.B7E-05 -4.73 1.3B F 1 M G 1 0 1400 1.97E-10 -9.71 . 10 NH3 1 M G 1 0 1410 3.1IE— IB -17.29 .OO NH3 2 MG 1 o 1 4 2 0 B.36E-26 -23. OB -.30 NH3 3 MG 1 0 1 430 S.44E-34 -33.26 -1.00 NH3 4 M G 1

0 1 440 4.21E-12 - 1 1 . SB 13.64 o o M B 1 PD4 ] H 0 1 7 4 0 1.75E-07 -6.76 -11.41 MG 1 H -1 0 1 7 5 0 2.26E-1S -14.63 12.64 CA 0 ] B R 1 H P04 0 1760 2.71E-16 -15.37 4.73 CA 0 B R 1 P04 0 1 950 3.73E-11 -10.43 -12.81 BA 1 H -1 0 1 9 6 0 1.61E-05 -4.79 .6B K 1 8 0 4 1 0 2 0 0 0 2.41E-06 -3.62 .78 NA 1 C O S - 1 0 2 0 1 0 3.26E-04 -3.49 .28 NA 1 8 0 4 1 0 2 7 2 0 6.79E-09 -B. 17 1.47 804 1 F E 2 1 0 2 7 3 0 1.13E-OB -7.94 .48 CL 1 F E 2 1 0 2 7 4 0 4.32E-14 -13.36 1.30 NH3 1 F E 2 1 0 2 7 3 0 B.90E-21 -20.05 2.10 NH3 2 F E Z 1 0 2760 3.00E-34 -33.32 7. 40 WW7 * p p - < to o Q a . ► 4 _» 0) o to s •« o ** Q IN tO - M o _J I- CO o « i (0 o to HI to * ♦ o o o o d > u a. o. S. id 2 J 8 2 2 U o til to K ID •'^M^iNto^iN** -»ro-4~»*4-4p*-«iN-4fNto-«cN ft. if It I I . I I I I I I I > to 1 1 to 1 I 10 N * tO 10 IN to to IN tO o to « to * o to N N N IO N O ? ID D J J J I I Z o o ID o ID ►4 O Id p O O M o id id u io z « o ilKDUUUZZC ZXUUXXXIIII1XI IlL U tL 0 X 0 c ft.Z U w l l O o >

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02300fe 2. 22E-02 -1. 65 17.3B H COS- 1 PERCENTAGE DISTRIBUTION OF COMPONENTS 0 OCA 0 60. B PERCENT BOUND IN SPECIES 4 1 CA 0 3.1 PERCENT BOUND IN SPECIES 4 1010 CA COS 1 O 6.2 PERCENT BOUND IN SPECIES 4 1020 CA B04 0 27.6 PERCENT BOUND IN SPECIES 420000 CA C03- 0 OS 0 43.1 PERCENT BOUND IN SPECIES 412SB0 S H o 45.4 PERCENT BOUND IN SPECIES 412570 S H 0 11.5 PERCENT BOUND IN SPECIES 420330 S FE2 0 OSR o 100.0 PERCENT BOUND IN BPECIES 4 SR 0 OK 0 78.3 PERCENT BOUND IN SPECIES 4 4 K O 1.7 PERCENT BOUND IN B P E C I E S 4 1760 K B04 0 ONA 0 77.3 PERCENT BOUND IN SPE C I E S 4 5 NA 0 OBR O 100.0 PERCENT BOUND IN SPECIES 4 105 BR 0 0MN2 O 64.5 PERCENT BOUND IN SPECIES 4 B MN2 0 10.4 PERCENT BOUND IN SPECIES 4 3270 HN2 H 1 COS 1 0 6.6 PERCENT BOUND IN SPE C I E S 4 3300 HN2 804 1 0 17.B PERCENT BOUND IN SPE C I E S 4 3310 HN2 CL 1 0 0S103 0 77.6 . PERCENT BOUND IN SPECIES 412720 SI03 H 2 0 0NH3 0 77.1 PERCENT BOUND IN SPECIES 412570 NH3 H 1 O 0804 0 73.7 PERCENT BOUND IN SPEC I E S 4 102 804 0 6.0 PERCENT BOUND IN SPEC I E S 4 1020 CA 604 1 0 13.1 PERCENT BOUND IN SPEC I E S 4 13B0 804 MS 1 o PERCENT BOUNDIN SPEC I E S 4 2010 NA 804 1 0 OCL 0 100.0 PERCENT BOUND IN SPE C 1 E B 4 103 CL 0 OF 0 BS.SPERCENT BOUNDINSPECIES 4 104 F 0 1.2 PERCENT BOUND IN SPECIES 4 1030 CA F 1 0 10.3 PERCENT BOUND IN SPECIES 4 1370 F MS 1 0 OHS 0 84.3 PERCENT BOUND IN B P E C I E S 4 2 MS 0 4.3 PERCENT BOUND INSPECIES 4 1370 MO H I COS 1 0 10.7 PERCENT BOUND IN SPECIES 4 13B0 804 MS 1 0 OH 0 2.0 PERCENT BOUND IN B P E C I E S 4 1010 CA H I COS 1 o 3.4 PERCENT BOUND IN SPE C I E S 4 1370 MB H 1 COS 1 2 1 1

