Radium Isotope Geochemistry in Groundwater Systems
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Radium Isotope Geochemistry in Groundwater Systems: The Role of Environmental Factors by David Stewart Vinson Division of Earth and Ocean Sciences Duke University Date:_______________________ Approved: ___________________________ Avner Vengosh, Supervisor ___________________________ Paul Baker ___________________________ Gary Dwyer ___________________________ Heileen Hsu‐Kim Dissertation submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Division of Earth and Ocean Sciences in the Graduate School of Duke University 2011
Transcript of Radium Isotope Geochemistry in Groundwater Systems
Radium Isotope Geochemistry in Groundwater Systems:
The Role of Environmental Factors
by
David Stewart Vinson
Division of Earth and Ocean Sciences Duke University
Date:_______________________ Approved:
___________________________
Avner Vengosh, Supervisor
___________________________ Paul Baker
___________________________
Gary Dwyer
___________________________ Heileen Hsu‐Kim
Dissertation submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Division of
Earth and Ocean Sciences in the Graduate School of Duke University
2011
ABSTRACT
Radium Isotope Geochemistry in Groundwater Systems:
The Role of Environmental Factors
by
David Stewart Vinson
Division of Earth and Ocean Sciences Duke University
Date:_______________________ Approved:
___________________________
Avner Vengosh, Supervisor
___________________________ Paul Baker
___________________________
Gary Dwyer
___________________________ Heileen Hsu‐Kim
An abstract of a dissertation submitted in partial fulfillment of the requirements for the degree
of Doctor of Philosophy in the Division of Earth and Ocean Sciences in the Graduate School of
Duke University
2011
Copyright by David Stewart Vinson
2011
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Abstract
Prior studies of groundwater systems have associated increasing salinity and anoxic
conditions with increasing radium (Ra) activities in water due to the decreasing effectiveness of
Ra removal processes. However, the components of salinity (e.g. Ca vs. Na and SO42‐ vs. Cl‐‐
dominated waters), and the relative importance of salinity‐sensitive vs. redox‐sensitive
processes for Ra mobilization, are less well understood. In this research, the response of Ra to
hydrochemical change was examined using a multiple tracer approach to obtain detailed
information on divalent cation and Ra mobility. A range of salinity and redox conditions was
examined in five field‐based studies in the United States and Morocco: (1) fresh waters in
fractured crystalline rocks in the Piedmont region of North Carolina; (2) the Willcox Basin, an
oxic alluvial basin‐fill aquifer in southeastern Arizona; (3) the Jordan sandstone aquifer, a
carbonate‐cemented quartz sandstone in southeastern Minnesota; (4) an unconfined coastal
aquifer undergoing salinization in the city of Agadir, Morocco; and (5) the confined, fresh to
saline Cretaceous and Pliocene aquifers of the Atlantic Coastal Plain in North Carolina.
In addition to analysis of major element concentrations, trace metal concentrations, and
224Ra, 226Ra, and 228Ra activities, complementary isotope systems were applied to gain insights
on the relative stability of chemical processes that remove radium and other alkaline earth
metals: (1) strontium isotope ratios (87Sr/86Sr) trace divalent cation release from sources such as
clay and carbonate minerals in the aquifer solids and also indicate conditions in which divalent
cation release (rather than uptake) is dominant; (2) boron concentrations and isotopes (δ11B)
coincide with the opposite condition in freshening conditions of the Atlantic Coastal Plain, in
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which divalent cations are removed in exchange for Na; and (3) sulfur and oxygen isotopes (δ34S,
δ18O) of sulfate trace sulfate sources and provide information on sulfate‐reducing conditions,
which can inhibit barite (BaSO4) from removing Ra by coprecipitation. In addition, other isotopic
and ion measurements trace salinity sources and groundwater residence time, including δ2H,
δ18O, 3H, Br‐/Cl‐, Na/Cl‐, and Ca/Na.
This dissertation documents correlations between salinity and radium in the brackish to
saline North Carolina coastal plain aquifer with total dissolved solids (TDS) up to ~18,000 mg L‐1
and to some degree in the Moroccan coastal aquifer, but even the lower‐salinity waters (TDS
<3000 mg L‐1) exhibit a range of Ra activities spanning approximately 3 orders of magnitude.
Among these low‐TDS waters, the highest Ra activities were observed in the anoxic Jordan
sandstone aquifer and the lowest were observed in the oxic Willcox Basin aquifer. Although the
main control on radium activities in fresh groundwater is the U‐ and Th‐series radionuclide
content of the aquifer solids, important secondary controls include the stability of redox‐
sensitive radium adsorption sites (Mn and Fe oxides), the relative dominance of divalent vs.
monovalent cations (e.g. the Ca/Na ratio), formation of the uncharged RaSO40 complex, and/or
the saturation state with respect to barite. These processes interact in varied ways in the field‐
based studies. Increasing radium activities and decreasing 222Rn/226Ra ratios in the North
Carolina fractured crystalline rock groundwater system are correlated with increasing Ba, Mn,
and Fe concentrations and decreasing dissolved oxygen concentrations, related to weathering
and/or organic carbon oxidation. Radium activities in the oxic, neutral to slightly basic Willcox
Basin are very low (median 226Ra activity 2 mBq L‐1), probably due to a combination of effective
Ra removal processes including adsorption to Mn and/or Fe oxides and the overall removal of
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divalent cations during groundwater evolution in this system. These are the same surface charge
conditions that release arsenic, of regional water concern, in this pH range. Radium in Jordan
aquifer groundwater is dependent on local variations in solid‐phase radionuclide levels,
probably hosted in the carbonate cement phase. Also, Ra is inefficiently adsorbed to the aquifer
solids in the aquifer’s anoxic conditions, resulting in the highest radium levels reported in this
dissertation (226Ra up to 420 mBq L‐1) despite apparent barite precipitation that partially
removes Ra. Radium‐224 activity in the Moroccan coastal aquifer is associated with salinity, but
Ra overall is apparently controlled by barite, indicated by conditions near BaSO4 saturation.
Radium activity in the saline waters of the Atlantic Coastal Plain aquifers is associated with TDS
concentrations, but the cation exchange properties of the aquifer may provide a major
mechanism of Ra removal in the Na‐HCO3‐ and Na‐Cl‐ waters. Overall, the complex interaction
between groundwater chemistry and Ra‐removing processes implies that in waters with TDS
below approximately 3,000 mg L‐1, dissolved solids concentration alone does not fully describe
radium’s response to hydrochemical conditions, but rather that aquifer‐specific examination of
Ra removal mechanisms is needed.
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Contents
Abstract........................................................................................................................................... iv
List of Tables ..................................................................................................................................xiv
List of Figures .................................................................................................................................xvi
Acknowledgements.......................................................................................................................xxii
1. Introduction ................................................................................................................................. 1
1.1 Overview of radium isotope geochemistry ......................................................................... 1
1.2 Health effects and drinking water standards ...................................................................... 1
1.3 The link between radium abundance and hydrogeochemical change ............................... 3
1.4 Modeling radium mobilization in groundwater systems .................................................... 4
1.5 Radium source mechanisms................................................................................................ 4
1.5.1 Alpha recoil..................................................................................................................... 4
1.5.2 Desorption and dissolution ............................................................................................ 5
1.6 Radium removal mechanisms ............................................................................................. 6
1.6.1 In situ radioactive decay................................................................................................. 6
1.6.2 Adsorption and cation exchange.................................................................................... 6
1.6.2.1 Salinity effects......................................................................................................... 7
1.6.2.2 pH effects ................................................................................................................ 8
1.6.2.3 Redox effects .......................................................................................................... 8
1.6.3 Coprecipitation into sulfate minerals............................................................................. 9
1.7 Controls on radium isotope ratios..................................................................................... 10
1.7.1 228Ra/226Ra .................................................................................................................... 10
1.7.2 224Ra/228Ra .................................................................................................................... 11
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1.7.3 222Rn/226Ra and 222Rn/224Ra........................................................................................... 11
1.8 Approach of this research ................................................................................................. 12
2. Relationships between radium and radon occurrence and hydrochemistry in fresh groundwater from fractured crystalline rocks, North Carolina (USA) ........................................... 17
2.1 Introduction ................................................................................................................... 17
2.2 Methods ............................................................................................................................19
2.2.1 Sampling procedures.................................................................................................... 19
2.2.2 Radium isotope analysis............................................................................................... 21
2.2.3 Major ions and trace elements .................................................................................... 23
2.2.4 Data handling ............................................................................................................... 23
2.3 Results ...............................................................................................................................23
2.3.1 Major ions, trace elements, and redox‐sensitive elements......................................... 23
2.3.2 Radionuclide activities in water .................................................................................. 24
2.3.3 Radionuclide content of granite rock samples............................................................. 27
2.4 Discussion ..........................................................................................................................27
2.4.1 Major element and redox chemistry of potential Ra sinks .......................................... 27
2.4.2 Radium and radon isotopes ......................................................................................... 30
2.4.2.1 228Ra/226Ra activity ratio........................................................................................ 31
2.4.2.2 224Ra/228Ra activity ratio ....................................................................................... 33
2.4.2.3 Radon‐222............................................................................................................ 34
2.4.2.4 222Rn/226Ra activity ratio....................................................................................... 35
2.4.3 Overview of radionuclide sources and removal mechanisms...................................... 36
2.5 Conclusions........................................................................................................................39
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3. Arsenic and other oxyanion‐forming trace elements in an alluvial basin aquifer: Evaluating sources and mobilization by isotopic tracers (Sr, B, S, O, H, Ra) ................................................... 52
3.1 Introduction.......................................................................................................................52
3.2 Geologic setting................................................................................................................. 55
3.2.1 Composition of the mountain blocks ........................................................................... 56
3.2.2 Basin geometry............................................................................................................. 56
3.2.3 Pleistocene‐Holocene Lake Cochise ............................................................................. 57
3.2.4 Groundwater flow and residence time ........................................................................ 57
3.3 Methods ............................................................................................................................58
3.3.1 Well sampling ............................................................................................................... 58
3.3.2 Major and trace element analysis ................................................................................ 59
3.3.3 Stable and radiogenic isotopes .................................................................................... 60
3.3.4 Radium isotopes........................................................................................................... 62
3.4 Results ...............................................................................................................................63
3.4.1 Groundwater salinity and major element chemistry ................................................... 63
3.4.2 Trace metals ................................................................................................................. 64
3.4.3 Environmental isotopes (δ2H, δ18O, 3H, 87Sr/86Sr, δ11B, δ34S) ....................................... 64
3.4.4 Radionuclides ............................................................................................................... 66
3.5 Discussion ..........................................................................................................................69
3.5.1 Recharge and groundwater residence time................................................................. 69
3.5.2 Isotopic tracing of sediment sources using Sr.............................................................. 74
3.5.3 Sulfate sources ............................................................................................................. 77
3.5.4 Significance of the Ca/Na ratio..................................................................................... 79
3.5.5 Trace element and isotope behavior: Implications for oxyanion‐forming elements ..80
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3.5.5.1 Arsenic .................................................................................................................. 80
3.5.5.2 Vanadium, chromium, selenium, molybdenum, and antimony........................... 82
3.5.5.3 Boron .................................................................................................................... 83
3.5.5.4 Radium and uranium ............................................................................................ 84
3.6 Conclusions........................................................................................................................85
4. Radium and strontium isotopes in the Jordan sandstone (Minnesota, USA): Carbonate control of a sandstone aquifer ................................................................................................................... 88
4.1 Introduction.......................................................................................................................88
4.2 Methods ............................................................................................................................93
4.2.1 Sampling of groundwater............................................................................................. 93
4.2.2 Major and trace element concentrations .................................................................... 93
4.2.3 Strontium isotopes and tritium.................................................................................... 93
4.2.4 Radium isotopes........................................................................................................... 94
4.2.5 Analysis of aquifer solids .............................................................................................. 95
4.2.6 Supplemental data sources .......................................................................................... 96
4.3 Results ...............................................................................................................................97
4.3.1 Major element chemistry and redox conditions.......................................................... 97
4.3.2 Strontium...................................................................................................................... 98
4.3.3 Barium, radium, and uranium ...................................................................................... 98
4.3.3.1 Water samples ...................................................................................................... 98
4.3.3.2 Aquifer solids ...................................................................................................... 101
4.4 Discussion ........................................................................................................................101
4.4.1 Groundwater residence time and evolution of dissolved solids................................ 101
4.4.2 Strontium isotope insights on the carbonate phase .................................................. 106
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4.4.3 The significance of 228Ra/226Ra.................................................................................... 108
4.4.4 Radium removal mechanisms .................................................................................... 110
4.4.4.1 Adsorption and redox conditions ....................................................................... 110
4.4.4.2 Coprecipitation into barite ................................................................................. 111
4.4.4.3 Sources of barium to the Jordan aquifer ............................................................ 113
4.4.5 Relationship to alpha recoil models of Ra isotopes in groundwater ......................... 116
4.4.6 Timing of U addition to aquifer solids ........................................................................ 118
4.5 Conclusions......................................................................................................................120
5. Geochemical evaluation of salinity sources and radium response to salinization of an unconfined coastal aquifer in Morocco....................................................................................... 131
5.1 Introduction.....................................................................................................................131
5.2 Hydrogeology of the study area ...................................................................................... 134
5.3 Methods ..........................................................................................................................135
5.3.1 Sample collection ....................................................................................................... 135
5.3.2 Major and trace element analysis .............................................................................. 135
5.3.3 Environmental isotopes (δ2H, δ18O, δ34SSO4, δ18OSO4,
87Sr/86Sr)................................... 136
5.3.4 Radium ....................................................................................................................... 137
5.3.5 Speciation calculations ............................................................................................... 138
5.4 Results .............................................................................................................................138
5.4.1 Major element chemistry........................................................................................... 138
5.4.2 Environmental isotope ratios ..................................................................................... 140
5.4.3 Radium, barium, and uranium ................................................................................... 144
5 Discussion ...........................................................................................................................148
5.5.1 Groundwater evolution and salinity source identification ........................................ 148
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5.5.2 Cation exchange and 87Sr/86Sr .................................................................................... 152
5.5.3 Boron’s sensitivity to seawater intrusion and ion exchange ..................................... 155
5.5.4 Sulfate sources and barite solubility .......................................................................... 157
5.5.5 Radium activities ........................................................................................................ 160
5.5.6 Radium isotope ratios ................................................................................................ 162
5.6 Conclusions......................................................................................................................166
6. Radium isotope response to salinity gradients in confined Atlantic Coastal Plain aquifers....168
6.1 Introduction.....................................................................................................................168
6.2 The Atlantic Coastal Plain aquifer system: Natural and increasing salinity .................... 170
6.3 Methods ..........................................................................................................................172
6.3.1 Water Sampling .......................................................................................................... 172
6.3.2 Major and trace elements .......................................................................................... 175
6.3.3 Environmental isotopes (δ11B, δ34SSO4, δ18OSO4,
87Sr/86Sr)........................................... 175
6.3.4 Radium isotopes......................................................................................................... 176
6.3.5 Speciation calculations ............................................................................................... 177
6.4 Results .............................................................................................................................177
6.4.1 Major and trace element chemistry........................................................................... 177
6.4.2 Strontium and boron.................................................................................................. 178
6.4.3 Sulfur and oxygen isotopes of sulfate ........................................................................ 184
6.4.4 Barium and Radium.................................................................................................... 184
6.5 Discussion ........................................................................................................................187
6.5.1 Major element chemistry and the aquifers’ response to salinization ....................... 187
6.5.2 Redox conditions ........................................................................................................ 189
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6.5.3 Strontium.................................................................................................................... 190
6.5.4 Boron.......................................................................................................................... 192
6.5.5 Radium and barium.................................................................................................... 195
6.5.5.1 Salinity effects..................................................................................................... 195
6.5.5.2 Redox effects and the barite removal mechanism for radium........................... 198
6.5.5.3 Radium isotope ratios......................................................................................... 202
6.6 Conclusions......................................................................................................................203
7. Synthesis ..................................................................................................................................205
7.1 The components of salinity: Cation and anion effects.................................................... 205
7.2 Salinity vs. redox effects.................................................................................................. 207
References ...................................................................................................................................210
Biography .....................................................................................................................................241
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List of Tables
Table 1: United States and international standards for radon and radium in drinking water. ....... 3
Table 2: First‐order rate constants for the terms in equation 1...................................................... 7
Table 3: Summary of major and trace element concentrations and pH. ND indicates not detected. ........................................................................................................................................ 41
Table 4: Major and trace element concentrations and pH............................................................ 42
Table 5: Summary of radionuclide activities and activity ratios. ................................................... 47
Table 6: Radionuclide activities and activity ratios........................................................................ 48
Table 7: Radium isotope activities in Rolesville Granite samples from cores WC‐1 and WC‐3. Errors are ±2σ from counting statistics. ........................................................................................ 51
Table 8: Major element results and mineral saturation index values. .......................................... 71
Table 9: Concentrations of oxyanion‐forming and other trace metals. See Table 10 for B and Sr concentrations. .............................................................................................................................. 72
Table 10: Isotope results and related trace element concentrations. .......................................... 73
Table 11: Simplified 238U and 232Th decay series including nuclides discussed in this paper. .......90
Table 12: Simplified list of hydrostratigraphic units discussed in this paper (modified from Mossler, 1990; Runkel et al., 2003). .............................................................................................. 91
Table 13: Names and well construction information for the wells sampled in this study. Six‐digit well identifiers are unique identifiers in the Minnesota Department of Health County Well Index (http://www.health.state.mn.us/divs/eh/cwi/). Samples with dates not listed are unfiltered, unpreserved supplemental water samples analyzed for anions and 87Sr/86Sr only. ...................122
Table 14: Major element chemistry and saturation index values (SI) of Jordan aquifer water samples. .......................................................................................................................................124
Table 15: Trace metal concentrations, Sr isotope ratios, and tritium and radium activities of Jordan aquifer groundwater. Italized tritium activites are from samples collected by the Minnesota Department of Health on different dates than indicated in Table 13.......................127
Table 16: Results of analysis of aquifer solids from Jordan aquifer wells. .................................. 130
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Table 17: Major element concentrations and saturation index values of groundwater samples......................................................................................................................................................141
Table 18: Element and ion ratios diagnostic of salinity sources in the coastal aquifer...............142
Table 19: Environmental isotope ratios, trace metal concentrations, and radium activities. ....147
Table 20: Spearman rank coefficients between major ion compositions, and radium and barium. Note (1) the decreasing values of ρ as radium half‐life increases (left to right); (2) the opposite trend associated with sulfate concentration; and (3) the close similarity between the behavior of Ba and 226Ra, but not other Ra isotopes. ..................................................................................... 149
Table 21: Major element concentrations (in mg L‐1) of the Yorktown and Cape Fear aquifers. Data in italics appeared in Vinson et al. (in press)....................................................................... 185
Table 22: Trace metal concentrations and isotope ratios of Yorktown and Cape Fear aquifer groundwater. Concentrations are in mg L‐1. Data in italics appeared in Vinson et al. (in press).186
Table 23: Components of salinity and redox conditions that are correlated with increasing (↑) or decreasing (↓) radium activities in the five studied groundwater systems. ..............................205
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List of Figures
Figure 1: Simplified 238U (black), 235U (green), and 232Th (blue) decay series including uranium, thorium, radium, and radon isotopes (modified from Molinari and Snodgrass, 1990). ................. 2
Figure 2: Diagram indicating the approximate ranges of salinity and redox conditions encountered in this study. ............................................................................................................. 14
Figure 3: Map of the study area showing the six sample groups. Numbers in parentheses are the number of sites in each group; political boundary shown is of Wake County, North Carolina. Geological data based on Clark et al. (2004); data for location map obtained from www.nationalatlas.gov. .................................................................................................................20
Figure 4: Relationships between pH, dissolved oxygen concentration, and total ions. ................ 25
Figure 5: Relationship between dissolved oxygen and other redox‐sensitive solutes. ................. 26
Figure 6: Relationship between dissolved oxygen, total ions, and activities of 222Rn and 226Ra in Rolesville Granite groundwater. .................................................................................................... 28
Figure 7: 228Ra and 224Ra activities in groundwater. Dashed line represents equilibrium value of 1. Counting error omitted where smaller than symbol size. ............................................................. 33
Figure 8: Relationship between 222Rn and 226Ra (left), 222Rn and 224Ra (right) and histograms of isotope activities in Rolesville Granite groundwater. .................................................................... 36
Figure 9: Relationship between dissolved oxygen and Mn (top), Ba and total ions (bottom), and the 222Rn/226Ra activity ratio. ......................................................................................................... 37
Figure 10: Apparent KD across dissolved oxygen and Mn (top) and Ba and total ions (bottom). Steps represent median 222Rn/226Ra in each of 10 concentration intervals containing approximately equal numbers of samples..................................................................................... 38
Figure 11: Map of Willcox Basin showing the basin’s extent, bedrock types in the mountain blocks, sample numbers, and 87Sr/86Sr ratios. Bedrock types derived from Richard et al. (2002); location of Apache Pass Fault Zone derived from Drewes (1985, 1986) and Erickson (1988). Approximate extent of cones of depression and groundwater flow directions derived water level contours in Oram (1993). Inset shows location of Willcox Basin in the Basin and Range region (USGS, Principal aquifers of the 48 conterminous United States, Hawaii, Puerto Rico, and the U.S. Virgin Islands, obtained from http://www.nationalatlas.gov). .............................................. 55
Figure 12: Hydrogen and oxygen isotope ratios in Willcox Basin groundwater. Dashed line represents the global meteoric water line (GMWL; δ2H = 8 × δ18O + 10‰; Craig, 1961). ............ 65
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Figure 13: Map showing δ18O of water samples in relationship to basin topography. Elevation data obtained from U.S. Geological Survey (http://seamless.usgs.gov). ...................................... 66
Figure 14: Relationships between 87Sr/86Sr and oxyanion‐forming trace element concentrations and Sr concentration......................................................................................................................67
Figure 15: Relationship between Ca/Na and 87Sr/86Sr and arsenic concentration. ....................... 68
Figure 16: δ11B in relation to B concentration, 87Sr/86Sr, and As concentration. Note the apparent separation between trend 1 plus the mountain blocks (low B concentration, high δ11B, high 87Sr/86Sr, low As) and trend 2 (high B concentration, low δ11B low 87Sr/86Sr, elevated As). Black bar represents the approximate δ11B range of rhyolite and granite (Barth, 1993)....................... 68
Figure 17: δ34SSO4 in relation to sulfate concentration, SO42‐/Cl‐ ratio, and Ca/SO4
2‐ ratio. ...........70
Figure 18: Map indicating δ 34S in relation to the possible gypsum source area in the Dragoon Mountains and the extent of Pleistocene lake deposits. .............................................................. 80
Figure 19: Schematic cross‐section of a generalized alluvial basin (modified from Robertson, 1991) depicting influences on observed Sr, S, B, O, and H isotope ratios and the Ca/Na ratio. ...87
Figure 20: Extent of Jordan‐Prairie du Chien aquifer system, confining units, and tritium activities in the Jordan aquifer and overlying Ordovician Prairie du Chien and St. Peter aquifers. Purple line represents extent of Prairie du Chien‐Jordan aquifer system derived from map data (Mossler and Book, 1984; Mossler and Bloomgren, 1990; Runkel, 1998; Mossler and Tipping, 2000; Mossler, 2001) or inferred from well logs. Tritium data are from published reports (Alexander and Alexander, 1989; Tipping, 1992, 1994; Wall and Regan, 1994; Zhang and Kanivetsky, 1996; Campion, 1997, 2002; Berg and Bradt, 2003; Petersen, 2005) and unpublished data from Minnesota Department of Health, Minnesota Geological Survey, and Minnesota Pollution Control Agency. Decorah confining unit extent is from Morey and Meints (2000); Des Moines till extent is from Hobbs and Goebel (1982)....................................................................................... 92
Figure 21: Cross‐section along line A‐A’ (Figure 20) indicating hydrogeologic setting, open intervals (box lengths), tritium activities (in tritium units, indicated by box colors), 87Sr/86Sr (black numbers), and 226Ra activity in mBq L‐1 (red numbers). Heavy green line represents land surface elevation; city names are discussed in text. See Table 12 for hydrostratigraphic unit abbreviations. ................................................................................................................................99
Figure 22: 87Sr/86Sr ratios in relation to 1/Sr, Sr/Ca, and Sr/Mg ratios and K concentration. Dashed line represents late Cambrian marine carbonates (McArthur et al., 2001). ..................100
Figure 23: Map showing radium‐226 activities in Jordan Aquifer groundwater and locations of wells from which cuttings were analyzed.................................................................................... 102
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Figure 24: Map showing 228Ra/226Ra activity ratios of Jordan aquifer groundwater. ..................103
Figure 25: Radium isotope activities (224Ra, 226Ra, 228Ra) plotted in relationship to each other. Line marked “MCL” indicates the EPA maximum contaminant level. ......................................... 104
Figure 26: Plots showing Ra and U isotope activities in aquifer solids and relationship between radium‐226 in wells analyzed for groundwater and solids. Groundwater radium‐226 data for well 151582 are from Lively et al. (1992). ........................................................................................... 105
Figure 27: Relationship between 228Ra/226Ra and 87Sr/86Sr in Jordan aquifer groundwater........109
Figure 28: Radium‐226 versus radium‐228 in regional Cambrian‐Ordovician aquifers. Data were obtained from this study, Lively et al. (1992), and the Minnesota Drinking Water Information System database, with other sources noted in figures. Prairie du Chien‐Jordan multi‐aquifer wells are plotted with the Prairie du Chien aquifer; some Mt. Simon aquifer wells include underlying Proterozoic sandstone aquifers. ................................................................................ 111
Figure 29: Plot of Ba vs. Ra molar concentration. Wells with anomalously high Ba/Ba are labeled by their unique numbers. ............................................................................................................ 114
Figure 30: Inset maps across the central Twin Cities metropolitan area, where radium concentrations are higher, showing 226Ra (top), Ba concentration (middle), and Ra/Ba ratio (bottom). Data are from samples in the Jordan sandstone aquifer and overlying Ordovician Prairie du Chien and St. Peter aquifers. Dashed lines represent inferred breaks from low to high concentrations of Ra and Ba. Note that the geographic patterns of Ra and Ba do not exactly coincide. .......................................................................................................................................115
Figure 31: Map showing sampling locations and chloride concentrations in Agadir area, western Morocco. Unlabeled square data points are from Tagma (2005). Base map content obtained from ESRI ArcGIS Online (World Topographic Map, World Terrain Base, and World Reference Overlay)........................................................................................................................................136
Figure 32: Plot showing relationship between bromide and chloride concentration. Dashed line represents seawater ratio............................................................................................................ 142
Figure 33: Boron concentrations in relationship to sodium and chloride concentrations. Dashed lines represent hypothetical dilution of seawater....................................................................... 143
Figure 34: δ18O plotted against δ2H of Agadir area groundwater. GMWL denotes the global meteoric water line, δ2H=8 x δ18O + 10‰ (Craig, 1961).............................................................. 144
Figure 35: Sulfur and oxygen isotope ratios of the sulfate ion for selected water samples. Lines indicate hypothetical two‐component mixing curves among possible sulfate sources. Box
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marked “gypsum” represents the range of δ34SSO4 and δ18OSO4 in Mesozoic marine sulfate
minerals (Claypool et al., 1980; Strauss, 1997)............................................................................ 145
Figure 36: Histogram showing 87Sr/86Sr values of groundwater in relation to potential Sr sources. Bars represent ranges of marine 87Sr/86Sr values (McArthur et al., 2001) of rocks encountered in the study area or the adjacent Atlas Mountains. ........................................................................ 145
Figure 37: Plots depicting radium‐224, ‐226, and ‐228 activities in relation to chloride, sulfate, and each other. Circled data points represent samples with 87Sr/86Sr > 0.709. ..........................146
Figure 38: Relationship between dissolved uranium and radium‐226 and sulfate concentration......................................................................................................................................................150
Figure 39: Relationship between chloride concentration and the (Ca+Mg)/HCO3‐ ratio. Note that
the trend extrapolates to (Ca+Mg)/HCO3‐ = 0.5 at zero salinity. ................................................. 150
Figure 40: Indicators of ion exchange in relation to each other and 87Sr/86Sr. Data indicated by x symbols are from Tagma (2005). ................................................................................................. 155
Figure 41: Map showing measured values and inferred contours of 87Sr/86Sr ratios. Base map content obtained from ESRI ArcGIS Online (World Topographic Map base map). ..................... 156
Figure 42: Sulfate and Ba2+ ion activities of Agadir groundwater. Lines represent barite saturation index values based on solubility product at 25°C. ..................................................... 158
Figure 43: Relationship between Ra2+ thermodynamic activity (derived from 226Ra) and Ba2+ and SO4
2‐ activities. .............................................................................................................................163
Figure 44: Relationship between anion concentrations and 228Ra/226Ra activity ratio. Circled points represent three wells with distinct 87Sr/86Sr > 0.709. ....................................................... 165
Figure 45: Schematic cross section of the Atlantic Coastal Plain aquifer system in North Carolina. Dashed polygons represent the approximate extent of sampling in the Cape Fear and Yorktown aquifers (modified from Winner and Coble (1996), plate 7). ...................................................... 172
Figure 46: Map showing the extent of the Yorktown, Upper Cape Fear, and Lower Cape Fear aquifers and well locations (data from North Carolina Division of Water Resources shapefiles)......................................................................................................................................................173
Figure 47: Historic water levels in Cape Fear aquifer wells sampled at this study and selected other Cape Fear wells at the same well nests. Data provided by North Carolina Division of Water Resources (http://www.ncwater.org). ........................................................................................ 174
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Figure 48: Ca, Mg, and Sr plotted relative to chloride, grouped by aquifer and water type. Dashed lines indicate seawater ion ratios. For clarity, the four shallowest samples are excluded from plots in the left two columns. ............................................................................................. 180
Figure 49: Sulfate concentration plotted relative to chloride concentration, grouped by aquifer. Dashed line represents the seawater ratio.................................................................................. 181
Figure 50: Plot of bromide vs. chloride concentrations in Cape Fear and Yorktown aquifer groundwater. Dashed line represents seawater ratio................................................................. 181
Figure 51: Ba and 226Ra plotted relative to chloride, grouped by aquifer and water type. Dashed line indicates seawater ratio. For clarity, the four shallowest wells are excluded from plots in the left two columns. .........................................................................................................................182
Figure 52: Environmental isotope ratios (strontium, boron, and sulfur) in Cape Fear vs. Yorktown aquifer groundwater. Dashed lines represent modern seawater values. ................................... 182
Figure 53: Plots showing boron concentrations and isotope ratios. ........................................... 183
Figure 54: δ34SSO4 plotted against δ18OSO4 for Yorktown and Cape Fear aquifer groundwater....184
Figure 55: Relationship between 87Sr/86Sr and 228Ra/226Ra. Labeled bands represent the 87Sr/86Sr ratios of Cretaceous and Pliocene marine carbonate (McArthur et al., 2001). Note that most Yorktown aquifer samples exhibit Pliocene marine 87Sr/86Sr and carbonate‐like median 228Ra/226Ra < 1; most Cape Fear samples exhibit 87Sr/86Sr modified by radiogenic continental inputs and sandstone‐like median 228Ra/226Ra > 1....................................................................... 192
Figure 56: Relationships between radium isotope (224Ra, 226Ra, 228Ra) activities and chloride concentrations. ............................................................................................................................197
Figure 57: Calculated chemical activities of Ra2+, RaSO40, and RaCl+ along gradients of sulfate
activity..........................................................................................................................................199
Figure 58: Ratio of sulfate‐complexed (uncharged) to divalent alkaline earth metal activities with increasing sulfate activity. ........................................................................................................... 199
Figure 59: Radium and barium in relationship to barite saturation. Dashed line represents barite saturation based on Ksp = 10
‐9.97 (Drever, 1997). Circled points are the shallowest two wells in each aquifer. ................................................................................................................................201
Figure 60: Differences in radium isotope ratios (228Ra/226Ra and 224Ra/228Ra) between the Yorktown and Cape Fear aquifers................................................................................................ 204
xxi
Figure 61: Relationship between total dissolved solids (TDS) and 226Ra activity in sand aquifers in this study. Dashed line is a linear regression of the Atlantic Coastal Plain waters with TDS > 103.5 mg L‐1 (slope ~15 mBq L‐1 per 1000 mg L‐1 TDS)........................................................................... 208
xxii
Acknowledgements
Throughout this research, many colleagues at Duke University provided valuable
collaboration and assistance. Avner Vengosh, my advisor in the Division of Earth and Ocean
Sciences, originally suggested an exploration of radium’s connections to groundwater
salinization and served as a constant source of guidance on conducting the research,
interpreting results, strengthening the research using a multi‐tracer approach from
complementary isotope systems, writing, and publication. He also secured funding without
which this research would not have been possible. Gary Dwyer provided frequent guidance on
major element, trace element, and isotope analysis, provided many hours of instruction in the
laboratory, and was a significant contributor to the data and interpretations presented in this
dissertation. Additional laboratory assistance at Duke University was provided by Hadas Raanan,
Nathaniel Warner, Laura Ruhl, Tong Zhang, Jonathan Karr, Paul Heine, Wes Willis, Heileen Hsu‐
Kim, Emily Klein, and Paul Baker. John Diemer of the University of North Carolina at Charlotte
allowed access to rock crushing equipment. Helpful discussions on isotope analysis with W.S.
Moore, Ralph Arnold, Matt Charette, Zoltan Szabo, and Thomas Bullen also significantly
improved this research. I was supported by the Nicholas School of the Environment at Duke
University through teaching and research assistantships, with partial summer support from the
North Carolina Division of Water Quality and Minnesota Department of Health. Sulfur isotope
analyses in Chapters 4 through 6 were supported by an International Association of
Geochemistry student research grant with the collaboration of Bernhard Mayer and Steve
Taylor of the University of Calgary. In addition to their roles mentioned above, Avner Vengosh,
xxiii
Paul Baker, Gary Dwyer, and Heileen Hsu‐Kim served as the PhD dissertation committee, and
generously offered their time and constructive suggestions throughout this research. Finally, the
research reported in this dissertation would not have been possible without the support and
patience of my wife, Christa Wagner Vinson. Other assistance specific to individual papers
within this research is acknowledged below.
Chapter 2: Numerous students and faculty at Duke University facilitated the large
private well sampling program. Lori Bennear, Elizabeth Pratson, and Soumya Balasubramanya
scheduled sampling events. Sampling and laboratory assistance was provided by Nick Baker,
Timothy Foley, Scott Geis, Nicole Hagan, Waruntorn Kanitpanyacharoen, James Leeper, Anna
Grzebien, and Kerry Schlichting. Phil Bradley and Michael Medina of the North Carolina
Geological Survey; Douglas Harned, Melinda Chapman, and Kristen McSwain of the U.S.
Geological Survey; and Richard Bolich of the North Carolina Division of Water Quality provided
access to unpublished data and core samples. Ted Campbell, also of the North Carolina Division
of Water Quality, was a source of many helpful discussions on radon in groundwater. GIS
assistance was provided by Joshua Tootoo of Duke University Children's Environmental Health
Initiative. This research was funded by U.S. Department of Agriculture Cooperative State
Research Education and Extension Service grant NCW‐2006‐03956. Two anonymous journal
reviewers provided constructive comments and Chemical Geology Guest Editors Jennifer
Roberts and Annette Summers Engel provided editorial handling for the published version of
this research.
Chapter 3: This research was the result of a collaboration with Jennifer McIntosh of the
University of Arizona Department of Hydrology and Water Resources, who identified the Willcox
xxiv
Basin as a location for the research, arranged sampling, participated in sampling excursions,
financially supported research costs, arranged elemental and isotopic analyses at the University
of Arizona, and participated in interpretation of the results. Stephen Osborn of the University of
Arizona also provided field assistance in the Willcox Basin. Tritium analyses were generously
provided by the University of Miami Tritium Laboratory. The quality of this research was also
significantly improved by the suggestions of Applied Geochemistry reviewers Thomas Bullen,
Alan Fryar, and Carol Frost, Associate Editor Russell Harmon, and Editor Ron Fuge.
Chapter 4: This research was the result of a lengthy collaboration with James Lundy of
the Minnesota Department of Health (MDH), Environmental Health Division, who obtained
access to sampling sites, sampled many of the waters presented in this study, arranged
additional analysis at the Minnesota Public Health Laboratory, and provided valuable knowledge
on regional hydrogeology and the occurrence of radium. This research would not have been
possible without MDH’s long‐term commitment to provide this large amount of logistical
assistance. Many helpful discussions on the hydrogeology of southeastern Minnesota with
Robert Tipping, Anthony Runkel (Minnesota Geological Survey); Bruce Olsen, Sheila Grow, Justin
Blum, Richard Soule, Tom Alvarez, Steve Robertson (Minnesota Department of Health); and E.
Calvin Alexander (University of Minnesota) greatly improved the research. Alan Knaeble of the
Minnesota Geological Survey provided access to samples of cuttings from Jordan aquifer wells.
Chapter 5: This research was done in collaboration between the faculty and students of
Duke University and Universite Ibn Zohr in Agadir, Morocco, including Tarik Tagma, Zineb
Lgourna, Najat Ettayfi, Nathaniel Warner, and Lhoussaine Bouchaou. The research was funded
xxv
by the North Atlantic Treaty Organization Science for Peace Program, Project EPS.Md SFP
983134.
Chapter 6: This research was made possible by logistical and sampling assistance from
personnel of the Dare County Water Department and the North Carolina Division of Water
Resources, especially Barbara Peck. The sampling of Dare County Water Department wells was
conducted with Haylee Schwartz of Duke University.
1
1. Introduction
1.1 Overview of radium isotope geochemistry
Four naturally‐occurring radium isotopes are present in nature, all derived from alpha
decay of thorium parents within the thorium‐232, uranium‐235, and uranium‐238 decay series
(Figure 1). The half‐lives of these isotopes occupy a wide range: 3.6 d for 224Ra, 11.4 d for 223Ra,
5.8 a for 228Ra, and 1600 a for 226Ra (Table 2). On a mass basis, more than 99% of natural radium
is present as 226Ra. However, on an activity (decay rate) basis, 224Ra, 226Ra, and 228Ra are about
equally abundant in the crust because both 232Th and 238U are also equally abundant on an
activity basis. Due to the low abundance of 223Ra, controlled by the natural 235U/238U activity
ratio (0.046 at secular equilibrium), it is of less concern for water quality and difficult to analyze
at high precision. Consequently, this research is focused on the abundances of 224Ra, 226Ra, and
228Ra in water resources and the geochemical mechanisms that control these variations.
1.2 Health effects and drinking water standards
In addition to understanding geochemical processes, one of the motivations for studying
radium abundances in water resources is the potential effects on human health. The health
effects of radium come not from chemical toxicity but from radioactivity. The primary route for
radiation exposure from radium is to the bones, where radium accumulates due to its chemical
similarity with calcium (Mays et al., 1985; Environmental Protection Agency, 1999). Because of
the elevated risk of bone cancer imparted by radium, the US Environmental Protection Agency
regulates the combined activity of 226Ra + 228Ra at the level of 5 picocuries per liter (pCi L‐1),
equivalent to 185 millibecquerels per liter (mBq L‐1) in the SI system (Table 1). Short‐lived
radium‐224, although a significant part of the gross alpha radioactivity of waters (Parsa, 1998), is
2
not regulated in the United States but is recommended for further study (Environmental
Protection Agency, 2000). In case of high levels of radium in drinking water, treatment by
softening (cation exchange) is effective, but concentrates radium in the brine, creating other
potential environmental and disposal concerns (Szabo et al., 2008, 2010).
Figure 1: Simplified 238U (black), 235U (green), and 232Th (blue) decay series including uranium, thorium, radium, and radon isotopes (modified from Molinari and Snodgrass, 1990).