0 73. S PERCENT BOUND IN SPECIES #12530 H 1 COS- 0 19.5 PERCENT BOUNDIN n SPECIES #12540 H 2 C03- <*>04 0 100.0 PERCENT BOUND IN SPECIES #20030 CA 5 H P04 3 0 0FE2 0 99.0 PERCENT BOUND IN SPECIES #20330 8 1 FE2 A V* 0C03- 0 l.B PERCENTBOUND IN SPECIES • i o i o CA 1 H COS “_ 4 1 0 3.2 PERCENTBOUNDIN SPECIES # 1370 MG 1 H COS _ 4 ** 1 0 6B.4 PERCENT BOUND IN BPECIES #12530 H 1 C03- 0 9. 1 PERCENT BOUND IN SPECIES #12540 H C03- 0 17.3 PERCENT BOUND IN BPECIES •20000 CA 1 C03- 2 1 2

Saturation indices for fluorite in marsh groundwaters. I - ionic strength. well I Free Free IAP Corrected SI ca F' log k so

JR Surface 0.02 1.37xl0-3 1.20xl0~ 5 1.97xl0-13 -10.04 -2.66

JR 457 cm 0.07 2.39xl0-3 1.44xl0-4 4.95xl0-11 -9.81 -0.49

JR 609 cm 0.08 2.76xl0-3 1.61xl0-4 7.15xl0-11 -9.79 -0.36

PFH 0.008 1.34xl0~ 3 4.11xl0-6 2.26xl0-14 -10.16 -3.48

TCD 0.01 1.76xl0~3 1.24xl0"5 2.70xl0-13 -10.13 -2.43

TOL 0.62 l.lxlO-2 1.42xl0-4 2.22xl0-10 -9.45 -0.20 213

Saturation indices for calcite and aragonite in marsh groundwaters. I ionic strength.

Free Free Corrected Corrected 2+ 2- lo9 Kso SI lo9 Kso SI Well I Ca COg calcite calcite aragonite arag.

OH _ "3 _ c Surface 0.02 1.37x10 3 7.31x10 0 -7.999 -0.0005 -7.855 -0.14

1311 3 fi 457 cm 0.07 2.39x10 8.6x10 ° -7.698 0.01 -7.554 -0.13

3^ T _ £ 609 cm 0.08 2.76x10 J 8.14x10 0 -7.661 0.01 -7.516 -0.13

PRH 0.008 1.34x10" 3 5.19xl0-6 -8.156 -0.002 -8.012 -0.14

TCD 0.01 1.76x10”2 4.31x10”® -8.122 -0.002 -7.978 -0.14

1WL . 0.62 l.lOxlO"2 3.74x10” 6 -7.212 0.17 -7.068 -0.32 21 4

Saturation indices £or fluorapatite and hydroxylapatite (free ion concentrations in moles/liter).