The solid‐phase distribution of radium is also a major control on the occurrence of radon
isotopes. Radon‐222 is a chemically inert, short‐lived, alpha‐emitting gas that is of health
concern in both indoor air and drinking water due to the lung cancer risk that results from
inhalation of alpha‐emitting radon decay products. The largest source of radon exposure to the
population is derived from soil gas and inhaled indoors. In contrast, radon degassed from
drinking water and inhaled is a smaller health risk to the overall population than radon derived
from soil gas, but is a significant health risk compared to other regulated waterborne substances
(National Research Council, 1999; Environmental Protection Agency, 2003). In the United States,
3
drinking water standards have been proposed (11 Bq L‐1 or alternative limit of 148 Bq L‐1 if
indoor air mitigation measures are in place; Table 1), but have not been enacted.
Table 1: United States and international standards for radon and radium in drinking water.
Isotope US Environmental Protection Agency standard
World Health Organization guideline
224Ra Not regulated 1000 mBq L‐1 226Ra 1000 mBq L‐1 228Ra
185 mBq L‐1 combined 226Ra+228Ra 100 mBq L‐1
222Rn 11‐148 Bq L‐1 (proposed) 100 Bq L‐1 Reference Environmental Protection
Agency, 2000 World Health Organization, 2008
1.3 The link between radium abundance and hydrogeochemical change
Because radium is a naturally‐occurring element broadly distributed across aquifer
materials, high radium in groundwater is not typically a result of radium transport. In fact, the
large majority of radium in an fresh water system remains on the solids. Elevated levels of
radium in groundwater are the result of solid‐phase radioactivity levels, the primary control, as
modified by groundwater chemistry. This overall balance is described, in the aggregate, by
adsorption and desorption rate parameters, distribution coefficients, and retardation constants
(e.g. Krishnaswami et al., 1982). This information is useful for radionuclide transport rates under
steady‐state conditions, but provides fewer insights about how the system may respond to
chemical change such as salinization or modification of redox conditions. One of the motivations
of this research is to better understand the sensitivity of radium in groundwater systems to
hydrogeochemical changes. For example, recent research in New Jersey demonstrated that the
nitrification of agricultural waste combined with high rates of groundwater pumping can lead to
higher levels of radium in water by decreasing the pH of a shallow acidic aquifer (Szabo et al.,
4
2005; Ayotte et al., 2011). This research aims to evaluate to what extent salinization and redox
gradients can similarly induce increases in radium in groundwater.
1.4 Modeling radium mobilization in groundwater systems
At steady‐state radium concentration, radium source terms are balanced by removal
terms:
in which P (in atoms per unit water volume per unit time) is the supply rate of the radium
isotope from all sources including alpha recoil (section 1.5.1), dissolution (section 1.5.2), and in
situ production from parent isotopes in solution (negligible for radium); iRas is the solid‐phase
(adsorbed) concentration of the radium isotope (in atoms per unit of water volume in contact
with the solid); iRaaq is the aqueous radium isotope atomic concentration (in atoms per unit of
volume); λRa‐i is the decay constant of the radium isotope, k1 is the first‐order rate constant of
radium removal; and k2 is the first‐order rate constant of desorption (Krishnaswami et al., 1982).
In this model of radium mobilization, the primary mechanism of radium removal is adsorption;
Ra‐Ba solid solution in sulfate minerals (coprecipitation; section 1.6.3) is neglected.
1.5 Radium source mechanisms
1.5.1 Alpha recoil
In nearly all groundwater systems, the primary source mechanism for radium in water is
alpha recoil, in which the energy associated with the alpha decay of the parent thorium isotopes
ejects the radium daughter products into solution. The rate of the alpha decay input from long‐
lived Th parents to shorter‐lived Ra daughters is governed by the solid‐phase atomic
concentration of the parent isotope and the half‐life of the daughter. According to the law of
(1)
5
radioactive ingrowth, about 5 half‐lives of the daughter are required for the activity of the
daughter to reach 95% equilibrium with the activity of the parent isotope. In porous systems,
the basic concepts of radioactive ingrowth are modified significantly: (1) parent nuclides must
be within recoil distance of the edge of the mineral (30‐40 nm; Sun and Semkow, 1998); (2)
approximately half of radionuclides within recoil distance of the mineral edge undergo recoil
into the water rather than the opposite direction, into the solid (Davidson and Dickson, 1986);
and (3) the extent and effective time scale of daughter product accumulation are significantly
affected by adsorption (Krishnaswami et al., 1982; Davidson and Dickson, 1986).
1.5.2 Desorption and dissolution
Compared to alpha recoil, net desorption and dissolution of radium sources in the
aquifer solids are negligible for short‐lived isotopes of radium, although the longest‐lived
isotope 226Ra may receive minor weathering inputs (Krishnaswami et al., 1982). This is because
the short‐lived isotopes decay rapidly relative to dissolution and desorption, whereas these
same isotopes are readily added to water by alpha recoil when their parents undergo decay. In
unusual situations, desorption or mineral dissolution are major radium sources to water,
including extremely rapid salinization from freshwater to brine levels of salinity (Wood et al.,
2004) and in saline geothermal systems (Zukin et al., 1987; Hammond et al., 1988). These result
in unusual radium compositions enriched in long‐lived radium isotopes while in the more typical
recoil‐adsorption scenarios, activities of the short‐lived Ra isotopes equal or exceed the
activities of the long‐lived Ra isotopes.
6
1.6 Radium removal mechanisms
1.6.1 In situ radioactive decay
In the oceans, where water‐rock interaction is minimal, the saline nature of seawater
supports radium as a semi‐conservative tracer in which its losses are assumed to be mostly to
radioactive decay in the water column. Radium isotopes can be used to estimate fluxes of
submarine groundwater discharge into the coastal ocean (e.g. Moore, 1996; Gonneea et al.,
2008; Lamontagne et al., 2008). In groundwater systems where water‐rock interaction is much
greater, the probability that radium will be chemically removed prior to its own decay is very
high, that is, k1 > λRa‐224. For this reason, the in situ radioactive decay term iRaaqλRa‐i in equation 1
is negligible in groundwater systems. The implication of this well‐accepted assumption is that
radium isotopes are not particularly applicable as tracers of flow velocity and residence time in
groundwater systems due to the high level of water‐rock interaction in aquifers. Instead, radium
isotopes in groundwater systems are sensitive tracers of water‐rock interaction.
1.6.2 Adsorption and cation exchange
Radium adsorption sites are abundant in most groundwater systems, and adsorption is
rapid, leading to a distribution of radium that is dominated by the solid phase. Effective radium
adsorption sites include Mn and Fe oxides (Ames et al., 1983c; Nirdosh et al., 1990; Sverjensky,
2006), although Mn oxides seem to be far more effective than Fe oxides (Moore and Reid,
1973); clays and zeolites (Nathwani and Phillips, 1979a; Ames et al., 1983a; Komarneni et al.,
2001; Tachi et al., 2001; Hidaka et al., 2007; Shao et al., 2009); micas (Ames et al., 1983b); and
organic matter (Nathwani and Phillips, 1979a). Given the salinity‐ and redox‐sensitive nature of
these adsorption sites, the removal of Ra by adsorption should be more efficient in fresh, oxic
7
waters. In general, adsorption of radium refers to inner‐sphere surface complexes, whereas
cation exchange refers to much more reversible sites in outer‐sphere complexes or the diffuse
ion swarm (Sposito, 2008), but are sometimes discussed interchangeably within the context of
radium removal mechanisms.
1.6.2.1 Salinity effects
The first‐order rate constant of adsorption k1 may vary by approximately 5 orders of
magnitude between freshwater and brine levels of salinity (Table 2; Krishnaswami et al., 1982,
1991). In field‐based studies, increasing radium levels with groundwater salinity are often
observed (Tanner, 1964; Bloch and Key, 1981; Kraemer and Reid, 1984; Herczeg et al., 1988;
Krishnaswami et al., 1991; Moise et al., 2000; Sturchio et al., 2001; Tomita et al., 2010).
Potential mechanisms include ionic strength (the increasing ion‐ion interactions at high salinity)
and complexation with chloride (RaCl+) or sulfate (RaSO40), which may be more soluble than Ra2+
in some environments. Limited evidence suggests that radium levels are associated with waters
dominated by divalent cations rather than monovalent cations, as indicated by Ca concentration
and/or the Ca/Na ratio (Nathwani and Phillips, 1979b; Langmuir and Melchior, 1985; Zukin et al.,
1987; Gascoyne, 1989; Kitto et al., 2005; Shao et al., 2009).
Table 2: First‐order rate constants for the terms in equation 1.
Parameter Half‐life Rate constant (s‐1) Reference λRa‐224 3.6 d 2.2 x 10‐6 λRa‐223 11.4 d 7.0 x 10‐7 λRa‐228 1600 yr 3.9 x 10‐9 λRa‐226 5.8 yr 1.4 x 10‐11
National Nuclear Data Center, NuDat 2.5 at http://www.nndc.bnl.gov/nudat2
k1 10‐5‐10‐1 k2 10‐6‐10‐5
Krishnaswami et al. (1982, 1991)
8
1.6.2.2 pH effects
Adsorption of radium is much less effective in acidic groundwaters because of the
increasing surface charge of Ra2+‐adsorbing minerals as pH decreases. For example, the point of
zero charge of Mn oxide is approximately 5, with positive surface charge below that pH value
(Stumm and Morgan, 1996). This is a documented cause of high‐Ra conditions in near‐surface,
acidic, low‐TDS groundwater in aquifer material of little or no pH buffering capacity (Cecil et al.,
1987; Dickson and Herczeg, 1992; Kozinski et al., 1995; Senior and Vogel, 1995; Bolton, 2000;
Lauria et al., 2004; Szabo et al., 2005; dePaul and Szabo, 2007). In some cases, the nitrification of
agricultural waste can exacerbate the low‐pH condition and significantly increase radium levels
(Szabo et al., 2005).
1.6.2.3 Redox effects
Redox effects on radium mobility have been observed qualitatively by correlations of
radium with dissolved oxygen, manganese, or iron concentrations. These effects seem to be due
to the redox‐sensitivity of Mn and Fe oxides because radium itself is not a redox‐sensitive
element. Redox effects have been observed in both fresh (Tanner, 1964; Felmlee and Cadigan,
1979; Szabo and Zapezca, 1987; Copenhaver et al., 1993) and saline waters and brines (Zukin et
al., 1987; Herczeg et al., 1988; Gonneea et al., 2008). When metal oxides undergo reductive
dissolution, the efficiency of radium adsorption is diminished significantly, permitting high
radium activities in groundwater. In addition to the redox‐sensitivity of Ra adsorption sites, at
somewhat more reducing conditions corresponding to progressive organic carbon oxidation
coupled to inorganic electron acceptors, sulfate‐reducing conditions can prevent barite
precipitation, which is a radium removal mechanism.
9
1.6.3 Coprecipitation into sulfate minerals
Groundwater Ra activities are commonly observed to be correlated with Ba
concentrations. Ba concentrations (Underwood et al., 2009) and Ra activities (Dickson, 1985;
Kitto et al., 2005) have been noted to be negatively correlated with increasing sulfate
concentration, which also implies that sulfate is linked to a net removal mechanism and
contradicts the expected association between Ra and overall salinity. However, the levels at
which an apparent sulfate control is observed are orders of magnitude undersaturated with
respect to RaSO4 so that the actual mechanism of Ra removal is Ba‐Ra solid solution
(coprecipitation) in barite, BaSO4, rather than pure radium salts (Langmuir and Melchior, 1985).
Barite precipitation occurs in a complex solid solution incorporating Ba2+, Sr2+, Pb2+, and Ra2+
(Zhu, 2004a), so effects from competing cations on the extent of Ba and Ra removal may be
significant. The presence of other competing cations (especially Sr2+) can diminish the radium
content of barite (Ceccarello et al., 2004; Shao et al., 2009). With significant changes in the
redox state of a system, reductive dissolution can release Ra that had been sequested in barite
and can prevent additional barite from precipitating (Phillips et al., 2001; Landa, 2003; Martin et
al., 2003). While barite precipitation is regarded as rapid enough to substantially limit radium in
groundwater systems (Zhu, 2004b; Curti et al., 2010), the kinetics of this process are significantly
more complicated than first‐order rate models. The rate of barite nucleation is higher from
more supersaturated solutions (Nancollas and Reddy, 1974; Liu et al., 1976; Nordstrom and Ball,
1989), so barite precipitation is probably slow at slight supersaturation and could be less
efficient than more rapid processes such as cation exchange (Shao et al., 2009) or adsorption
(Krishnaswami et al., 1982). Thus, it is not straightforward to incorporate Ra removal into barite
into the first‐order Ra removal (adsorption) rate constant k1.
10
1.7 Controls on radium isotope ratios
Radium isotope ratios are commonly expressed as short‐lived to long‐lived isotope
ratios (SRa/LRa). Due to the small relative mass differences between the radium isotopes, there
is no known fractionation. Instead, differences in Ra isotope ratios are due to the different
decay constants of the isotopes in relationship to the rates of radium‐removing processes.
1.7.1 228Ra/226Ra
It is generally accepted that the 228Ra/226Ra ratio is closely associated with the solid‐
phase 232Th/238U ratio (Asikainen, 1981a; Andrews et al., 1989; Dickson, 1990; Sturchio et al.,
1993; Szabo et al., 1997; Tricca et al., 2001; Vengosh et al., 2009), which at its crustal average
value suggests an average 228Ra/226Ra value ≈ 1. This predicts that 228Ra/226Ra values are at or
above 1 in sandstones (Dickson et al., 1987; Lively et al., 1992; Szabo et al., 1997; Szabo et al.,
2005; Vengosh et al., 2009), and are significantly below 1 in carbonate aquifers (Hileman and
Lee, 1993; Moise et al., 2000; Sturchio et al., 2001; Condomines et al., 2010), marine phosphates
(Moise et al., 2000; dePaul and Szabo, 2007), and U‐mineralized areas (Dickson et al., 1984) due
to the preferential enrichment of U in those deposits. However, several arguments have been
advanced to explain significant deviations from the rock‐derived ratio. Low ratios may be
obtained by preferential solubility of U relative to Th concentrating the parents of 226Ra in
fractures and flow channels (Dickson et al., 1987; Rihs and Condomines, 2002); or by successive
alpha‐recoil in which 226Ra is the third alpha recoil step from 238U, but 228Ra is only the first recoil
step from 232Th, resulting in a greater recoil flux of 226Ra relative to 228Ra than would be expected
from the Th/U ratio (Davidson and Dickson, 1986). 228Ra/226Ra ratios higher than rock ratios
could be caused by preferential leaching of thorium during weathering processes (Dickson,
11
1985; Sturchio et al., 1993); rapid intrusion of high‐salinity, low‐radium seawater in which 228Ra
is replenished more rapidly by recoil during 226Ra resulting in elevated 228Ra/226Ra during the
adjustment period (Lamontagne et al., 2008; Otero et al., 2011); or millennial‐scale variation in
weathering rates in old, slow‐circulating groundwater (Reynolds et al., 2003).
1.7.2 224Ra/228Ra
The 224Ra/228Ra ratio is a short‐lived to long‐lived ratio within the 232Th decay series.
Because both isotopes undergo recoil at the same rate, but short‐lived 224Ra is replaced by recoil
more rapidly than 228Ra, and perhaps also because of successive recoil effects in which 224Ra’s
parent 228Th is more likely to be near the mineral surface than 228Ra’s parent 232Th, median
values of 224Ra/228Ra in most studies are at or slightly above 1 and rarely exceed 2 (Davidson and
Dickson, 1986; Copenhaver et al., 1993; Sturchio et al., 1993; Luo et al., 2000; Moise et al., 2000;
Focazio et al., 2001; Tricca et al., 2001; Reynolds et al., 2003; Ruberu et al., 2005). Values of
224Ra/228Ra > 2 are unusual and seem to be associated with waters subjected to rapid shifts in
solid‐phase radioactivity along the flowpath and/or very short residence time (Dickson et al.,
1987; Lamontagne et al., 2008). Values in 224Ra/228Ra significantly below 1 are associated with
mineral dissolution, but this is mainly documented in geothermal systems (section 1.5.2).
1.7.3 222Rn/226Ra and 222Rn/224Ra
Studies in which both radium and radon have been analyzed provide additional
quantification of radium’s mobility. It is commonly documented that 222Rn activities are 3 to 6
orders of magnitude higher than 226Ra in fresh water systems (Krishnaswami et al., 1982;
Copenhaver et al., 1993; Tricca et al., 2001), and are especially high in fractured crystalline rock
groundwater systems (Asikainen, 1981b; Wanty et al., 1991; Senior and Vogel, 1995). In saline
12
aquifers, the 222Rn/226Ra ratio may be on the order of ~100 (Krishnaswami et al., 1991; Moise et
al., 2000). Due to the lack of retardation along flowpaths and its common recoil‐derived source
in the 238U decay series, 222Rn has been assumed to be proportional to the recoil flux of all
isotopes in the decay series. Thus, the ratio of 222Rn to other radionuclides provides an estimate
of retardation or distribution coefficients for those radionuclides (Krishnaswami et al., 1982).
222Rn/224Ra, a variation on the 222Rn/226Ra ratio, negates half‐life differences and primarily
documents the differences in retardation between the inert Rn and the strongly‐adsorbed Ra. In
most groundwater systems, values of 222Rn/224Ra are similar to 222Rn/226Ra, corrected for relative
224Ra/226Ra abundance (Krishnaswami et al., 1982).
1.8 Approach of this research
In this research, additional tracers from complementary isotopic systems are used to
evaluate specific radium‐mobilizing mechanisms along salinity and redox gradients. These
tracers include: (1) ion ratios that relate to cation exchange including, for example, Ca/Na and
Na/Cl‐ ratios; (2) 87Sr/86Sr ratios, which record sources of divalent cation release to groundwater,
such as cation exchange from clays in response to seawater intrusion; (3) boron concentrations
and δ11B, which can also coincide with conditions in which Ca and Na are being exchanged in
response to salinization and freshening; and (4) sulfur and oxygen isotopes of sulfate, which are
used to confirm the redox‐stability of barite (BaSO4), a mineral that can remove radium by
coprecipitation. Although radium isotopes are reasonably well understood to trace water‐rock
interaction processes in field‐based and modeling studies, this multi‐tracer approach represents
a novel application in combining radium isotopes with environmental isotopes that address
specific hypotheses related to radium mobilization. In field‐based studies of groundwater along
13
salinity and redox gradients, critical mechanisms might include, for example, inefficient Ra
adsorption indicated by divalent cation release, efficient Ra adsorption or cation exchange,
and/or redox effects on Ra removal processes. The five chapters reported in this dissertation
represent a range of hydrochemical conditions: salinity ranging from very dilute to
approximately 50% of seawater salinity, redox conditions ranging from oxic to sulfate‐reducing
(Figure 2), and containing radium activities ranging across three orders of magnitude.
Chapter 2 examines radium reactivity under low‐salinity conditions in which slightly‐
acidic to neutral pH waters in fractured crystalline rocks were examined in North Carolina. These
waters exhibit total dissolved solids commonly < 500 mg L‐1 and redox conditions ranging from
oxic to Fe oxide reducing. This and other research in the saprolite‐fractured crystalline rock
groundwater system in the southeastern United States documents that redox conditions, cation
concentrations, and groundwater residence time tend to coincide, as anoxic, high‐Fe conditions
can be the result of biogeochemical activity and/or weathering of Fe2+‐bearing minerals such as
biotite. These higher‐cation, anoxic conditions exhibit the highest levels of radium in the study
and the lowest values of the 222Rn/226Ra ratio. In contrast, the lowest‐TDS, oxic waters tend to
exhibit the highest levels of 222Rn and 222Rn/226Ra. Overall, it appears that one order of
magnitude of Ra mobility variation is documented along the gradients from oxygen‐saturated to
anoxic and from low cation concentrations to high cation concentrations. This research, with
contributing authors Daniella Hirschfeld, Gary S. Dwyer, and Avner Vengosh, was published in
the journal Chemical Geology (Vinson et al., 2009).
14
Figure 2: Diagram indicating the approximate ranges of salinity and redox conditions encountered in this study.
Chapter 3 evaluates radium abundance in the Willcox Basin, an oxic basin‐fill aquifer in
Arizona with total dissolved solids ranging from very dilute to ~1200 mg L‐1. Substantial variation
in major element chemistry is apparent, but redox conditions of the waters in this study are
consistently oxic. Throughout the study area, extremely low levels of all Ra isotopes were
observed, except in the portion of the study area underlain by fractured crystalline rocks. The
application of radium in this study is to assess the efficiency of adsorption sites (e.g. Fe and Mn
oxides) that may be shared between Ra and As and other oxyanion‐forming trace elements.
Essentially, the adsorption of radium under these conditions is very effective because all of the
mechanisms of Ra removal are at their maximum efficiency under the oxic, neutral to slightly
basic pH, low salinity, divalent‐cation removing conditions documented in the Willcox Basin.
15
These redox‐pH conditions that efficiently remove radium favor the desorption of arsenate,
meaning that there is no overlap between elevated levels of oxyanion‐forming elements and
elevated Ra. This research is in collaboration with contributing author Jennifer C. McIntosh of
the University of Arizona Department of Hydrology and Water Resources.
Chapter 4 evaluates the mechanisms of high‐Ra, yet low‐salinity groundwater from the
Jordan sandstone in Minnesota, a Cambrian quartz sandstone present in Minnesota, Wisconsin,
Iowa, and Illinois. Waters are fresh, with Ca‐Mg‐HCO3‐ composition near saturation with calcite
and dolomite, exhibiting low concentrations of other dissolved ions and TDS < 1000 mg L‐1.
Redox conditions range from oxic to Fe oxide reducing, although the majority of the aquifer,
including the high‐Ra portion of the aquifer, is Fe oxide‐reducing as indicated by elevated
concentrations of Fe and Mn. The largest range of 226Ra, and the highest overall values of 226Ra
in this thesis were observed in the Jordan aquifer, despite its low salinity. Because the
radionuclide levels of the aquifer rocks are not particularly high, analysis of radium distribution
coefficients suggests that the redox condition of the aquifer contributes to inefficient Ra
adsorption. The overall range of radium levels in groundwater is a result of both the redox
conditions and the approximately one order of magnitude variation in solid‐phase radionuclide
levels. The relatively inefficient removal of radium is partially mitigated by coprecipitation of Ra
into barite. Both Ba and 226Ra seem to be influenced by a geographic distribution in the aquifer
of unknown origin. This chapter is a collaboration with contributing author James R. Lundy of
the Minnesota Department of Health.
Chapter 5 reports the radium variations under of an unconfined coastal aquifer in
Morocco that is undergoing salinization. Application of multiple geochemical and isotopic
tracers reveals several salinity sources, including high‐sulfate material from the adjacent Atlas
16
Mountains foothills, oxidized marine sulfide, and urban wastewater, and thus the conditions of
the aquifer differ from classic seawater intrusion. In general, the aquifer appears to exhibit oxic
conditions indicated by detectable dissolved oxygen, nitrate, and uranium. Indications of cation
behavior suggest that divalent cations are released by cation exchange, which does not favor Ra
removal. Radium‐224 is correlated with chloride concentration, which implies that the inhibited
adsorption by salinity is most closely linked to the short‐lived isotope’s rapid response. Radium
is probably strongly removed by coprecipitation into barite. This removal mechanism is best
correlated with long‐lived 226Ra, which is consistent with the relatively low rate of 226Ra
replacement by alpha recoil. This research was conducted with contributing authors Tarik
Tagma and Lhoussaine Bouchaou of Universite Ibn Zohr.
Chapter 6 examines radium variations in anoxic, sulfate‐reducing, fresh to saline
Cretaceous and Pliocene aquifers of the Atlantic Coastal Plain in North Carolina. This aquifer
system represents a well‐studied ion exchange environment in which divalent cations are
rapidly exchanged for Na that saturates the marine‐derived clays in the aquifer solids. Along a
salinity gradient of chloride concentration reaching up to ~10,000 mg L‐1 (TDS ~18,000 mg L‐1),
radium activities are correlated with chloride concentration, but in the few fresh samples
containing higher proportions of divalent cations, indicating weaker cation exchange, Ra exhibits
even higher levels relative to the saline groundwater. Given the strongly reducing conditions
and their apparent interference with Ra removal by adsorption and coprecipitation, it seems
that a rapid and highly efficient Ra removal process is cation exchange onto the aquifer solids,
which may explain the relatively low levels of Ra in the aquifer. This study also shows that the
association of Ra with salinity can be enhanced by the presence of the uncharged RaSO40
complex in high‐sulfate waters.
17
2. Relationships between radium and radon occurrence and hydrochemistry in fresh groundwater from fractured crystalline rocks, North Carolina (USA)
2.1 Introduction ∗
Groundwater sources of drinking water may contain naturally‐occurring radionuclides
from the uranium‐238, uranium‐235, and thorium‐232 series, especially uranium (U), radium
(Ra), and radon (Rn), all of which are the basis of current (U, Ra) or proposed (Rn) drinking water
standards of the U.S. Environmental Protection Agency (EPA). Uranium is a source of kidney
toxicity as well as alpha radiation (EPA, 2000). Radium is a bone‐seeking form of natural
radioactivity when consumed in drinking water (Mays et al., 1985). Radon is a radioactive gas for
which both ingestion and inhalation present health risks, although inhalation may provide about
one order of magnitude larger exposure (National Research Council, 1999). U, Ra, and Rn may
co‐occur in high concentrations in the case of mineralized areas (e.g. Asikainen and Kahlos,
1979), but in most cases, strong disequilibrium exists among U, Ra, and Rn in the groundwater
of crystalline rocks. Disequilibrium is observed because of geochemical conditions that
∗ ‐ This chapter was published as Vinson, D.S., Vengosh, A., Hirschfeld, D., and Dwyer,
G.S., 2009, Relationships between radium and radon occurrence and hydrochemistry in fresh
groundwater from fractured crystalline rocks, North Carolina (USA): Chemical Geology, v. 260,
no. 3‐4, p. 159‐171, doi:10.1016/j.chemgeo.2008.10.022.
*
18
preferentially mobilize U and/or Ra as well as the inert nature of Rn. For example, waters may
be high in U but low in other nuclides (Asikainen, 1981b); orders of magnitude higher in Rn than
in Ra (King et al., 1982; Wanty et al., 1991); or high in Ra but relatively low in U in deep, anoxic
waters (Andrews et al., 1989; Gascoyne, 1989).
In addition to being of significance as natural drinking water contaminants, Ra isotopes
are useful tools to evaluate the behavior and reactivity of radionuclides in aquifer systems. This
is due to (1) the large range of half‐lives (3.7 d for 224Ra, 11.4 d for 223Ra, 5.8 a for 228Ra, and 1600
a for 226Ra); and (2) the highly reactive nature of Ra, especially used in conjunction with the inert
nature of radon‐222 (half‐life 3.8 d), the daughter of 226Ra. In fresh waters, the relationships
between Ra isotopes and between Ra and Rn suggest that Ra is generated primarily by alpha
recoil from the decay of its Th parents on the aquifer solids, rather than by weathering
processes, and that transport of Ra along a flowpath is retarded several orders of magnitude
relative to conservative solutes (Krishnaswami et al., 1982; Dickson, 1990; Porcelli, 2008). Thus,
Ra isotope ratios and Rn/Ra ratios can indicate the balance between physical processes (alpha
recoil) and geochemical processes such as Ra adsorption.
Radium mobility or adsorption effectiveness is understood to vary with chemical
parameters including pH (Cecil et al., 1987; Dickson and Herczeg, 1992; Bolton, 2000; Szabo et
al., 2005), salinity (Kraemer and Reid, 1984; Sturchio et al., 2001; Wood et al., 2004), reduced
conditions (Szabo and Zapezca, 1987; Herczeg et al., 1988), supersaturation with respect to
barite (Gilkeson et al., 1984; Grundl and Cape, 2006), microbial sulfate reduction affecting barite
stability (Phillips et al., 2001; Martin et al., 2003), and microbial Fe oxide reduction (Landa et al.,
1991). Ra adsorption is a rapid influence in freshwater at near‐neutral pH (Krishnaswami et al.,
1982), but the effects of redox processes and relatively low ion concentrations on adsorption
19
and desorption are not well constrained in existing field investigations. In addition, many
previous studies of Ra in fresh waters have not included detailed characterization of redox
conditions or analysis of short‐lived Ra isotopes with which to apply the full range of Ra half‐
lives to geochemical conditions. In this research, we focus chemical and Ra isotope investigation
on fresh, near‐neutral, bicarbonate‐dominated groundwaters from fractured crystalline rocks in
North Carolina (USA). The objectives of the study are to evaluate mechanisms of Ra and Rn
mobilization in shallow groundwater and to elucidate the hydrochemical influences, if any, on
observed Ra and Rn activities under these conditions.
2.2 Methods
2.2.1 Sampling procedures
117 private wells were sampled in Wake County, North Carolina, clustered into six
groups (Figure 3): (1a) and (1b) are a complexly mapped area mostly composed of late
Proterozoic to early Paleozoic metasedimentary and metavolcanic rocks (felsic gneiss, felsic
schist, phyllite, meta‐tuff, and scattered ultramafic rocks) and associated felsic intrusions
including quartz diorites (Stoddard et al., 1991; Clark et al., 2004); (2) is primarily underlain by
the early Paleozoic Raleigh Gneiss, composed of sedimentary and granitic components
(Stoddard et al., 1991; Clark et al., 2004); and (3a), (3b), and (3c) are underlain by the late
Paleozoic Rolesville granitic pluton, a large, complex unit composed of monzogranite, granite,
granodiorite and pegmatite in places (Speer, 1994; Clark et al., 2004). The study area is also
intruded by Mesozoic dikes of mafic composition (Clark et al., 2004).
20
Figure 3: Map of the study area showing the six sample groups. Numbers in parentheses are the number of sites in each group; political boundary shown is of Wake County, North
Carolina. Geological data based on Clark et al. (2004); data for location map obtained from www.nationalatlas.gov.
Samples were collected at wellhead taps where available, or at outdoor taps after
ensuring that samples were not affected by chemical treatment systems. Dissolved oxygen (DO)
and pH were measured from a continuous flow at the bottom of a polyethylene bucket using YSI
DO200 and pH100 meters, respectively. Also from this continuous flow, a radon sample was
collected into a submerged glass vial, and an unfiltered alkalinity sample was collected into a
21
polyethylene bottle. Trace metal/cation samples were filtered in the field directly into new,
high‐purity acid‐washed polyethylene bottles containing high‐purity HNO3 using syringe‐tip
0.45 µm filters. Anion samples were similarly field‐filtered. In the lab, Ra isotopes were
concentrated from 52 L samples onto 10 g bundles of Mn oxide‐coated acrylic fibers (Moore and
Reid, 1973; Reid et al., 1979) at a flow rate of < 1 L min‐1. 19 of the 117 samples were
concentrated onto one column of Mn oxide fibers and the remaining 98 samples were
concentrated onto two sequential (A and B) columns (e.g. Luo et al., 2000). Radon‐222 samples
were shipped overnight to a commercial laboratory for determination by liquid scintillation
counting and are decay‐corrected to the time of collection.
2.2.2 Radium isotope analysis
Analysis of 224Ra was carried out as soon as possible after concentration onto Mn oxide
fibers by alpha counting using a delayed coincidence counter (Moore and Arnold, 1996). For
fibers with less than 20 total alpha counts per minute at time of analysis, 224Ra was quantified
using the 220Rn channel, corrected for chance coincidence events. For fibers with greater than 20
counts per minute, 224Ra was quantified by total counts, corrected for activity attributable to
223Ra (Garcia‐Solsona et al., 2008) and for 224Ra‐220Rn disequilibrium during short run times. 224Ra
was calibrated by an old, equilibrated 232Th solution that was transferred onto Mn oxide fibers.
224Ra was decay‐corrected to the time of collection and for 224Ra supported by small amounts of
228Th adsorbed onto the fibers after 3‐6 weeks. 226Ra analysis was performed by radon counting
after 20 days of incubation in a sealed, evacuated glass column on a Durridge RAD7 alpha
counting instrument. 226Ra standards are from a NIST solution transferred onto Mn oxide fibers.
228Ra analysis was performed on a Canberra broad‐energy Ge gamma spectrometer. In order to
22
produce a more efficient (e.g. Buesseler et al., 1995) and homogeneous (e.g. Herranz et al.,
2006) counting geometry, 228Ra analysis was performed on fibers of the A column that had been
compressed into a disk geometry 65 mm in diameter and ~5 mm thick in a metal canister. 228Ra
was quantified from a weighted average of counts from the 338 keV and 911 keV peaks of 228Ac
and corrected for the partitioning of Ra between the A and B columns using relative 224Ra
activities. The standard for gamma spectrometric analysis was U‐Th ore standard DL‐1a
(Canadian Certified Reference Materials Project), loaded to resemble the geometry of the
compressed fibers. For radium isotope analyses, propagation of error equations were applied to
background subtraction, delayed coincidence counter corrections, and addition of radium
activities from multiple columns using the methods of (Ivanovich and Murray, 1992) and Garcia‐
Solsona et al. (2008). Reported 2σ counting error (Table 6) approximates precision determined
by repeated analyses of samples and standards at comparable activities: (1) replicate analysis of
a 226Ra standard (equivalent to a 67 mBq L‐1 sample) indicates relative 2σ of 4% (n= 8); (2)
routine duplicate runs of 224Ra standards (equivalent to a 10 mBq L‐1 sample) indicate an average
difference of 9% between duplicates (n= 9 pairs); (3) replicate analysis during one day of a low‐
activity sample containing 3 mBq L‐1 of 224Ra indicates relative 2σ of 12% for 224Ra (n= 5); and (4)
replicate analysis of a sample equivalent to 38 mBq L‐1 of 228Ra indicates relative 2σ of 16% (n=
6).
Rolesville Granite samples collected from cores WC‐1 and WC‐3 at the Raleigh
Hydrogeologic Research Station (McSwain et al., in review), located within the study area (Figure
3), were crushed to powder, sealed into completely filled plastic canisters, incubated for three
weeks and analyzed for 226Ra and 228Ra using the 609 keV peak of 214Bi and the 911 keV peak of
228Ac, respectively. These were calibrated with the DL‐1a standard in the same geometry.
23
2.2.3 Major ions and trace elements
Ca, Mg, Na, Sr, Ba, Fe, Mn, and Si concentrations were determined by direct current
plasma spectrometry and K was determined by flame atomic absorption spectrometry, both
calibrated using solutions prepared from plasma‐grade single‐element standards. U
concentrations were determined by inductively‐coupled plasma mass spectrometry calibrated
with plasma‐grade U standard. As the wells used in this study contain metal surfaces, some
metal concentrations may be influenced by sampling materials. However, one blank each
evaluating bottles and filtration indicate minimal effects on major cations, Fe, Mn, and U. Major
anion (Cl‐, NO3‐, SO4
2‐) concentrations were determined by ion chromatography, and bicarbonate
concentrations were determined by titration to pH 4.5.
2.2.4 Data handling
Speciation and saturation index values were calculated using the PHREEQC geochemical
code (Parkhurst and Appelo, 1999; Post, 2006) modified with Ra speciation coefficients from
Langmuir and Riese (1985). Bedrock geology at each sampling site was assigned from 1:100,000‐
scale map data based on (Clark et al., 2004). Correlation coefficients reported in the text are
Spearman rank coefficients.
2.3 Results
2.3.1 Major ions, trace elements, and redox‐sensitive elements
The overall chemical composition of the investigated groundwater from the different
lithological settings is dominated by Na‐Ca‐Mg bicarbonate, except for a few samples with
anions dominated by a combination of bicarbonate and nitrate, chloride, or sulfate (Table 3).
One anomalous high‐sulfate water, well 64 in Group 2 (Table 4), was excluded from plots and
24
correlations due to its non‐representative chemistry. Average major element chemistry exhibits
few apparent differences between the three main rock types in the study area. Wells in the
Raleigh Gneiss exhibit lower average DO concentrations than the other rock types. Also, the
Rolesville Granite wells exhibit lower average Ca, Mg, K, and bicarbonate concentrations than
the other rock types (Table 3). However, concentrations of major elements and redox‐sensitive
elements vary significantly between wells. Most samples fall along a continuum between (1)
slightly acidic (pH 5.0‐6.0), oxic, low‐total dissolved solids (TDS) waters, and (2) near‐neutral,
oxic to anoxic, higher‐TDS waters (Figure 4). These trends illustrate the linked nature of the
major element and redox chemistry of the waters. Dissolved oxygen concentration is inversely
associated with Mn, Fe, and HCO3‐ concentrations (Figure 5), reflecting the sensitivity of these
solutes to redox processes. As barite (BaSO4) supersaturation is a potential sink for Ra,
saturation index values were calculated from major element chemistry and Ba concentrations;
however, waters in this study rarely approach barite saturation (median ‐1.9; range ‐3.4 to 0.2),
indicating a negligible role for barite.
2.3.2 Radionuclide activities in water
In contrast to major element chemistry, radionuclide activities exhibit large variation
between the Rolesville Granite and the other rock types in the study area. Median radon‐222
activity (249 Bq L‐1) in the Rolesville Granite exceeds the other rock units by about one order of
magnitude (Table 5), and median radium‐226 activity (18.6 mBq L‐1) is about a factor of 2 higher
in the Rolesville Granite (Table 5). Median uranium concentration (5 nM) is 1‐2 orders of
magnitude higher in the Rolesville Granite than in the other rock types (Table 5). Also, 222Rn
activity shows an inverse relationship with total dissolved ions (TDI) and direct relationship with
25
Figure 4: Relationships between pH, dissolved oxygen concentration, and total ions.
DO, whereas 226Ra is directly associated with TDI and inversely with DO (Figure 6). Relative to
EPA drinking water standards applied to public water systems, 77% of samples from the
Rolesville Granite exceed the proposed alternative maximum contaminant level (MCL) for 222Rn
of 148 Bq L‐1. 10% exceed the MCL for uranium of 126 nM (30 µg L‐1), and 6% exceed the MCL
for 226Ra+228Ra of 185 mBq L‐1. Radionuclide levels above the MCLs are uncommon in the other
rock types (Table 6). In the Rolesville Granite, U is correlated with pH (ρ= 0.52), HCO3‐ (ρ= 0.49),
26
Figure 5: Relationship between dissolved oxygen and other redox‐sensitive solutes.
27
Ca (ρ= 0.48), and sulfate (ρ= 0.47). However, no clear relationships were observed between U
and redox conditions or other radionuclides except for a relatively weak association with 226Ra
(ρ= 0.36). Because radionuclide activities and activity ratios are often lognormally distributed
(e.g. King et al., 1982), geometric mean and geometric standard deviation are presented in Table
5.
2.3.3 Radionuclide content of granite rock samples
Gamma spectrometric analysis of crushed Rolesville Granite bulk samples indicates 226Ra
activity from 20.8 to 53.5 Bq kg‐1 and 228Ra from 42.4 to 72.1 Bq kg‐1. This produces 228Ra/226Ra of
solids ranging from 0.99 to 2.68. These activities and activity ratios vary within each core and
between the two cores analyzed (Table 7).
2.4 Discussion
2.4.1 Major element and redox chemistry of potential Ra sinks
The high silica concentrations, dominance of bicarbonate among the dissolved anions,
and correlation between pH and TDI (ρ= 0.55; Figure 4) indicate that most of the solutes in the
investigated groundwater are derived from silicate weathering, which consumes protons,
contributes cations, and converts soil gas CO2 to bicarbonate (Drever, 1997). Bicarbonate may
also be derived from organic carbon oxidation or calcite dissolution, although the latter should
be insignificant in long‐exposed granites in areas where denudation rates are low, such as the
southeastern United States (White et al., 1999). The mass balance of dissolved oxygen (up to
~0.3 mM at atmospheric saturation, ranging to near zero in some samples) suggests that a large
proportion of bicarbonate (median 0.98 mM for all samples, 0.57 mM in the Rolesville Granite)
is derived from the open system of unsaturated zone CO2 rather than organic carbon oxidation
28
Figure 6: Relationship between dissolved oxygen, total ions, and activities of 222Rn and 226Ra in Rolesville Granite groundwater.
in the saturated zone. Consequently, the majority of cations are also derived from the shallow
weathering zone. The drinking water wells sampled in this study commonly contain open‐hole
intervals tens of meters long and produce water from an unknown configuration of fractures.