Well I Free Ca2+ Free F Free P04 3 Free OH

JR 0 cm 0.02 1.37xl0” 3 1 .20xl0-5 1.14xl0-16 1.29xl0-7

JR 457 cm 0.07 2.39xl0-3 1.44xl0” 4 1 .86xl0-1 6 1.15xl0-7

JR 609 cm 0.08 2.76xl0-3 1.61xl0-4 1 .86xl0-16 9.33x10”®

PFH 0.008 1.34xl0” 3 4.llxlO”6 5.84xl0-17 1.32xl0-7

TCD 0.01 1.76xl0~3 1.24xl0“5 4.83xl0-17 9.77xl0-8

TOL 0.62 l.lxlO-2 1.42xl0~4 1.16xl0-16 1.15xl0-7

LAP Corrected IAP Corrected Well OH apatite F apatite l09 KOH SIOH log k f SIF

JR 0 cm 3.03xl0”35 -27.5 -7.03 2.93xl0-34 -28.8 -4.77

JR 457 cm 2.40x10-34 -26.6 -7.01 8.50x10”33 27.8 -4.21

JR 609 cm 3.lOxlO-34 -26.5 -7.01 1.39xl0” 32 -27.7 -4.14

PRH 1.06xl0”35 -27.9 -7.02 5.95xl0“ 35 -29.2 -4.99

TCD 1.36xl0"35 -27.8 -7.01 1.54xl0” 34 -29.1 -4.68

TWL 5.37xl0-33 -25.2 -7.09 1.89xl0-31 -26.4 -4.32 APPENDIX II

X-RAY DIFFRACTION ANALYSES

X-ray diffraction analyses were performed on limestone samples to determine whether or not the minerals calcite, aragonite, hydroxyapatite, fluorapatite and fluorite were present. The analyses were conducted using a General Electric XRD-3 diffractometer. Limestone samples were prepared simply by pulverizing with a mortar and pestle and then gluing the powder to a glass slide with a non-crystalline adhesive. X-rays were generated with a copper tube, emitting copper K alpha radiation of wavelength 1.5405 angstroms, and each sample was scanned over a 10 to 90 angle range. For each mineral the analytical peaks are given in Table (A). 216

Table A. Analytical peaks.

Mineral Rel intensity d-Value Angle

Calcite 100 3.035 29.40 18 2.095 43.14 18 2.285 39.40 17 1.913 47.48 17 1.875 48.50

Aragonite 100 3.396 26.20 52 3.273 27.22 46 2.700 33.15

Hydroxyapatite 100 2.817 31.74 81 2.723 32.86 59 2.779 32.18 41 1.945 46.66

Fluorapatite 100 2.880 31.94 60 2.702 33.12 55 2.772 32.26

Fluorite 100 1.931 47.02 94 3.153 28.22 217

3 * 0 3 8 6 wweii eo»m enoM « 9.04 I.M ft. 10 U— eaco. + to. too 10 19 It Ciiaw e>——iff (Cittirc)

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MAJOW COWWtCTION

CiCO,

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ANGLE IN DEGREES XRD of Limestones PRH 400 cm

ANGLE IN DEGREES j -II-1 I .1 XRD of Limestones JR 0 cm (watertable)

50 40 ANGLE IN DEGREES XRD of Limestones JR 300 cm

j i I

I l

J i

i

| 1 I 5060 40 30 20 10 ANGLE IN DEGREES XRD of Limestones JR 600

T ■' !*i]TP: :T I l; !;iii 224

angle in degrees APPENDIX III

COMPUTERIZED ALKALINITY MODEL

Given the data foe pH, ionic strength, excess alkalinity, excess calcium, excess magnesium, excess strontium, along with the total sulfide, total anmonium, total phosphate and total silicate, this program is designed to calculate the alkalinity of typical Bermuda groundwaters and determine the contribution of the various components. A description of the algorithms used in the program is discussed in Chapter 2 of the text.

ABBREVIATIONS

SULFIDE SPECIES

ST, Total sulfide 50, Hydrogen sulfide H?S 51, Bisulfide HS- 52, Sulfide S2-

AMMONIUM SPECIES

NT, Total ammonium NO, Ammonium NH^ N l , Ammonia NH^

SILICA SPECIES

IT, Total silicate 10, Silicious acid H.SiO. 11, Tri H silicate H^SiO^- 22 6

PHOSPHATE SPECIES

PT, Total phosphate PO, Phosphoric acid H,PO. Pi, Di H phosphate H-PO, P2, Biphosphate HEO. P3, Phosphate P04

ALKALINITY

AA, Analytical (titration) alkalinity AC, Excess alkalinity HA, Model alkalinity MB, Contribution of sulfide, ammonia, phosphate and silicate to alkalinity DA, Difference between analytical alkalinity (AA) and model alkalinity (MA) PD, Percent different of DA

ICNIC STRENGTH Q, Davies equation tion Dl, Component of Davies equation a for singly charged ion D2, Component of Davies equation a for doubly charged ion D3, Component of Davies equation a for triply charged ion Gl, activity coefficient for a singly charged ion G2, activity coefficient for a doubly charged ion G3, activity coefficient for a triply charged ion

IONIZATION FRACTIONS

Al, Alpha 1 ionization fraction A2, Alpha 2 ionization fraction A3, Alpha 3 ionization fraction Kl, Kal, first dissociation (acidity) constant K2, Ka2, second dissociation (acidity) constant K3, Ka3, third dissociation (acidity) constant

For sulfide pKal - 7.02 For ammonium - 9.2 PKal 13.9 PRa 2 -

For silicate pKfll - 9.86 For phosphate - 2.15 PKal 13.1 - 7.20 PKa2 - PKa2 - 12.35 PKa3

(Morel, 1983) 227

10 TEXT : HOME : PRINT 100 INPUT "ENTER PH ' ;PH: PRINT 110 INPUT " TOTAL SULFIDE 'iST: PRINT 120 INPUT “ TOTAL PHOSPHATE '; PT: PRINT 130 INPUT " TOTAL SILICATE ";IT: PRINT 140 INPUT " TOTAL AMMONIA '; NT: PRINT 150 INPUT " EXCESS CALCIUM ' ; CA: PRINT 160 INPUT “ EXCESS MAUHES. *;MG: PRINT 170 INPUT " EXCESS STRONT. ;SR: PRINT iao INPUT " ANALYTICAL ALK ': AA: PRINT 185 INPUT " CONSER. ALK ' ; AC: PRINT 190 INPUT " IONIC STRENGTH ' ; I : PRINT 200 220 = ((I " 0 .5 / (1 + I - 0.5)) - (0.2 * I)) 230 0 .5 * « 235 2 * 0 240 4 .5 * e 245 10 " 1)1 250 10 " D2 255 10 * D3 257 260 U = ( 10 " PH) / G1 265 OH = ( ( 1 0 - 14) t ( H * G1 ) ) / <31 270 280 HEM + 300 K1 - 10 * - S. 24 310 AO = H / (H + K l) 320 A1 = Kl / (H + Kl): REM *** AMMONIUM *** 330 NO = AO * NT:N1 = A1 * NT 350 REM ******************* 400 Kl = (10 * - 7.02) / (G1 * G1):K2 = (10 * - 13.9) / G2 410 E = (H * 2) ♦ (H * Kl) + (Kl * K2) 420 AO = (H * 2) / E: REM *** SULFIDE *** 430 A1 = (H * K l) / E 440 A2 = (K l * K2) / E 450 SO = AO * S T :S I = A1 * S T :S 2 = A2 * ST 460 480 REM **************** 1*1* 500 Kl s (10 * - 9.86) /