However, once groundwater passes through the weathered zone (soil and saprolite) into the
fracture network of relatively unweathered bedrock, where the open‐hole interval is located
29
(Daniel and Dahlen, 2002), the effect of mineral weathering on major element chemistry is
expected to be minimal due to the relatively short residence time in the fractures and the high
ratio of water to mineral surfaces. Therefore, major solute composition is determined during the
early stages of groundwater evolution.
Major species that are sensitive to redox conditions in groundwater include DO, nitrate,
Mn, Fe, sulfate, and bicarbonate. The large variations in DO, Fe, and Mn concentrations (Figure
5) indicate that redox processes play an important role in influencing water chemistry. In
fractured crystalline rocks, variations in redox condition may be due to abiotic reductants
including iron(II) in biotite (Gascoyne, 1997) or iron(II) and sulfide in pyrite (Gascoyne, 1997;
Tarits et al., 2006). Biotite is abundant in the Rolesville Granite (Kosecki and Fodor, 1997), and
pyrite has been observed in a pluton of similar age and setting in South Carolina (Speer et al.,
1981). Redox processes may also be driven by organic carbon oxidation (Banwart et al., 1996,
1999; Gascoyne, 2004). In the saturated zone, organic carbon could be oxidized to bicarbonate
by aerobic respiration or other terminal electron‐accepting processes including denitrification,
Mn oxide reduction, and Fe oxide reduction.
Mn concentrations over 2 µM and Fe concentrations over 10 µM are almost invariably
associated with DO concentrations below 50 µM (Figure 5). The elevated Mn and Fe in anoxic
fractures indicate conditions favorable for reductive dissolution of Mn and Fe oxides. No
evidence for sulfate reduction, such as a sulfide odor, was observed, suggesting that Mn and/or
Fe oxide reduction represent the most reducing biogeochemical conditions that prevail in a
significant number of wells. Finally, elevated bicarbonate concentrations, when associated with
decreasing DO concentrations (Figure 4), may indicate the effects of organic carbon oxidation,
30
although alternative mechanisms of producing bicarbonate and consuming DO, discussed
above, may be of equal or greater significance.
Our results indicate negative correlation of DO with total dissolved ion (TDI)
concentrations (ρ= ‐0.63; Figure 4), which suggests that DO is affected by some combination of
organic carbon oxidation and oxidation of minerals such as biotite or pyrite related to
weathering. Although determination of the rates of these processes is beyond the scope of this
paper, other studies indicate a negative relationship between DO and apparent residence time
in Piedmont groundwater and overall residence times of about 10‐50 a (e.g. Nelms and Harlow,
2003; D. Harned, USGS, Raleigh, NC, unpub. data, 2007). These relationships suggest the
potential importance of relatively slow processes such as mineral oxidation on water chemistry,
as the role of abiotic reductants in controlling groundwater redox conditions is expected to be
more important in older, slowly‐circulating groundwater systems (Gascoyne, 1997, 2004). In
comparison, consumption of dissolved oxygen by biogeochemical reactions in bedrock fractures
can be as rapid as several days (Puigdomenech et al., 2000).
2.4.2 Radium and radon isotopes
Overall activities of radium are higher in the Rolesville Granite than in the
metasedimentary and metavolcanic rocks and in the Raleigh Gneiss (Table 5). Also, Ra isotope
activities co‐occur, that is, high 226Ra from the 238U decay series tends to be accompanied by high
224Ra and 228Ra from the 232Th series (ρ= 0.78, 0.79, and 0.87 for the 226Ra‐228Ra, 224Ra‐226Ra, and
224Ra‐228Ra pairs, respectively, in the overall data set). In groundwater from the Rolesville
Granite, where Ra activities are higher and relative counting error lower, the strongest
correlations (ρ≥ 0.40) of 226Ra with other elements are with Mn > Ca > Ba > DO (ρ= 0.51, 0.41,
31
0.41, and ‐0.40, respectively). The strongest correlations of 228Ra are with Ba > K > Mn (ρ= 0.57,
0.49, and 0.40, respectively). Overall, Ra is positively correlated with Ba and to a lesser degree
with other alkaline earth metals. Ra is also correlated with redox‐sensitive solutes such as DO
and Mn in a consistent fashion: more reduced conditions are generally associated with higher Ra
(Figure 6). Although acidity (pH below 5) may significantly inhibit Ra adsorption (Nirdosh et al.,
1990; Moon et al., 2003), and pH variations are associated with variations in DO and TDI (Figure
4), pH does not exhibit strong correlations with Ra or Rn (ρ= 0.07, 0.23, and 0.00 between pH
and 222Rn, 226Ra, and 228Ra, respectively, in the Rolesville Granite). Mn oxides and clays exhibit
negative surface charge above pH 5 (Stumm and Morgan, 1996), so within the pH range in this
study, Ra adsorption is relatively unaffected by variations in pH. Also, because anion
concentrations in these waters are 1‐2 orders of magnitude below levels expected to
significantly complex Ra (Langmuir and Riese, 1985), aqueous complexation of Ra is negligible.
Speciation calculations indicate that virtually all aqueous Ra is present as Ra2+ (in Rolesville
Granite, median 99.7% as Ra2+, range 91.5%‐100%). The remaining Ra is complexed as RaSO40.
Results are essentially identical for Ba (median 98.1% as Ba2+, range 79.2‐100%). Thus, Ra occurs
primarily as Ra2+, which is strongly subject to adsorption at near‐neutral pH.
2.4.2.1 228Ra/226Ra activity ratio
The 228Ra/226Ra ratio in groundwater may be expected to resemble the 232Th/238U ratio
of the host rock (Asikainen, 1981a; Porcelli, 2008) assuming secular equilibrium within the rock.
Overall variations in 228Ra/226Ra occur with rock type; water samples from the Rolesville Granite
exhibit lower median 228Ra/226Ra than the Raleigh Gneiss and metasedimentary and
metavolcanic rocks (Table 5). Within the Rolesville Granite, the median 228Ra/226Ra of
32
groundwater is 0.68 (Table 5), but the median Th/U activity ratio calculated from Th and U
content of Rolesville Granite rock samples (median 17.2 mg kg‐1 and 4.4 mg kg‐1, respectively) is
1.2 (n= 42; data from Costain et al., 1977; Speer and Hoff, 1997; McSwain et al., in review).
These radionuclide levels are not unusually high for granites (e.g. Gascoyne, 1989; Speer and
Hoff, 1997). In addition, the median 228Ra/226Ra of water samples is lower than the median
228Ra/226Ra ratio measured directly from Rolesville Granite bulk rock samples (1.36; Table 7).
These ratios imply that some excess 226Ra relative to 228Ra is available to fracture surfaces,
compared to bulk rock ratios. Three possible mechanisms could reduce 228Ra/226Ra in water
relative to the bulk rock: (1) relative 226Ra enrichment could be derived from preferential recoil
processes within each decay series, enhancing the mobility of successive decay products. 226Ra is
the third alpha decay in its decay series, but 228Ra is the first decay step and 224Ra is the second
in the 232Th decay series. Thus, 226Ra would be more available due to the recoil effect placing it
closer to the mineral surface (Davidson and Dickson, 1986); (2) U, being more soluble than Th,
could be mobilized by weathering or the groundwater system to provide additional 226Ra
precursors to adsorption sites (e.g. Michel, 1984). Although some fractionation of 228Ra/226Ra is
observed relative to bulk rock samples, we note no predominance of 226Ra that would indicate
major net accumulations of U on fracture surfaces because wells with elevated overall Ra tend
to contain both 228Ra and 226Ra. Localized minerals containing both U and Th and/or adjacent
adsorption sites in fractures may be necessary, which implies that primary and/or secondary
minerals are significant Ra sources rather than net additions from long‐range U transport in
groundwater; and (3) dissolution of aquifer minerals would generate lower 228Ra/226Ra due to
the half‐life difference between 228Ra and 226Ra and the slow rate of dissolution (Hammond et
al., 1988). Dissolution as a direct source of Ra is ruled out given that short‐lived 224Ra is present
33
at similar activities as 228Ra (Figure 7; section 2.4.2.2), indicating rapid local interaction with the
aquifer rocks. Also, the strong retardation of Ra in fresh waters (retardation factor ~104 – 105;
(Krishnaswami et al., 1982; Dickson, 1990) would prevent 226Ra from accumulating along a
flowpath relative to 228Ra.
Figure 7: 228Ra and 224Ra activities in groundwater. Dashed line represents equilibrium value of 1. Counting error omitted where smaller than symbol size.
2.4.2.2 224Ra/228Ra activity ratio
Because all isotopes of Ra exhibit identical chemical behavior but are generated by
recoil at different rates, short‐lived to long‐lived isotope ratios may indicate the rates of Ra
contribution or removal. High ratios may indicate that the effective first‐order rate constant of
Ra immobilization exceeds the ingrowth rate (decay constant) of long‐lived Ra from alpha recoil,
thus indicating that Ra sinks are effective (Krishnaswami et al., 1982; Martin and Akber, 1999;
Moise et al., 2000). Low short‐lived to long‐lived isotope ratios may indicate a Ra source from
34
primary mineral dissolution in which short‐lived Ra decays faster than the dissolution rate
(Hammond et al., 1988). On average, 224Ra/228Ra activity ratios are close to the expected
equilibrium value of 1 (Figure 7; Table 5). This probably indicates that (1) Ra sources are broadly
distributed on fracture surfaces or are at least very close to the sampling point, rather than
being mobilized in an earlier phase of groundwater evolution that would allow significant 224Ra
decay; and (2) the rates of critical Ra‐mobilizing processes do not generate large changes in
224Ra/228Ra. Thus, we conclude that Ra is not released significantly by mineral dissolution, and
the rate constant of Ra adsorption allows long‐ and short‐lived isotopes to be generated by
recoil without large fluctuations in Ra isotope ratios.
2.4.2.3 Radon‐222
Radon‐222 activity in water may be interpreted as recoil flux from mineral surfaces
(Krishnaswami et al., 1982) or more complexly as a combination of recoil and diffusion of Rn
from microfractures resulting from a higher effective surface area for Rn relative to Ra (Rama
and Moore, 1984; Davidson and Dickson, 1986; Andrews et al., 1989). In this study, 222Rn is
strongly associated with rock type, and in particular, groundwater from the Rolesville Granite
exhibits approximately one order of magnitude higher median 222Rn (Table 5). In the granite,
222Rn activity is positively correlated with closed casing depth (ρ= 0.61), somewhat correlated
with DO (ρ= 0.42; Figure 6), and negatively correlated to some degree with most major cations
including Ba (ρ= ‐0.40). Overall, high Rn occurs mainly in oxic, low‐TDS waters (Figure 6). As Rn is
inert, these relationships indicate (1) that more oxic, low‐TDS conditions support more adsorbed
Ra on fracture surfaces per volume of water, consistent with a recoil‐dominated interpretation
35
(Krishnaswami et al., 1982); or (2) that oxic fractures contain a larger surface area per volume of
water, consistent with a diffusion‐based interpretation (Rama and Moore, 1984).
2.4.2.4 222Rn/226Ra activity ratio
222Rn/226Ra ratio distributions are generally consistent between rock types (Table 5),
suggesting that 222Rn/226Ra provides information on Ra adsorption in crystalline rock fractures
regardless of the Ra and Rn sources in each rock type. If the assumption that 222Rn activity in the
water is proportional to the recoil flux of Rn and Ra from the aquifer rocks is used, then the
Rn/Ra activity ratio is roughly equivalent to the distribution coefficient KD for radium (i.e.,
adsorption‐desorption equilibrium between aqueous and adsorbed Ra; Krishnaswami et al.,
1982; Dickson, 1990). The highest activities of Ra and Rn occur in distinct samples and thus
Rn/Ra occupies a large range. Given the approximately lognormal distributions of Rn and Ra,
median values of 222Rn/226Ra and 222Rn/224Ra on the order of 104 approximate KD (Figure 8).
Among redox‐sensitive solutes, dissolved oxygen concentration is positively correlated with
222Rn/226Ra (ρ= 0.52; Figure 9), and Mn is negatively correlated with 222Rn/226Ra (ρ= ‐0.56; Figure
9). Also, 222Rn/226Ra exhibits correlation with alkaline earth metals (ρ= ‐0.48 for Ba, ‐0.47 for Ca, ‐
0.45 for Sr, and ‐0.41 for TDI). As Ba in particular may compete with Ra for adsorption sites,
these trends could indicate that adsorption of Ra is related to cation concentrations, probably
derived from the weathering zone. Similar correlations are evident if 222Rn/224Ra is used to
represent apparent KD instead of 222Rn/226Ra. The significance of this approach is that the similar
half‐lives of 222Rn and 224Ra (3.8 d and 3.7 d, respectively) negate dissolution contributions of
long‐lived 226Ra (Krishnaswami et al., 1982). The median apparent KD estimated from 222Rn/226Ra
decreases by about one order of magnitude between oxygen‐saturated and anoxic conditions
36
and along the range of Mn, TDI, and Ba concentrations (~20,000 to ~2000; Figure 10). In
summary, trends in 222Rn/226Ra suggest that adsorption of recoil‐produced Ra varies as a result
of local chemical conditions.
Figure 8: Relationship between 222Rn and 226Ra (left), 222Rn and 224Ra (right) and histograms of isotope activities in Rolesville Granite groundwater.
2.4.3 Overview of radionuclide sources and removal mechanisms
At small scales, much of the U in granite is distributed in secondary minerals in
microfractures and adsorbed to Fe oxides, Mn oxides, and clays (Speer et al., 1981; Guthrie and
Kleeman, 1986; Wathen, 1987). In addition, significant disequilibrium of solid‐phase U‐ and Th‐
series radionuclides at the scale of meters indicates long‐term mobility of radionuclides,
probably by groundwater (e.g. Michel, 1984). Long‐lived U may be mobile in groundwater
systems until reaching reducing zones, at which it is strongly adsorbed to Fe oxides (Ames et al.,
37
1983c; Waite et al., 1994). However, the strong positive correlation of 226Ra and 228Ra obtained
in this study indicates that both 226Ra and 228Ra are present within recoil range of fracture
surfaces, indicating the co‐occurrence of Th and U in primary and/or secondary minerals near
fracture surfaces.
Figure 9: Relationship between dissolved oxygen and Mn (top), Ba and total ions (bottom), and the 222Rn/226Ra activity ratio.
38
Figure 10: Apparent KD across dissolved oxygen and Mn (top) and Ba and total ions (bottom). Steps represent median 222Rn/226Ra in each of 10 concentration intervals containing
approximately equal numbers of samples.
The relatively high minimum values of apparent KD in fractured crystalline rocks (~103)
indicate that Ra is rapidly adsorbed in fresh waters in bedrock fractures to balance input from
alpha recoil, and that Ra does not exhibit the lower apparent KD commonly observed in
sandstone aquifers (~101‐103; King et al., 1982; Lively et al., 1992; Reynolds et al., 2003) or saline
39
aquifers (~100‐102; Sturchio et al., 2001). Net adsorption of Ra is less effective under the low‐DO,
higher‐TDS conditions observed in this study because (1) anoxic fractures contain lower surface
area of redox‐sensitive Ra adsorption sites, such as Mn oxides; and/or (2) higher concentrations
of alkaline earth metals, such as Ba, compete with recoil‐generated Ra for adsorption sites. The
down‐flow transition from the saprolite to water‐producing bedrock fractures involves
increasing water‐to‐mineral surface ratio, decreasing clay content, and decreasing mineral‐
water contact time. Thus, cations in waters passing from highly to lightly weathered zones may
experience increased competition for adsorption sites. Our results indicate that in the saprolite‐
fractured crystalline rock system, major ion concentrations and redox state (dissolved oxygen
consumption, Mn oxide reduction, and Fe oxide reduction) are linked. Fractures receiving water
from more weathered recharge sources contain higher concentrations of alkaline earth metals
and bicarbonate. As these waters also tend to exhibit lower DO concentrations, they have also
been more involved in organic carbon oxidation and/or biotite and pyrite oxidation.
Consequently, the distinction between the individual contributions of cation concentrations
versus redox conditions on Ra mobilization is not straightforward, and we cannot differentiate
the redox and cation exchange effects on apparent KD.
2.5 Conclusions
Although the main influence on radionuclides in groundwater is the U and Th content on
fracture surfaces, linkage between weathering‐derived ion concentrations, redox conditions,
and radionuclide mobility in bedrock fractures occurs across a range of rock types and
radionuclide activities. The primary source of Ra and Rn in groundwater is recoil from parent
radionuclides on fracture surfaces, which is balanced by adsorption of Ra. While groundwater
40
chemistry is mostly acquired during recharge through the soil and saprolite, radionuclides are
generated by recoil in the saturated bedrock fracture network. Thus, the efficiency of adsorption
of Ra is influenced by the redox‐sensitive stability of adsorption sites and/or competition with
other divalent cations. In low‐DO samples, elevated Mn and Fe concentrations suggest that their
oxides are subject to reductive dissolution and thus less efficient at removing Ra. This
interpretation is consistent with observed trends in Ra and Rn activities, as well as variable
222Rn/226Ra ratios correlated with DO, Mn, Ba, and TDI concentrations. Overall, the
approximately one order of magnitude variation in the apparent Ra distribution coefficient
across the range of dissolved oxygen and alkaline earth metal concentrations indicates that in
fresh groundwaters from fractured crystalline rocks, water chemistry is an important influence
of on Ra adsorption, although secondary in importance to the radionuclide content of the solids.
41
Table 3: Summary of major and trace element concentrations and pH. ND indicates not detected.
Ca Mg Na K Fe Mn Sr Ba Cl‐ NO3‐ SO4
2‐ HCO3‐ SiO2 pH DO
All data (n=117) (mM) (mM) (mM) (mM) (µM) (µM) (µM) (µM) (mM) (mM) (mM) (mM) (mM) (µM) Mean 0.30 0.13 0.46 0.05 4.27 0.85 0.81 0.23 0.16 0.097 0.050 1.10 0.57 6.4 129 Median 0.20 0.11 0.42 0.05 0.25 0.15 0.68 0.14 0.09 0.036 0.010 0.98 0.56 6.3 147 2σ 0.55 0.21 0.39 0.05 34.92 3.75 1.13 0.57 0.51 0.280 0.265 1.47 0.32 1.4 160 Minimum 0.02 0.01 0.11 0.02 ND ND 0.11 0.01 0.04 ND ND 0.13 0.19 4.7 6 Maximum 1.38 0.55 1.34 0.17 167.50 12.92 2.97 1.65 2.41 0.643 1.248 3.75 1.08 8.6 269 Rolesville Granite (n=52) Mean 0.27 0.09 0.47 0.05 1.86 0.44 0.88 0.28 0.16 0.125 0.044 0.90 0.59 6.1 155 Median 0.14 0.05 0.43 0.04 0.23 0.11 0.68 0.15 0.11 0.094 0.007 0.57 0.58 6.0 173 2σ 0.64 0.20 0.35 0.03 0.90 1.69 1.25 0.74 0.23 0.232 0.141 1.68 0.32 1.3 159 Minimum 0.03 0.01 0.11 0.02 ND ND 0.23 0.01 0.04 ND ND 0.13 0.22 4.7 13 Maximum 1.38 0.55 1.25 0.12 58.32 5.12 2.97 1.65 0.61 0.460 0.332 3.75 1.08 7.5 269 Raleigh Gneiss (n=24) Mean 0.25 0.14 0.38 0.06 6.85 1.71 0.69 0.17 0.11 0.079 0.085 1.00 0.56 6.4 76 Median 0.19 0.12 0.32 0.07 0.63 0.44 0.46 0.13 0.08 0.011 0.021 1.05 0.54 6.4 38 2σ 0.40 0.19 0.42 0.05 1.41 5.90 1.12 0.24 0.20 0.321 0.511 1.02 0.41 1.3 144 Minimum 0.08 0.02 0.21 0.03 ND 0.05 0.23 0.06 0.05 ND 0.003 0.31 0.19 5.5 6 Maximum 0.98 0.38 1.07 0.12 51.71 12.92 2.74 0.51 0.53 0.643 1.248 2.31 1.07 8.2 188 Metasedimentary & metavolcanic rocks (n=34) Mean 0.36 0.18 0.50 0.06 6.99 1.01 0.78 0.18 0.22 0.075 0.038 1.41 0.55 6.6 118 Median 0.28 0.17 0.45 0.06 0.23 0.12 0.68 0.12 0.08 0.004 0.012 1.40 0.53 6.5 131 2σ 0.48 0.19 0.43 0.06 3.22 4.21 0.96 0.31 0.87 0.331 0.185 1.21 0.28 1.4 143 Minimum 0.02 0.02 0.23 0.03 ND ND 0.11 0.02 0.05 ND ND 0.25 0.30 5.1 9 Maximum 0.91 0.43 1.34 0.17 167.50 9.23 1.83 0.55 2.41 0.638 0.502 2.92 0.83 8.6 263
Table 4: Major and trace element concentrations and pH.