670 REM ft***************************** 1000 MA = AC - H ♦ OH + (2 * (CA + MG + SR )) + Nl + S I ♦ ( 2 * S2) + II - PO + P2 + (2 * P3) 1005 MB = N1 ♦ SI + (2 * S2) 4- II - PO + P2 + (2* P3 > 1010 DA = AA - MA;PD = (DA / AA) * 100 1015 1020 TEXT : HOME 1030 PR# I 1035 PRINT GOTO 1300 1100 PRINT HYDROGEN SULFIDE H2S = " SO: PRINT 1110 PRINT "BISULFIDE HS- = “ S i: PRINT 1120 PRINT "SULFIDE S2- = " S2: PRINT 1130 PRINT PHOSPHORIC ACID H3P04= " PO: PRINT 1140 PRINT "DI H PHOSPHATE H2P04 = " PI: PRINT 1150 PRINT ■BIPHOSPHATE HP04 = ” P 2 : PRINT 1160 PRINT "PHOSPHATE P04 = " P 3 : PRINT 1170 PRINT ■SILICIOUS ACID H 4S I04= " 10: PRINT 1180 PRINT •TRJ H SILICATE H 3S I04= " 11: PRINT 1190 PRINT •AMMONIUM NH4+ = " NO: PRINT 1200 PRINT •AMMONIA NH3 = " N1 : PRINT 1210 PRINT ‘ANALYTICAL ALK = " AA: PRINT 1220 PRINT 'MODEL ALK = " MA: PRINT 1230 'DIF IN ALK = - DA. PRINT 1240 PRINT 'PERCENT D IF ALK = " PD: PRINT 1250 PRINT 'CONTRIB. S.N.P.SI = " MB: PRINT 1254 PRINT : PRINT : PRINT PRINT 1255 PRINT "ION. STR.=";I 1256 PRINT "GAMMA 1 = ";G1 1257 PRINT "GAMMA 2 = ";G 2 1256 PRINT "GAMMA 3 = ';G3 1260 PR# 0 1270 GOTO 1400 1280 1300 PRINT "TOT SULFIDE = ST: PRINT 131 0 PRINT "TOT PHOSPHATE = PT: PRINT 1320 PRINT "TOT AMMONIUM = NT: PRINT 1330 PRINT "TOT SILICATE = IT : PRINT 1340 PRINT " pH = PH: PRINT 1345 PRINT "CONSER. ALK = AC: PRINT 1350 PRINT “EX CALCIUM = CA: PRINT 136 0 PRINT "EX MAGNES = MG: PRINT 1370 PRINT "EX STRON. = SR: PRINT 1380 PRINT : PRINT : PRINT GOTO 1100 1390 PRINT "IONIC STREN. = I 1400 END APPENDIX IV

DATA

units

Depth: cm Temperature: C pH: - Alkalinity: roeq/1 Redox Potential (Eh): mvolts Conductivity: Mseimans Iron: pg/1 Manganese: pg/1

Itie following constituents are reported in mmoles/1 .

Sodium Potassium Calcium Magnesium Strontium Chloride Sulfate Bromide Fluoride Oxygen Sulfide

The following constituents are reported in Umoles/1.

Nitrate Nitrite Ammonium Phosphate Silicate PRH BAIL FEB 11

DEPTH F. TEMP SODIUM POTASSIUM CALCIUM MAGNESIUM STRONTIUM CHLORIDE SULFATE FLUORIOE BROMIDE 0 21.6 1.54 0.29 1.62 0.27 0.009 1.36 0.12 0.011 0.01 152 21.7 1.44 0.29 1.32 0.27 0.007 1.29 0.13 0.026 0.03 305 22.0 1.44 0.30 1.29 0.25 0.007 1.31 - 0.011 0.22 366 21.6 1.44 0.30 1.46 0.25 0.007 1.31 0.13 0.013 -

DEPTH QXIGEN NITRATE NITRITE AMMONIUM PHOSPHATE SILICATE 0 0 20.1 0.27 47.9 3.0 68.5 152 0 20.2 0.21 53.1 2.3 23.9 305 0 19.9 0.13 54.9 2.7 23.9 366 0.001 20.4 0.04 53.2 2.0 19.1 230

DEPTH Na/CI K/CI Ca/CI Mg/Cl Sr/Cl S04/CI f;ci Br/CI 0 1.13 0.21 1.19 0.2 0.007 0.1 0.01 0.01 152 1.12 0.23 1.03 0.21 0.006 0.1 0.02 0.02 305 1.10 0.23 0.10 0.19 0.006 . 0.01 0.17 366 1.10 0.23 1.12 0.19 0.006 0.1 0.01 - P R H P U M P FEB 11

DEPTH F. TEMP L. TEMP pH ALKALINITY CONDUCT SODIUM POTASSIUM CALCIUM MAGNESIUM 0 23.4 20.0 7.120 6.312 780 1.49 0.34 2.79 0.30 152 22.4 20.5 7.171 6.276 755 1.49 0.34 2.81 0.30 305 22.7 20.5 7.116 6.251 660 1.54 0.38 2.88 0.32 396 22.9 20.5 7.179 6.311 680 1.54 0.35 2.80 0.31

DEPTH STRONTIUM CHLORIDE SULFATE FLUORIDE BROMIDE 0 0.014 1.30 0.12 0.004 0.012 152 0.015 1.31 0 13 0.005 0.014 305 0.014 1.31 0.13 0.005 0.012 396 0.015 1.30 0.14 0.005 0.012

DEPTH C0CVGB4 NITRATE NITRITE AMMONIUM PHOSPHATE IRON MANGANESE 0 0.03 1.5 0.21 57.2 1.3 45.7 40.4 152 0.05 2.3 0.25 44.2 1.5 127.9 40.9 305 0.04 4.1 0.19 50.7 1.2 46.7 39.2 396 0.15 3.5 0.21 39.0 1.0 28.7 39.3

DEPTH Na/CI K/CI ca/ci Mg/Cl Sr/Cl S04/CI F/CI Br/CI 0 1.15 0.26 2.14 0.23 0.01 0.09 0.003 0.01 152 1.14 0.26 2.14 0.23 0.01 0.10 0.004 0.01 305 1.18 0.29 2.20 0.24 0.01 0.10 0.004 0.01 396 1.19 0.27 2.15 0.24 0.01 0.11 0.004 0.01

4 JR BAIL MAR 1

DEPTH F. TEMP L. TEMP pH Eh ALKALINITY CONDUCT SODIUM POTASSIUM CALCIUM MAGNESIUM 0 20.4 23.0 - 330 640 3.1 0 28 2.1 0.7 305 20.4 23.0 - -85 1200 9.8 0.39 2.6 1.5 457 20.8 23.0 - -1 2 0 2000 11.1 0 46 2.8 1.8 610 21.3 23.0 - -1 3 0 2600 14.9 0.49 3.0 2.0

DEPTH STRONTIUM CHLORIDE SULFATE FLUORIDE BROMIDE 0 0.009 3.0 0.28 0.013 0.005 305 0.012 8.6 0.26 0.008 0.017 457 0.014 12.5 0.42 0.035 0.031 610 0.014 15.7 0.56 0.038 0.029 232

DEPTH CKVGS4 NITRATE NITRITE AMMONIUM PHOSPHATE SILICATE SULFIDE IRON MANGANESE 0 0.07 123.5 0.21 7.1 0.6 60.4 - 4.2 2.7 305 0 25.2 0.96 54.9 1.3 6S.9 0.05 11.7 16.7 457 0.06 - 0.03 74.5 1.7 30.5 - 10.6 18.1 610 0 - 0.03 69.1 1.7 31.0 0.05 9.6 19.9

DEPTH Na/CI K/CI Ca/CI Mg/Cl Sr/Cl S04/CI F/CI Br/CI 0 1.04 0.10 0.71 0.24 0.003 0.10 0.004 0.002 305 1.13 0.05 0.30 0.18 0.001 0.03 0.001 0.002 457 0.88 0.04 0.22 0.14 0.001 0.03 0.002 0.002 610 0.94 0.03 0.19 0.13 0.001 0.03 0,002 0.002 JR PUMP MAR 1

DEPTH F. TEMP LTEMP pH Efl ALKALINITY COMXJCT SODIUM POTASSIUM CALCIUM MAGNESIUM 0 21.5 23.9 7.392 -80 6.748 1200 7.9 0.34 2.6 1,3 152 21.3 24.0 7.271 -93 7.178 3850 23.5 0.64 3.4 3.0 305 21.5 24.0 7.257 -95 7.314 4300 27.1 0 68 3.6 3.4 457 21.6 24.0 7.357 -85 7.475 4800 32,0 0.78 3.8 4.0

DEPTH STRONTIUM CHLORIDE SULFATE FLUORIDE BROMIDE 0 0.014 8.6 0.4 0.017 0.02 152 0.021 27.3 1.0 0.108 0.04 305 0.024 31.1 1.2 0.121 0.04 457 0.025 35.2 1.5 0.138 0.05 233

DEPTH OXYG04 NITRATE NITRITE AMMONIUM PHOSPHATE SILICATE SULFIDE IRON MANGANESE 0 0.03 51.1 0.21 36.4 0.6 0.03 19.0 17.4 152 0 24.3 0.37 53.3 0.9 0.04 22.3 24.0 305 0 25.2 0.25 58.5 0.7 0.06 19.1 26.8 457. 0.01 14.7 0.29 55.9 0.8 0.07 22.3 26.4