Group Well Sampling Rock Type Well Casing pH DO Mg Ca Sr Ba Na K Cl‐ NO3‐ SO4
2‐ HCO3‐ SiO2 Mn Fe
ID date depth (m) depth (m) (µM) (mM) (mM) (µM) (µM) (mM) (mM) (mM) (mM) (mM) (mM) (mM) (µM) (µM)
1 A 20 19 Jan 2007 Metasedimentary/
Metavolcanic124 18 8.57 127 0.27 0.24 0.58 0.32 0.36 0.09 0.06 0.008 0.006 1.37 0.56 0.11 0.41
1 A 22 21 Jan 2007 Metasedimentary/
Metavolcanic125 20 6.32 89 0.20 0.50 0.39 0.07 0.57 0.07 0.08 0.066 0.039 1.90 0.41 0.04 0.25
1 A 23 21 Jan 2007 Metasedimentary/
Metavolcanic179 18 7.50 87 0.43 0.91 1.56 0.35 0.53 0.17 0.42 0.109 0.028 2.92 0.45 0.05 0.18
1 a 24 21 Jan 2007 Metasedimentary/
Metavolcanic57 21 7.20 222 0.08 0.13 0.37 0.08 0.23 0.04 0.08 0.110 0.005 0.53 0.54 0.15 0.26
1 a 25 21 Jan 2007 Metadiorite/
Metagranodiorite107 14 6.79 145 0.32 0.80 1.55 0.93 0.70 0.09 0.38 0.028 0.126 2.28 0.70 0.32 0.34
1 a 26 26 Jan 2007 Ultramafic 124 19 7.12 171 0.14 0.17 0.38 0.13 0.27 0.09 0.07 0.111 0.007 0.89 0.55 0.16 0.491
a 27 26 Jan 2007 Metadiorite/
Metagranodiorite 6.54 240 0.11 0.16 0.50 0.05 0.29 0.03 0.10 0.162 0.004 1.16 0.37 0.15 0.23
1 a 28 26 Jan 2007 Metasedimentary/
Metavolcanic112 8 7.50 149 0.14 0.22 0.42 0.05 0.34 0.07 0.06 0.016 0.016 1.23 0.47 0.15 0.18
1 a 29 26 Jan 2007 Metasedimentary/
Metavolcanic92 6.04 151 0.22 0.25 0.66 0.23 0.46 0.06 0.10 0.010 0.006 1.33 0.66 0.07 0.08
1 a 30 26 Jan 2007 Metasedimentary/
Metavolcanic81 15 6.20 180 0.23 0.72 1.05 0.12 0.47 0.06 0.39 0.495 0.004 2.37 0.51 0.12 0.21
1 a 31 28 Jan 2007 Metadiorite/
Metagranodiorite76 13 6.33 47 0.21 0.64 1.20 0.48 0.46 0.09 0.11 0.022 0.039 1.74 0.59 1.04 0.68
1 a 32 28 Jan 2007 Metasedimentary/
Metavolcanic 6.21 134 0.33 0.56 1.87 0.56 0.67 0.09 1.01 0.602 0.007 1.02 0.50 0.03 0.31
1 a 34 28 Jan 2007 Metasedimentary/
Metavolcanic75 31 6.51 160 0.13 0.19 0.63 0.24 0.42 0.06 0.07 0.005 0.005 1.02 0.72 0.20 ND
1 a 35 28 Jan 2007 Metadiorite/
Metagranodiorite50 18 6.15 96 0.21 0.33 1.10 0.62 0.62 0.05 0.18 0.028 0.052 1.95 0.79 0.07 0.72
1 a 36 02 Feb 2007 Metasedimentary/
Metavolcanic 6.62 263 0.11 0.12 0.50 0.03 0.31 0.04 0.08 0.007 0.004 0.76 0.43 ND ND
1 a 37 02 Feb 2007 Metasedimentary/
Metavolcanic 6.88 133 0.14 0.63 0.79 0.05 0.58 0.04 0.10 0.066 0.243 1.57 0.37 0.04 0.15
1 a 38 02 Feb 2007 Metasedimentary/
Metavolcanic81 19 7.45 32 0.20 0.91 1.41 0.13 0.93 0.07 0.09 0.002 0.501 2.07 0.42 0.18 0.46
42
43
Group Well Sampling Rock Type Well Casing pH DO Mg Ca Sr Ba Na K Cl‐ NO3‐ SO4
2‐ HCO3‐ SiO2 Mn Fe
ID date depth (m) depth (m) (µM) (mM) (mM) (µM) (µM) (mM) (mM) (mM) (mM) (mM) (mM) (mM) (µM) (µM) 1
a 39 02 Feb 2007 Metasedimentary/
Metavolcanic55 20 6.21 46 0.30 0.78 1.20 0.35 0.42 0.08 0.52 0.001 0.095 2.08 0.55 9.24 16.95
1 a 40 02 Feb 2007 Metasedimentary/
Metavolcanic69 12 6.13 218 0.26 0.50 1.64 0.54 0.67 0.07 0.44 0.638 0.015 1.25 0.63 0.06 0.22
1 a 48 11 Feb 2007 Metasedimentary/
Metavolcanic93 11 5.78 0.10 0.22 0.58 0.06 0.32 0.03 0.09 0.025 0.005 1.42 0.41 0.07 0.35
1 a 49 11 Feb 2007 Metasedimentary/
Metavolcanic185 15 6.31 0.31 0.35 1.20 0.18 0.48 0.06 0.14 0.150 0.006 1.68 0.56 0.03 0.14
1 b 19 19 Jan 2007 Metasedimentary/
Metavolcanic104 12 6.87 19 0.12 0.34 1.23 0.04 0.86 0.05 0.06 ND 0.010 1.75 0.29 0.17 0.16
1 b 41 09 Feb 2007 Metadiorite/
Metagranodiorite 7.16 68 0.17 0.30 0.74 0.11 0.40 0.06 0.06 0.001 0.049 1.35 0.56 0.78 1.36
1 b 42 09 Feb 2007 Raleigh gneiss 46 29 7.08 13 0.34 0.56 0.94 0.06 0.54 0.06 0.09 0.001 0.039 2.31 0.53 1.16 8.37
1 b 43 09 Feb 2007 Metasedimentary/
Metavolcanic104 6.51 0.15 0.28 0.50 0.22 0.44 0.04 0.07 0.001 0.009 1.30 0.81 0.02 0.03
1 b 44 09 Feb 2007 Metasedimentary/
Metavolcanic70 38 6.32 12 0.20 0.29 0.33 0.10 0.39 0.06 0.06 ND 0.030 1.93 0.62 3.91 15.45
1 b 45 09 Feb 2007 Metasedimentary/
Metavolcanic92 28 5.99 184 0.02 0.06 0.16 0.16 0.35 0.04 0.05 0.001 0.005 0.58 0.65 0.00 0.32
1 b 50 16 Feb 2007 Metasedimentary/
Metavolcanic 5.60 127 0.30 0.49 1.76 0.44 1.34 0.07 2.41 0.080 0.025 1.42 0.79 0.08 0.03
1 b 51 16 Feb 2007 Metasedimentary/
Metavolcanic99 47 8.28 10 0.17 0.32 1.02 0.02 0.71 0.06 0.07 0.001 0.016 1.67 0.32 0.06 ND
1 b 52 16 Feb 2007 Raleigh gneiss 113 37 7.82 38 0.13 0.40 1.79 0.09 1.06 0.07 0.07 0.001 0.032 1.87 0.48 0.55 0.59
1 b 53 16 Feb 2007 Metasedimentary/
Metavolcanic75 32 6.29 27 0.12 0.10 0.35 0.04 0.44 0.03 0.06 0.001 0.022 1.10 0.69 3.08 10.99
1 b 54 16 Feb 2007 Metadiorite/
Metagranodiorite 6.30 168 0.05 0.08 0.28 0.09 0.21 0.02 0.06 0.008 0.003 0.48 0.64 0.06 0.00
1 b 55 23 Feb 2007 Metasedimentary/
Metavolcanic 6.70 130 0.17 0.34 0.65 0.11 0.42 0.09 0.06 0.002 0.005 1.45 0.56 0.19 0.25
1 b 56 23 Feb 2007 Metasedimentary/
Metavolcanic86 20 6.20 164 0.04 0.09 0.29 0.16 0.31 0.03 0.06 0.001 0.005 0.64 0.60 0.23 ND
1 b 57 23 Feb 2007 Metasedimentary/
Metavolcanic75 24 7.21 36 0.18 0.66 0.69 0.27 0.46 0.04 0.07 0.001 0.051 2.19 0.81 3.66 2.00
1 b 58 23 Feb 2007 Raleigh gneiss 39 6.29 15 0.23 0.39 0.69 0.17 0.32 0.07 0.12 0.001 0.079 1.58 0.52 5.33 22.97
44
Group Well Sampling Rock Type Well Casing pH DO Mg Ca Sr Ba Na K Cl‐ NO3‐ SO4
2‐ HCO3‐ SiO2 Mn Fe
ID date depth (m) depth (m) (µM) (mM) (mM) (µM) (µM) (mM) (mM) (mM) (mM) (mM) (mM) (mM) (µM) (µM) 1 b 59 23 Feb 2007
Metasedimentary/ Metavolcanic
5.12 208 0.02 0.03 0.17 0.41 0.48 0.05 0.18 0.124 ND 0.24 0.47 0.11 0.20
1 b 60 25 Feb 2007 Metasedimentary/
Metavolcanic80 25 6.65 166 0.15 0.23 0.48 0.03 0.47 0.03 0.05 ND 0.011 1.09 0.83 0.65 0.11
1 b 61 25 Feb 2007 Metasedimentary/
Metavolcanic 6.95 83 0.24 0.51 1.30 0.03 0.62 0.09 0.10 ND 0.012 2.06 0.47 0.08 0.40
1 b 62 25 Feb 2007 Metasedimentary/
Metavolcanic 7.26 81 0.24 0.43 0.78 0.19 0.42 0.07 0.09 ND 0.016 1.90 0.43 1.37 1.10
1 b 63 25 Feb 2007 Metasedimentary/
Metavolcanic76 32 6.10 188 0.06 0.04 0.18 0.03 0.27 0.03 0.07 0.028 0.017 0.26 0.52 0.22 0.31
1 b 74 04 Mar 2007 Metasedimentary/
Metavolcanic69 33 6.91 153 0.23 0.25 0.56 0.04 0.50 0.07 0.06 ND 0.012 1.44 0.48 0.05 0.21
1 b 75 04 Mar 2007 Metasedimentary/
Metavolcanic70 57 6.81 33 0.23 0.22 0.46 0.02 0.45 0.06 0.06 ND 0.032 1.38 0.70 3.18 18.37
2 11 14 Jan 2007 Raleigh gneiss 24 20 6.45 11 0.32 0.18 0.41 0.25 0.22 0.10 0.08 0.002 0.046 1.23 1.07 12.93 51.712 64 02 Mar 2007 Raleigh gneiss 73 10 6.86 30 0.17 9.26 28.37 0.13 7.34 0.07 0.92 ND 13.261 0.31 0.24 1.42 1.62
2 65 02 Mar 2007 Metasedimentary/
Metavolcanic75 44 6.32 25 0.09 0.18 0.95 0.35 0.28 0.04 0.06 ND 0.024 0.92 0.47 6.68 167.52
2 66 02 Mar 2007 Raleigh gneiss 61 16 7.01 14 0.15 0.40 0.94 0.12 0.40 0.07 0.08 ND 0.106 1.37 0.60 4.48 7.882 67 02 Mar 2007 Raleigh gneiss 52 15 6.49 43 0.14 0.32 0.78 0.08 0.37 0.07 0.07 0.009 0.075 1.14 0.68 0.23 0.092 68 09 Mar 2007 Raleigh gneiss 28 6.39 172 0.11 0.20 0.79 0.13 0.30 0.04 0.08 0.033 0.003 0.88 0.57 0.06 0.692 69 09 Mar 2007 Raleigh gneiss 6.30 13 0.14 0.13 0.38 0.13 0.33 0.07 0.06 0.001 0.076 1.06 0.90 0.86 1.742 70 09 Mar 2007 Raleigh gneiss 122 31 5.94 172 0.08 0.12 0.36 0.07 0.31 0.06 0.08 0.086 0.006 1.44 0.64 0.05 ND2 71 09 Mar 2007 Raleigh gneiss 37 20 7.05 7 0.14 0.19 0.46 0.30 0.33 0.07 0.05 ND 0.038 1.05 0.65 3.20 20.972 72 04 Mar 2007 Raleigh gneiss 6.04 154 0.04 0.11 0.30 0.14 0.21 0.03 0.09 0.077 0.038 0.45 0.30 0.43 ND2 73 04 Mar 2007 Raleigh gneiss 69 28 5.54 186 0.04 0.10 0.36 0.11 0.29 0.04 0.07 0.030 0.003 0.53 0.52 0.32 0.632 76 04 Mar 2007 Raleigh gneiss 64 29 6.36 9 0.19 0.25 0.59 0.46 0.28 0.05 0.06 ND 0.017 1.18 0.62 4.93 25.992 77 18 Mar 2007 Raleigh gneiss 6.44 22 0.14 0.17 0.46 0.16 0.23 0.09 0.06 ND 0.096 0.81 0.89 1.55 12.172 78 18 Mar 2007 Raleigh gneiss 93 21 5.97 175 0.04 0.08 0.27 0.10 0.22 0.04 0.07 0.067 0.004 0.42 0.48 0.09 0.152 79 18 Mar 2007 Raleigh gneiss 75 31 6.39 119 0.12 0.21 0.77 0.18 0.29 0.07 0.07 0.011 0.006 1.06 0.63 0.28 0.372 80 18 Mar 2007 Raleigh gneiss 5.94 173 0.07 0.23 0.96 0.20 0.41 0.07 0.20 0.643 0.006 0.31 0.49 0.15 1.042 81 18 Mar 2007 Raleigh gneiss 165 14 6.11 95 0.10 0.16 0.48 0.06 0.37 0.05 0.12 0.085 0.019 0.76 0.55 0.22 0.16
45
Group Well Sampling Rock Type Well Casing pH DO Mg Ca Sr Ba Na K Cl‐ NO3‐ SO4
2‐ HCO3‐ SiO2 Mn Fe
ID date depth (m) depth (m) (µM) (mM) (mM) (µM) (µM) (mM) (mM) (mM) (mM) (mM) (mM) (mM) (µM) (µM) 2 82 25 Mar 2007 Raleigh gneiss 61 43 6.16 0.03 0.08 0.39 0.07 0.32 0.03 0.07 0.014 0.003 0.65 0.60 0.06 ND2 83 25 Mar 2007 Raleigh gneiss 5.56 0.06 0.13 0.34 0.28 0.28 0.03 0.15 0.368 0.003 0.37 0.23 0.16 0.222 84 25 Mar 2007 Raleigh gneiss 70 24 6.46 0.11 0.15 0.49 0.16 0.24 0.05 0.07 0.022 0.003 0.82 0.47 0.05 0.282 85 25 Mar 2007 Raleigh gneiss 73 42 8.24 13 0.38 0.97 2.78 0.11 0.91 0.12 0.16 0.005 1.248 1.29 0.33 1.79 1.452 86 25 Mar 2007 Raleigh gneiss 5.87 0.08 0.23 0.41 0.51 0.45 0.12 0.53 0.364 0.021 0.42 0.19 0.47 0.093 a 7 12 Jan 2007 Granite 55 20 6.31 176 0.02 0.07 0.43 0.20 0.35 0.05 0.05 0.001 0.003 0.64 0.83 0.03 0.173 a 8 14 Jan 2007 Granite 46 7 6.80 188 0.12 0.33 1.03 0.07 0.53 0.04 0.10 0.198 0.046 1.03 0.60 0.11 ND3 a 9 14 Jan 2007 Granite 18 13 5.74 213 0.05 0.18 0.73 0.13 0.41 0.05 0.11 0.234 0.001 0.44 0.53 0.15 0.073 a 10 14 Jan 2007 Granite 38 16 6.01 215 0.06 0.18 0.93 0.14 0.33 0.06 0.08 0.061 0.003 0.74 0.53 2.45 3.613 a 12 14 Jan 2007 Granite 6.33 178 0.20 0.24 0.63 0.30 0.43 0.03 0.16 0.198 0.016 0.91 0.75 0.07 0.073 a 13 15 Jan 2007 Granite 49 29 6.50 13 0.05 0.08 0.49 0.31 0.45 0.06 0.04 ND 0.005 1.42 1.08 1.17 58.323 a 14 15 Jan 2007 Granite 61 18 5.77 171 0.03 0.10 0.79 0.30 0.42 0.04 0.17 0.107 0.004 0.52 0.72 0.14 0.343 a 16 15 Jan 2007 Granite 92 18 5.68 268 0.01 0.03 0.31 0.11 0.19 0.04 0.05 0.011 0.002 0.29 0.52 0.00 0.343 a 17 15 Jan 2007 Granite 75 20 6.02 260 0.02 0.08 0.44 0.20 0.25 0.05 0.08 0.117 0.004 0.41 0.54 0.05 0.803 a 18 19 Jan 2007 Granite 147 21 6.87 118 0.04 0.20 0.95 0.25 0.53 0.05 0.08 0.062 0.085 0.78 0.46 0.35 ND3 a 21 19 Jan 2007 Granite 43 24 6.00 136 0.02 0.07 0.43 0.20 0.33 0.05 0.05 0.004 0.003 0.61 0.68 0.23 1.853 a 33 28 Jan 2007 Granite 55 28 6.62 248 0.02 0.06 0.35 0.12 0.33 0.04 0.05 0.004 0.004 0.53 0.67 0.10 0.273 a 47 11 Feb 2007 Granite 52 6.61 15 0.22 0.73 2.31 1.09 0.59 0.04 0.15 0.002 0.150 2.17 0.70 1.68 6.813 a 89 09 May 2007 Granite 5.28 200 0.03 0.08 0.52 0.31 0.43 0.06 0.36 0.192 0.002 0.21 0.32 0.09 0.103 a 90 09 May 2007 Granite 69 25 6.04 33 0.06 0.13 0.43 0.12 0.30 0.05 0.10 0.074 0.141 0.60 0.61 0.84 1.143 a 92 11 May 2007 Granite 6.48 259 0.04 0.09 0.41 0.02 0.36 0.04 0.06 0.011 0.006 0.58 0.79 0.08 ND3 a 93 11 May 2007 Granite 58 31 6.17 214 0.02 0.07 0.42 0.17 0.27 0.06 0.07 0.039 0.004 0.46 0.68 0.11 0.103 a 95 20 May 2007 Granite 5.25 228 0.02 0.03 0.20 0.28 0.11 0.05 0.07 0.049 ND 0.25 0.22 0.06 ND3 a 96 20 May 2007 Granite 5.81 135 0.06 0.15 0.99 0.15 0.57 0.04 0.14 0.237 0.002 0.65 0.81 0.10 0.103 a 97 20 May 2007 Granite 45 24 5.93 255 0.02 0.06 0.28 0.18 0.29 0.05 0.12 0.184 0.003 0.30 0.50 0.02 0.033 a 99 21 Oct 2007 Granite 92 5.87 236 0.04 0.10 0.43 0.06 0.39 0.04 0.05 0.010 0.005 0.56 0.80 0.06 0.133 b 1 14 Dec 2006 Granite 343 12 6.30 17 0.07 0.24 0.85 0.13 0.69 0.03 0.28 0.156 0.021 2.09 0.47 0.52 1.833 b 2 14 Dec 2006 Granite 119 19 6.92 44 0.09 0.30 0.91 0.36 0.45 0.03 0.10 0.086 0.011 1.18 0.55 0.26 0.073 b 3 14 Dec 2006 Granite 275 18 5.96 68 0.11 0.51 0.81 0.06 0.55 0.04 0.28 0.304 0.086 1.10 0.43 0.30 0.073 b 4 14 Dec 2006 Granite 7.44 26 0.34 0.84 3.01 1.64 1.25 0.04 0.10 0.002 0.087 3.24 0.54 1.19 1.58