DEPTH Na/CI K/CI Ca/CI Mg/Cl Sr/Cl S04/CI F/CI Br/CI 0 0.91 0.04 0.3 0.15 0.002 0.05 0.002 0.003 152 0.86 0.02 0.12 0.11 0.001 0.04 0.004 0.001 305 0,87 0.02 0.11 0.11 0.001 0.04 0.004 0.001 457 0.91 0.02 0.11 0.11 0.001 0.04 0.004 0.002 JR BAIL MAY 16

DEPTH F. TEMP LTEMP pH Eh ALKALINITY 00NCUCT SODIUM POTASSIUM CALCIUM MAGNESIUM 0 20.9 25.9 7 150 350 5.378 670 4.7 0 33 152 21.0 25 9 7.017 -180 7.313 1700 12.3 0 42 305 21.3 25.9 6.882 -120 7.944 2000 14.9 0 48 457 21.2 26.0 6.927 -123 7.847 2300 16.8 0.49 610 21.3 26.0 6.938 -120 7.356 3000 21.2 0.57

DEPTH STRONTIUM CHLORIDE SULFATE FLUORIDE BROMIDE 0 4.0 0.37 -- 1 5 2 12.2 0.50 0.03 0.02 305 14.8 0.61 0.04 0.01 4 5 7 16.9 0.66 0.04 0.03 610 22.3 0.80 0.10 0.04 234

DEPTH OXVGEN NITRATE NITRITE AMMONIUM PHOSPHATE 0 0.04 136.0 0.52 0.09 0.5 152 0 55.3 0.34 2.08 2.3 305 0 41.9 0.57 2.48 1.7 4 5 7 0 44.5 0.16 3 28 1.8 610 0 42.3 0.11 3.61 1.7

DEPTH K/CI Ca/CI Mg/Cl Sr/Cl S04/CI F/CI Br/CI O 0.08 - 0.09 -- 1 5 2 0.03 - -- 0.04 0.002 0.002 305 0.03 - - 0.04 0.003 0.001 4 5 7 0.03 - - - 0.04 0.003 0.002 610 0.02 - - 0.04 0.004 0.002 JR PUMP MAY 16

DEPTH F. TEMP L. TEMP pH Eh ALKAUNITY OCMXCT SODIUM POTASSIUM CALCIUM MAGNESIUM 0 21.3 25.9 7.108 -75 6.906 1500 9.8 0.48 2.4 1.6 152 21.3 25.2 7.040 -70 7.098 1350 16.4 0.51 3.0 2.1 305 21.3 25.2 7.028 -75 7.871 4900 41.3 0.99 4.2 5.5 457 21.7 25.2 7.064 -65 7.686 4400 43.5 0.79 4.2 5.6 610 21.6 25.2 6.974 -70 7.764 3800 49.1 0.97 4.5 5.7

DEPTH STRONTIUM CHLORIDE SULFATE FLUORIDE BROMIDE 0 0.01 11.3 0.5 0.01 0.02 152 0.01 17.5 1.4 0.06 0.03 305 0.02 48.4 3.3 0.17 0.08 457 0.02 48.9 3.5 0.16 0.07 610 0.02 57 4.7 0.18 0.08 235

DEPTH OXVGS4 NITRATE NITRITE AMMONIUM PHOSPHATE SILICATE SULFIDE IRON MANGANESE 0 0.002 8.9 0.45 1.8 1.7 14.5 0.056 31.9 27.0 152 0.002 22.8 0.33 2.4 1.7 t4 .9 0.066 31.9 29.8 305 0 7.1 0.10 3.4 1.5 8.5 0.074 24 5 22.9 457 0 13.3 0.13 3.3 2.5 3.8 0.073 26.6 27.3 610 0 12.5 0.13 3.5 3.9 9.7 0.077 27.6 32.5

DEPTH Na/Cl K/CI Ca/CI Mg/Cl Sr/CJ S04/CI F/CI Br/CI 0 0.86 0.04 0.21 0.14 0.0010 0.05 0.001 0.002 152 0.85 0.02 0.08 0.11 0.0005 0.07 0.003 0.002 305 0.94 0.03 0.17 0.12 0.0008 0.08 0.003 0.002 457 0.89 0.02 0.09 0.11 0.0004 0.07 0.003 0.001 610 0.86 0.02 0.08 0.10 0.0004 0.08 0.003 0.001 REFERENCES

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