46
Group Well Sampling Rock Type Well Casing pH DO Mg Ca Sr Ba Na K Cl‐ NO3‐ SO4
2‐ HCO3‐ SiO2 Mn Fe
ID date depth (m) depth (m) (µM) (mM) (mM) (µM) (µM) (mM) (mM) (mM) (mM) (mM) (mM) (mM) (µM) (µM) 3 b 5 12 Jan 2007 Granite 127 10 5.23 241 0.02 0.08 0.41 0.12 0.40 0.03 0.10 0.187 0.003 0.35 0.57 0.06 0.343 b 6 12 Jan 2007 Granite 104 9 7.10 12 0.55 1.38 2.47 1.55 0.60 0.05 0.14 0.007 0.189 3.76 0.45 5.12 2.763 b 15 15 Jan 2007 Granite 55 13 6.43 139 0.08 0.26 0.92 0.04 0.55 0.02 0.09 0.118 0.076 1.16 0.70 0.12 0.253 b 46 11 Feb 2007 Granite 5.02 35 0.05 0.14 0.87 0.73 0.62 0.12 0.44 0.346 0.029 0.28 0.47 0.08 0.623 b 87 09 May 2007 Granite 5.37 147 0.07 0.19 1.44 0.35 0.78 0.07 0.61 0.318 0.041 0.38 0.58 0.09 ND3 b 88 09 May 2007 Granite 21 6 5.46 235 0.02 0.07 0.38 0.03 0.47 0.03 0.10 0.346 0.006 0.15 0.49 0.14 0.263 b 91 11 May 2007 Granite 6.10 177 0.03 0.26 0.53 0.14 0.42 0.03 0.10 0.051 0.048 0.79 0.38 0.41 3.853 b 94 13 May 2007 Granite 5.46 103 0.02 0.18 0.55 1.38 0.34 0.08 0.24 0.206 0.008 0.41 0.31 0.41 ND3 b 98 10 Sep 2007 Granite 6.59 158 0.19 0.52 1.07 0.40 0.51 0.03 0.15 0.062 0.095 1.51 0.57 0.05 0.183 c 100 21 Oct 2007 Granite 31 5.68 253 0.08 0.13 0.59 0.03 0.57 0.04 0.17 0.354 0.001 0.29 0.67 0.05 0.183 c 101 21 Oct 2007 Granite 47 5.75 133 0.06 0.14 0.56 0.20 0.30 0.04 0.08 0.142 0.108 0.24 0.42 1.13 3.953 c 102 21 Oct 2007 Granite 124 16 6.75 69 0.19 1.35 2.15 0.14 0.57 0.04 0.18 0.051 0.130 3.19 0.51 0.11 0.123 c 103 26 Oct 2007 Granite 206 0.09 0.13 0.50 0.13 0.33 0.04 0.11 0.104 0.009 0.48 0.58 0.05 0.233 c 104 26 Oct 2007 Granite 69 17 7.10 162 0.11 0.98 1.63 0.15 0.53 0.04 0.17 0.069 0.087 2.03 0.54 1.37 0.263 c 105 26 Oct 2007 Granite 73 19 7.25 36 0.24 0.96 2.27 0.01 0.67 0.04 0.21 0.030 0.332 2.16 0.50 0.04 0.083 c 106 11 Nov 2007 Granite 58 16 5.76 175 0.05 0.17 1.00 0.29 0.57 0.06 0.16 0.460 0.001 0.38 0.73 0.12 0.053 c 107 11 Nov 2007 Granite 44 18 7.47 14 0.38 0.70 0.99 0.11 0.53 0.04 0.14 0.014 0.117 2.35 0.64 1.82 0.243 c 108 11 Nov 2007 Granite 130 6.03 175 0.12 0.50 2.11 0.17 0.67 0.05 0.36 0.112 0.262 0.94 0.52 0.11 0.053 c 109 11 Nov 2007 Granite 40 19 6.04 236 0.03 0.08 0.41 0.09 0.40 0.04 0.13 0.153 0.005 0.29 0.73 0.09 1.153 c 110 11 Nov 2007 Granite 61 14 4.73 163 0.05 0.09 0.84 0.86 0.53 0.07 0.42 0.196 0.004 0.13 0.31 0.28 0.213 c 111 18 Nov 2007 Granite 45 16 5.10 200 0.04 0.09 0.44 0.10 0.38 0.03 0.12 0.043 0.010 0.43 0.65 ND 0.063 c 112 18 Nov 2007 Granite 87 34 5.68 172 0.10 0.20 1.29 0.24 0.61 0.06 0.31 0.340 0.005 0.58 0.75 0.09 0.153 c 113 18 Nov 2007 Granite 5.90 156 0.07 0.14 0.75 0.14 0.42 0.05 0.10 0.021 ND 0.76 0.66 0.35 0.623 c 114 18 Nov 2007 Granite 5.91 213 0.06 0.13 0.63 0.13 0.43 0.04 0.17 0.232 ND 0.38 0.68 0.09 0.593 c 115 18 Nov 2007 Granite 86 6.54 213 0.05 0.13 0.35 0.05 0.36 0.02 0.10 0.047 0.009 0.57 0.81 0.10 0.333 c 116 07 Dec 2007 Granite 88 0.05 0.14 0.65 0.09 0.44 0.04 0.09 0.063 0.010 0.68 0.78 0.40 0.333 c 117 07 Dec 2007 Granite 83 216 0.03 0.09 0.55 0.06 0.35 0.02 0.17 0.102 0.001 0.33 0.50 0.04 1.80
ND indicates not detected; these were treated as zero concentrations in data analysis.
47
Table 5: Summary of radionuclide activities and activity ratios.
All data 222Rn (Bq L‐1) 224Ra (mBq L‐1) 226Ra (mBq L‐1) 228Ra (mBq L‐1) U (nM) 222Rn/224Ra 222Rn/226Ra 228Ra/226Ra 224Ra/228RaMean 170 17.4 21.4 17.9 25 22,420 17,710 1.17 1.03Median 77 8.5 9.4 9.3 1 11,190 9,023 0.95 0.962σ 451 73.0 81.5 66.7 172 63,214 56,879 1.59 0.93Geometric 76 8.6 9.8 9.5 3 8,693 7,603 0.94 0.932 geom std 2.9 2.1 2.4 2.0 4.4 3.1 2.9 1.37 0.9Minimum <3 0.6 0.7 ND ND 194 201 0.15 0.20Maximum 1,113 324.0 348.6 234.1 593 170,900 247,800 4.26 2.60Rolesville Granite (n=52) Mean 335 26.4 36.9 26.4 55 36,450 24,700 0.76 1.08Median 249 11.6 18.6 12.3 5 17,910 12,060 0.68 0.982σ 503 100.6 109.1 89.3 245 80,981 76,245 1.24 0.89Geometric 249 14.0 22.4 13.9 9 17,776 11,114 0.62 1.012 geom std 1.7 1.9 1.9 2.0 4.3 2.8 2.8 1.24 0.8Minimum 23 2.4 1.4 2.0 ND 331 307 0.15 0.43Maximum 1,113 324.0 348.6 234.1 593 170,900 247,800 4.26 2.47Raleigh Gneiss (n=24) Mean 44 11.5 11.1 12.3 ND 13,680 10,880 1.23 0.89Median 30 5.6 6.0 5.5 ND 6,904 6,142 0.98 0.912σ 79 44.0 42.0 42.3 2 34,947 28,712 1.44 0.71Geometric 31 5.4 6.1 6.7 2 5,673 5,071 1.07 0.822 geom std 2.0 2.3 1.9 2.1 2.4 3.1 2.8 1.05 0.8Minimum <3 0.7 0.9 ND ND 194 201 0.51 0.35Maximum 152 108.7 104.9 102.7 5 76,890 60,170 3.11 1.86Metasedimentary & metavolcanic rocks (n=34)Mean 35 9.8 8.3 10.8 2 6,839 12,130 1.68 1.12Median 31 6.7 4.0 6.6 1 4,530 9,693 1.69 0.992σ 58 25.7 28.9 34.1 9 13,183 25,629 1.56 1.09Geometric 25 6.6 4.6 7.0 1 3,672 5,594 1.49 0.992 geom std 2.1 1.7 1.9 1.8 3.2 2.6 3.0 1.04 1.1Minimum <3 0.6 0.7 ND ND 245 247 0.39 0.20Maximum 96 76.0 83.3 96.3 17 23,470 50,830 3.53 2.60ND indicates not detected. a‐ Samples below detection excluded from geometric mean and geometric standard deviation
48
Table 6: Radionuclide activities and activity ratios.
Group
Well ID
U (nM)
224Ra (mBq L‐1)
226Ra (mBq L‐1)
228Ra (mBq L‐1)
222Rna
(Bq L‐1)
1 a 20 1 11.3 ± 1.1 5.1 ± 0.9 9.4 ± 1.9 49 1 a 22 17 6.1 ± 0.7 5.5 ± 0.7 5.6 ± 2.4 3 1 a 23 5 7.0 ± 0.8 4.0 ± 0.8 6.1 ± 2.3 23 1 a 24 ND 3.4 ± 0.4 4.9 ± 0.9 2.9 ± 1.9 79
1 a 25 1 12.5 ± 0.9 6.5 ± 1.1 10.7 ± 3.3 <3 1 a 26 1 1.2 ± 0.2 1.5 ± 0.9 4.2 ± 2.9 52 1 a 27 ND 2.3 ± 0.3 2.6 ± 0.9 4.9 ± 1.8 40 1 a 28 1 3.3 ± 0.4 1.6 ± 0.6 3.2 ± 2.3 7 1 a 29 1 6.4 ± 0.7 4.8 ± 1.0 17.0 ± 5.9 47 1 a 30 14 2.5 ± 0.3 0.7 ± 0.9 1.3 ± 1.9 32 1 a 31 6 13.3 ± 0.9 7.0 ± 0.9 13.8 ± 2.9 27 1 a 32 ND 12.1 ± 0.7 6.2 ± 1.1 12.7 ± 2.3 <3 1 a 34 ND 1.8 ± 0.3 3.7 ± 0.6 ND 40 1 a 35 1 1.6 ± 0.2 2.2 ± 0.7 4.6 ± 2.7 19 1 a 36 ND 5.7 ± 0.6 3.7 ± 0.7 9.3 ± 3.8 95 1 a 37 7 7.1 ± 0.7 1.6 ± 0.8 3.1 ± 3.3 50 1 a 38 16 20.7 ± 1.3 20.5 ± 1.6 8.0 ± 6.2 5 1 a 39 1 13.2 ± 1.2 22.1 ± 1.4 14.3 ± 4.3 14 1 a 40 ND 2.4 ± 0.3 1.2 ± 0.9 1.5 ± 2.4 11 1 a 48 1 4.4 ± 0.5 2.6 ± 0.6 2.7 ± 1.9 37 1 a 49 1 7.0 ± 0.7 5.0 ± 0.8 7.2 ± 4.8 81 1 b 19 2 13.7 ± 0.8 7.1 ± 1.1 12.3 ± 4.2 13 1 b 41 ND 6.2 ± 0.7 7.3 ± 1.1 8.1 ± 3.4 12 1 b 42 ND 13.0 ± 0.9 7.0 ± 1.0 21.8 ± 5.0 14 1 b 43 ND 3.7 ± 0.4 2.6 ± 1.0 6.1 ± 3.2 3 1 b 44 ND 10.2 ± 1.0 13.9 ± 1.1 14.8 ± 4.9 <3 1 b 45 ND 4.7 ± 0.6 2.9 ± 0.6 3.7 ± 2.4 48 1 b 50 ND 8.5 ± 1.0 4.0 ± 0.6 5.2 ± 1.9 8 1 b 51 3 8.0 ± 1.0 5.0 ± 0.9 8.5 ± 3.6 58 1 b 52 5 22.1 ± 1.7 11.3 ± 1.3 31.2 ± 7.1 32 1 b 53 1 2.4 ± 0.3 3.9 ± 0.8 2.8 ± 2.3 8 1 b 54 ND 1.6 ± 0.3 3.7 ± 0.7 3.9 ± 4.1 7 1 b 55 ND 0.6 ± 0.2 1.5 ± 0.5 3.1 ± 4.9 <3 1 b 56 1 5.0 ± 0.6 3.7 ± 0.6 7.0 ± 3.6 70 1 b 57 ND 76.0 ± 6.1 83.3 ± 3.8 96.3 ± 23.2 41 1 b 58 ND 15.4 ± 0.7 10.0 ± 1.1 17.0 ± 4.7 14 1 b 59 ND 23.0 ± 1.5 6.9 ± 0.9 23.5 ± 3.6 31 1 b 60 ND 8.2 ± 0.6 5.1 ± 0.9 9.4 ± 1.3 76 1 b 61 4 5.9 ± 0.8 1.3 ± 0.9 3.5 ± 2.3 66 1 b 62 ND 5.0 ± 0.7 2.5 ± 0.7 ND 5 1 b 63 ND 13.4 ± 1.1 2.7 ± 0.9 7.7 ± 3.9 66 1 b 74 1 4.5 ± 0.5 3.2 ± 0.8 3.5 ± 1.8 80
49
Group
Well ID
U (nM)
224Ra (mBq L‐1)
226Ra (mBq L‐1)
228Ra (mBq L‐1)
222Rna
(Bq L‐1)
1 b 75 1 19.4 ± 1.2 19.3 ± 1.2 41.3 ± 7.5 14 2 11 ND 5.9 ± 0.4 8.7 ± 1.7 6.3 ± 1.4 5
2 64 ND 71.4 ± 6.2 68.3 ± 2.3 83.1 ± 23.5 13 2 65 ND 8.4 ± 0.7 18.9 ± 1.6 14.6 ± 4.5 25 2 66 ND 9.6 ± 0.7 10.2 ± 1.2 5.2 ± 3.8 7 2 67 1 1.5 ± 0.2 3.4 ± 0.6 1.9 ± 1.6 25 2 68 ND 6.4 ± 0.8 2.9 ± 0.8 5.5 ± 1.5 24 2 69 ND 16.3 ± 1.2 21.0 ± 1.5 17.3 ± 5.6 57 2 70 ND 5.6 ± 0.8 10.3 ± 1.1 10.1 ± 2.3 79 2 71 ND 4.1 ± 0.5 4.8 ± 0.9 10.0 ± 4.4 <3 2 72 ND 14.8 ± 0.9 7.0 ± 0.9 13.2 ± 3.1 42 2 73 ND 3.2 ± 0.5 2.7 ± 0.6 ND 100 2 76 ND 1.7 ± 0.2 2.1 ± 0.5 2.7 ± 3.1 5 2 77 ND 10.8 ± 0.7 16.7 ± 1.4 9.8 ± 1.9 97 2 78 ND 3.6 ± 0.4 5.3 ± 0.9 6.3 ± 1.6 90 2 79 ND 0.7 ± 0.2 0.9 ± 0.6 1.1 ± 3.2 51 2 80 ND 4.0 ± 0.4 6.0 ± 1.2 4.4 ± 2.5 67 2 81 3 6.8 ± 0.4 7.2 ± 1.1 5.3 ± 2.1 152 2 82 ND 2.4 ± 0.3 2.9 ± 0.9 1.8 ± 3.0 78 2 83 ND 1.8 ± 0.3 3.5 ± 0.7 3.0 ± 2.6 30 2 84 ND 0.8 ± 0.2 2.2 ± 0.8 2.2 ± 2.4 13 2 85 ND 4.1 ± 0.5 5.0 ± 0.9 3.7 ± 9.2 20 b
2 86 ND 108.7 ± 6.6 104.9 ± 3.3 102.8 ± 8.9 21 3 a 7 6 9.3 ± 1.0 19.0 ± 1.3 10.3 ± 2.1 1070 3 a 8 102 11.1 ± 0.6 11.9 ± 1.1 10.0 ± 1.5 375 3 a 9 80 21.1 ± 2.4 58.4 ± 3.2 20.8 ± 1.7 705 3 a 10 19 9.5 ± 0.7 15.6 ± 1.4 10.9 ± 1.5 187 3 a 12 2 12.2 ± 0.6 10.1 ± 1.0 9.3 ± 2.2 168 3 a 13 ND 4.9 ± 0.6 15.3 ± 1.5 7.9 ± 1.9 371 3 a 14 2 18.9 ± 1.0 25.8 ± 1.8 19.8 ± 3.9 207 3 a 16 ND 14.4 ± 1.2 17.3 ± 1.7 12.4 ± 2.5 453 3 a 17 5 22.9 ± 1.4 18.6 ± 1.4 9.4 ± 2.1 225 3 a 18 83 45.8 ± 1.5 44.5 ± 2.5 31.2 ± 5.9 219 3 a 21 6 11.7 ± 1.1 23.6 ± 1.8 12.1 ± 2.3 259 3 a 33 8 7.5 ± 0.9 20.6 ± 1.6 17.0 ± 5.9 868 3 a 47 1 324.0 ± 15.2 348.6 ± 7.6 234.1 ± 34.5 107 3 a 89 22 29.5 ± 2.0 28.2 ± 1.7 13.8 ± 4.6 159 3 a 90 8 57.6 ± 3.1 75.4 ± 2.9 64.1 ± 10.7 395 3 a 92 28 6.5 ± 1.1 18.5 ± 1.4 5.0 ± 3.2 1113 3 a 93 ND 9.2 ± 0.8 14.9 ± 1.2 14.3 ± 4.3 499 3 a 95 ND 6.8 ± 0.7 14.2 ± 1.4 10.3 ± 3.6 368 3 a 96 ND 8.6 ± 0.8 17.7 ± 1.6 11.1 ± 3.7 118 3 a 97 2 16.9 ± 1.6 16.8 ± 1.5 17.6 ± 5.4 443 3 a 99 36 9.8 ± 1.7 21.6 ± 1.9 10.1 ± 5.7 679
50
Group
Well ID
U (nM)
224Ra (mBq L‐1)
226Ra (mBq L‐1)
228Ra (mBq L‐1)
222Rna
(Bq L‐1)
3 b 1 1 9.2 ± 0.9 48.0 ± 2.2 14.3 ± 6.0 23 3 b 2 60 22.0 ± 1.5 41.1 ± 2.7 23.8 ± 2.2 295 3 b 3 340 22.0 ± 2.6 65.5 ± 2.7 15.1 ± 3.2 310 3 b 4 4 30.6 ± 1.6 33.4 ± 1.7 44.3 ± 2.6 70 3 b 5 3 2.4 ± 0.3 6.1 ± 0.8 2.0 ± 1.5 92 3 b 6 489 7.0 ± 1.1 84.1 ± 2.6 16.3 ± 2.1 64 3 b 15 90 6.7 ± 1.0 32.5 ± 1.9 4.8 ± 2.9 240 3 b 46 2 31.1 ± 1.8 14.4 ± 1.4 31.1 ± 6.5 130 3 b 87 9 14.4 ± 1.0 25.8 ± 1.7 20.3 ± 5.5 221 3 b 88 1 9.6 ± 0.8 13.9 ± 1.4 6.9 ± 2.4 103 3 b 91 593 10.4 ± 1.0 42.3 ± 2.0 13.9 ± 4.5 163 3 b 94 ND 12.3 ± 1.2 33.2 ± 1.7 25.3 ± 8.1 30 3 b 98 107 16.7 ± 0.8 29.1 ± 2.1 6.8 ± 2.5 265 3 c 100 4 7.7 ± 0.9 6.6 ± 0.9 5.5 ± 3.1 638 3 c 101 9 47.6 ± 3.4 56.5 ± 2.8 68.4 ± 17.7 117 3 c 102 348 103.4 ± 6.5 184.3 ± 5.8 132.7 ± 26.7 156 3 c 103 ND 4.7 ± 0.4 7.1 ± 1.3 4.5 ± 3.2 364 3 c 104 97 164.9 ± 8.5 129.1 ± 4.8 199.6 ± 34.8 227 b
3 c 105 174 4.6 ± 0.4 1.4 ± 1.1 6.1 ± 6.6 358 3 c 106 4 14.5 ± 2.0 31.4 ± 2.0 15.3 ± 7.4 380 3 c 107 21 4.1 ± 0.5 6.1 ± 0.9 3.2 ± 2.5 188 3 c 108 34 74.8 ± 5.2 69.7 ± 3.5 75.8 ± 17.7 223 3 c 109 2 9.1 ± 1.9 13.0 ± 1.3 5.5 ± 4.0 594 3 c 110 ND 22.5 ± 1.7 18.0 ± 1.5 17.1 ± 6.1 135 3 c 111 2 11.6 ± 1.0 12.0 ± 1.3 9.6 ± 3.4 187 3 c 112 ND 13.1 ± 1.1 14.6 ± 1.4 18.9 ± 5.9 695 3 c 113 4 16.8 ± 1.3 16.7 ± 1.6 11.8 ± 4.2 411 3 c 114 2 7.9 ± 0.9 16.7 ± 1.4 7.6 ± 3.5 471 3 c 115 36 4.2 ± 0.6 6.5 ± 1.1 3.1 ± 2.4 548 3 c 116 9 4.4 ± 1.1 13.6 ± 1.1 2.7 ± 3.0 643 3 c 117 1 6.3 ± 0.5 8.1 ± 1.2 8.2 ± 2.7 78
Errors are ±2σ from counting statistics. ND indicates not detected; these were treated as zero activities in data analysis. a‐ Values less than the reporting limit of 3 Bq L‐1 were treated as zero values for data analysis b‐ Rn sample collected on a later date than indicated on Table 4.
51
Table 7: Radium isotope activities in Rolesville Granite samples from cores WC‐1 and WC‐3. Errors are ±2σ from counting statistics.
Sample 226Ra
(Bq kg‐1)
228Ra (Bq kg‐1)
228Ra/226Ra
WC‐1‐41.4 43.0 ± 5.6 42.4 ± 4.6 0.99 WC‐1‐59.6 31.3 ± 3.8 61.5 ± 3.8 1.97 duplicate can 27.3 ± 2.5 53.3 ± 2.5 1.95 WC‐1‐70.0 53.5 ± 5.6 56.5 ± 5.0 1.06 duplicate can 51.9 ± 5.7 54.8 ± 5.2 1.06 WC‐3‐48.5 44.6 ± 4.4 72.1 ± 4.2 1.62 WC‐3‐66.3 20.8 ± 5.8 55.6 ± 5.8 2.68 WC‐3‐117.3 47.8 ± 5.9 51.9 ± 4.9 1.09
52
3. Arsenic and other oxyanion‐forming trace elements in an alluvial basin aquifer: Evaluating sources and mobilization by isotopic tracers (Sr, B, S, O, H, Ra)
3.1 Introduction
Occurrence of oxyanion‐forming trace elements, primarily arsenic (As), imposes
limitations on the use of groundwater from alluvial basin aquifers in the semi‐arid western
United States (Robertson, 1989; Welch et al., 2000; Smedley and Kinniburgh, 2002). Population
growth, degradation of existing water resources, and the 2001 reduction of the U.S.
Environmental Protection Agency (EPA) maximum contaminant level (MCL) for arsenic to
10 µg L‐1 have emphasized the need to provide sustainable sources of high‐quality drinking
water for the region. Other related elements may co‐occur with As and form oxyanions that are
mobile under similar conditions, including boron (B), chromium (Cr), vanadium (V), molybdenum
(Mo), selenium (Se), and antimony (Sb). Numerous published studies have implicated both
lithologic and geochemical controls on these elements. Volcanic rocks, especially silicic volcanic
rocks, and clastic sediments derived from volcanic rocks contain leachable As‐bearing phases
that yield significantly higher As than other contributing sediment sources to groundwater
(Welch et al., 2000; Smedley and Kinniburgh, 2002). In addition, the mobility of oxyanion‐
forming trace elements is enhanced by the oxic and near‐neutral to slightly basic (pH ~6.5‐9.5)
conditions in alluvial basin‐fill aquifers (Robertson, 1989). Under these conditions, oxyanion‐
forming trace elements are dominantly found in soluble, oxidized species including HAsO42‐,
H2AsO4‐, MoO4
2‐, CrO42‐ , SeO4
2‐, VO3OH2‐, VO2(OH)
2‐, and Sb(OH)6‐ (Baes and Mesmer, 1976; Fujii
and Swain, 1995; Amrhein et al., 1998). Potential adsorption sites for these oxyanions (e.g. Fe
53
oxides) are present on aquifer solids in oxic, near‐neutral conditions, with desorption promoted
at increasing pH (e.g. Robertson, 1989; Scanlon et al., 2009).
This study focuses on groundwater in Willcox Basin, a typical alluvial basin in semi‐arid
southeastern Arizona (Figure 11). The groundwater in this basin is characterized by
concentrations of As commonly exceeding 10 µg L‐1 (Towne and Freark, 2001), which is typical
for basins of the region (Robertson, 1989; Coes et al., 2002). Because the observed As
concentration in groundwater is a function of lithologic source, weathering, and
adsorption/desorption equilibrium, the purpose of this study is to apply several isotopic tracers,
including strontium, boron, sulfur, and radium isotopes, in order to evaluate (1) the variations in
sediment source that affect oxyanion‐forming natural trace elements’ abundance in
groundwater in Willcox Basin and (2) the mechanisms that control their mobilization from the
aquifer lithology.
Because there is only one naturally‐occurring isotope of As, and because O isotopes of
arsenate and arsenite ions are subject to rapid isotopic exchange with water limiting their use as
process tracers (Larese‐Casanova and Blake, 2010), tracing of As sources could be obtained
indirectly using other isotopic proxies that are sensitive to potential geologic sources of
oxyanion‐forming elements and geochemical mobilization processes. The ratio of radiogenic to
nonradiogenic strontium (87Sr/86Sr) is a sensitive tracer for sediment sources and silicate
weathering processes in groundwater systems (Franklyn et al., 1991; Bullen et al., 1996; McNutt,
1999; Harrington and Herczeg, 2003). It is hypothesized that Sr would not directly relate to
geochemical processes controlling oxyanions because it is a divalent cation, but rather would
indicate the origin of sediments (Sr sources) within alluvial basins. Due to the presence of
Proterozoic granitic rocks containing K‐feldpar adjacent to the Willcox Basin, it is inferred that
54
highly radiogenic Sr would be present in sediments derived from these older materials, whereas
the Tertiary silicic volcanic rocks would be characterized by unradiogenic Sr isotope
composition. The boron isotope ratio is another tracer that could provide indication for the
origin of oxyanion‐forming elements and/or chemical controls on these elements. For example,
boron could possibly be used as an analogue to arsenic because both exhibit pH‐dependent
adsorption‐desorption reactions that are associated with B isotope fractionation and because As
and B commonly co‐occur in volcanic‐derived materials (e.g. Welch and Lico, 1998; Fujii and
Swain, 1995; Scanlon et al., 2009). A major difference between B and As(V) is that the dominant
B species at neutral pH is the uncharged B(OH)3, whereas the dominant As(V) species are the
charged H2AsO4‐ (below pH ~7) and HAsO4
2‐ (above pH ~7). The potential application of sulfur
isotopes is that some of the sulfate in the system is hypothesized to be derived from oxidation
of volcanic sulfide, which is also a possible As source. In addition, S isotopes can diagnose the
presence of sulfate reduction, which could influence oxyanion mobility by converting oxyanion‐
forming trace elements to less soluble or uncharged forms. Radium and uranium isotope
occurrence may supply additional understanding of oxyanionic trace elements because uranium
is a mobile anion in oxic conditions, primarily as carbonate complexes, whereas the major
adsorption sites of radium (e.g. Mn and Fe oxides) are most stable in oxic conditions. The state
of these metal oxides can influence both radionuclide activities and oxyanion concentrations.
Because metal oxides can adsorb As(V) and Ra at near‐neutral pH, the disequilibrium between
226Ra and 238U in the 238U decay series permits inferences on the reactivity of the U‐decay chain
nuclides and indirectly also on sinks of other oxyanion‐forming trace elements. Additionally,
isotopes of the water molecule (3H, δ2H, and δ18O) are reported to provide information on
groundwater residence time and the relationship to modern precipitation.
55
Figure 11: Map of Willcox Basin showing the basin’s extent, bedrock types in the mountain blocks, sample numbers, and 87Sr/86Sr ratios. Bedrock types derived from Richard et al. (2002); location of Apache Pass Fault Zone derived from Drewes (1985, 1986) and Erickson (1988). Approximate extent of cones of depression and groundwater flow directions derived water level contours in Oram (1993). Inset shows location of Willcox Basin in the Basin and Range region (USGS, Principal aquifers of the 48 conterminous United States, Hawaii, Puerto Rico,
and the U.S. Virgin Islands, obtained from http://www.nationalatlas.gov).
3.2 Geologic setting
The complex geologic history of the Willcox Basin affects both groundwater evolution
and the availability of oxyanion‐forming trace elements. As a tectonically and hydrologically
closed basin, isotopic end members in the Willcox Basin are locally defined and represent a
mixture of the mobile sources exposed in the basin.
56
3.2.1 Composition of the mountain blocks
The mountain blocks bounding the Willcox Basin and other tectonic basins of southern
Arizona contain a complex assortment of crystalline and sedimentary rocks from ~1.7 Ga to
Tertiary‐Quaternary volcanic rocks. Briefly, the main rock types surrounding the Willcox Basin
include: (1) 1.7 Ga metavolcanic and metasedimentary rocks (felsic tuff, basalt, quartzite,
sandstone, and conglomerate; Condie et al., 1985); (2) 1.4 Ga granitic rocks (Erickson, 1981); (3)
Cambrian through Triassic sedimentary rocks (conglomerate, sandstone, and carbonate rocks;
(Hayes, 1978; Drewes, 1985); (4) Late Cretaceous to early Tertiary volcanic (felsic to
intermediate tuff and breccias) and intrusive rocks (granite and granodiorite; Drewes et al.,
1988); and (5) mid‐Tertiary felsic volcanic and intrusive rocks (Cooper, 1960; Drewes, 1996;
Figure 11). Of these, the Tertiary volcanic rocks and Proterozoic rocks are the most extensive
surrounding the Willcox Basin and are closest to the sample locations in this study (Figure 11).
The presence of outliers of Tertiary rocks in the southern Willcox Basin (Figure 11) implies that
these rocks are also abundant in the subsurface. Although it may be inferred that the fluvial
basin‐fill sediments are dominated by nearby bedrock lithologies, there is inadequate
subsurface data to characterize the spatial variation in the composition of basin‐fill sediments.
3.2.2 Basin geometry
Depth to bedrock increases steeply from the edges of the Willcox Basin and from
bedrock outliers (Drewes, 1986; Drewes et al., 1988; Richard et al., 2007). At its center, the
Willcox Basin exceeds 1500 m depth (Richard et al., 2007). Overall, the depth of basin‐fill
deposits greatly exceeds the depth of most of the wells in this study. The Apache Pass Fault
(Figure 11), a major Precambrian structural feature, is preserved in the Dos Cabezas Mountains
57
and beneath the eastern portion of the Willcox Basin. This fault has been hypothesized to supply
high‐As water to wells along its trace (Towne and Freark, 2001), and, in general, basement‐
penetrating faults in the southwestern USA have been observed locally to contribute waters
with high levels of salinity, CO2, and/or natural trace elements affecting water quality to basin‐
fill aquifers (e.g. Coes et al., 2002; Newell et al., 2005; Earman et al., 2008; Crossey et al., 2009).
3.2.3 Pleistocene‐Holocene Lake Cochise
The surficial deposits at the center of the Willcox Basin represent periodic filling and
drying of a large closed‐basin lake during the Pleistocene and early Holocene epochs. A
distinctive shoreline deposit at 1274 m elevation (14 m above the present dry lake bed) formed
approximately 13.5 ka and lower high stands occurred during the early Holocene (Waters,
1989). At least 40 m of black, sulfidic clay is beneath the present‐day dry lake bed (Schreiber et
al., 1972; Vine et al., 1979). At greater depth (up to 300 m), the sediments become dominated
by alternating calcite and gypsum layers (Vine et al., 1979), consistent with evaporation at the
center the basin being a major drain for the groundwater system during the basin’s history.
3.2.4 Groundwater flow and residence time
Based on relationships between precipitation and water levels, Gardner and Heilweil
(2009) suggested that mountain‐front recharge and infiltration from ephemeral streams near
the basin edges are the major source of recharge to the basin, and that natural groundwater
flow directions follow the surface topography of the Willcox Basin. At the center of the basin,
fine‐grained lake bed deposits locally create a shallow (<30 m depth), perched groundwater
system (Oram, 1993) and may be responsible for the historical artesian conditions in the
underlying basin‐fill aquifer (Meinzer et al., 1913). Although historically, the main discharge
58
point for the hydrologically‐closed basin has been evapotranspiration at Willcox Playa, a more
significant mechanism today is water withdrawals for irrigation. The main irrigated areas,
located north and southeast of the Willcox Playa, are the sites of large cones of depression
(Mann et al., 1978; Oram, 1993; Arizona Department of Water Resources, 2005; Gardner and
Heilweil, 2009) that significantly affect groundwater flow (Figure 11). Quantitative data on
groundwater residence time are sparse; however, one published carbon‐14 activity near the
center of the basin (26% modern carbon; Robertson, 1991) is consistent with Holocene‐aged
groundwater, rather than the basin‐fill groundwater being solely a remnant of elevated
Pleistocene recharge.
3.3 Methods
3.3.1 Well sampling
Samples were collected from 15 domestic and high‐capacity irrigation wells and one
spring. Wells were sampled during a period of water use at a tap as close to the wellhead as
possible. Well depths are known in some cases (Table 8), but the lengths of well screens
probably vary significantly among the irrigation and domestic drinking water wells sampled for
this study. pH, temperature, and dissolved oxygen were measured in the field using meters that
were calibrated daily. Trace metal, cation, anion, and Sr and B isotope samples were filtered in
the field using 0.45 µm nylon filters. Major cation/trace metal samples were filtered directly into
new polyethylene bottles that had been cleaned with trace metal grade HCl and HNO3, then
rinsed with deionized water having resistivity > 18 MΩ cm‐1. Major cation/trace metal samples
were immediately acidified with high‐purity HNO3 (Fisher Optima). Sample aliquots for anion
and water stable isotope analyses were stored in DI‐washed HDPE bottles and glass scintillation
59
vials, respectively, with no headspace. Unfiltered samples were also collected for tritium and
sulfur isotope analysis; nitric acid was added to the S isotope samples to prevent sulfate
reduction. All samples (except radium) were kept on ice in the field and refrigerated in the
laboratory prior to analysis. 40 liter unfiltered samples were collected for radium isotopes in
completely filled polyethylene containers and kept closed until processing upon return to the
laboratory later the same day. An acidified blank was prepared by filling an acid‐washed sample
bottle with unfiltered deionized water (> 18 MΩ cm‐1) in the lab, with the same subsequent
transportation, acid addition, and handling as water samples.
3.3.2 Major and trace element analysis
Major cation concentrations (Na, K, Ca, Mg, Si) and Fe were determined using a Perkin
Elmer Optima 5300 inductively‐coupled plasma optical emission spectrometer, and major anion
concentrations (F‐, Cl‐, Br‐, NO3‐, SO4
2‐) were analyzed using a Dionex ICS‐3000 ion
chromatograph with AS23 analytical column. Alkalinity concentrations were determined by the
Gran‐Alk titration within 24 h of sampling (Gieskes and Rogers, 1973). Trace element (B, Sr, V,
Cr, Mn, As, Se, Mo, Sb, Ba, and U) concentrations were determined by inductively‐coupled
plasma mass spectrometry on a VG Plasmaquad 3 instrument. Calibration standards were
prepared from serial dilution of certified trace metal solution NIST1643e. Samples and standards
were diluted to similar proportions with an internal standard solution of 10ppb In, Tm, and Bi in
2% HNO3 to monitor and correct for drift. Dilutions were carried out with solutions prepared
from quartz‐distilled deionized water and nitric acid. Instrument drift was also monitored and
corrected for each element by analysis of one of the calibration standards at regular intervals.
To allow for U analysis, NIST1643e was spiked with a plasma‐grade single‐element U solution
60
prior to serial dilution. Trace element analysis included a field blank for which resulting
concentrations were <1 µg L‐1 except B (1.0 µg L‐1), Sr (2.1 µg L‐1), and Ba (2.9 µg L‐1). Boron and
arsenic speciation were calculated using the PHREEQC geochemical code (Parkhurst and Appelo,
1999). The code was directed to use dissolved oxygen measurements to calculate pe under
assumptions of redox equilibrium. Arsenic speciation was not calculated for well 2 due to the
anoxic nature of the well.
3.3.3 Stable and radiogenic isotopes
Hydrogen isotope analysis was conducted by Cr metal reduction at 750°C (Gehre et al.,
1996) and oxygen analysis by CO2 equilibration at 15°C (Craig, 1957), both on a Finnigan Delta S
gas‐source isotope ratio mass spectrometer at the University of Arizona. Isotope ratios were
standardized relative to VSMOW and VSLAP and are reported as per mil relative to VSMOW.
Precision is within 0.9‰ for δ2H and 0.08‰ for δ18O based on replicate analyses. Sulfur isotope
ratios of sulfate were analyzed using a ThermoQuest Finnigan Delta PlusXL continuous flow
isotope ratio mass spectrometer at the University of Arizona. 34S/32S ratios are normalized to the
CDT standard and reported as δ34S with precision of 0.15‰. Tritium activities were analyzed at
the University of Miami Tritium Laboratory by gas proportional counting with precision (1σ) of
approximately 0.1 tritium unit (TU).
Due to the generally low B concentrations, samples were prepared for B isotope analysis
by preconcentrating filtered, non‐acidified water samples to approximately 500 µg B L‐1 at ~40°C
in ultra‐clean Teflon vials in a laminar‐flow hood using B‐free filters and treating with 30% H2O2
(Fisher Optima) to oxidize labile organic matter. Approximately 2 ng B (less in the case of low‐B
samples) was loaded directly onto degassed Re filaments in a B‐free synthetic seawater solution
61
prepared from plasma‐grade standards (Dwyer and Vengosh, 2008) and analyzed by negative
ion thermal ionization mass spectrometry (TIMS) on a ThermoFisher TRITON thermal ionization
mass spectrometer at Duke University. Samples were analyzed by stabilizing 11BO2 signal at a
maximum of ~1 V (significantly less in the case of low‐B samples), then 11B/10B ratios were
measured for approximately 30 minutes (60 cycles). B isotope ratios were normalized to the
long‐term average 11B/10B of the NIST 951 standard (4.0067±0.25‰; 1σ, n=238) and reported as
δ11B. A two‐step protocol was used to accept measured 11B/10B ratios: (1) Potential interference
at mass 42 (from CNO) was monitored by verifying CN at mass 26 (Hemming and Hanson, 1992)
to be < 5000 counts per second using the secondary electron multiplier. Although the
relationship between CN and CNO in negative‐ion TIMS is not completely understood, the CN
signal on mass 26 is used as a proxy for possible interference at mass 42 and ratios are rejected
in cases of CN > 5000 counts per second; and (2) stable ratios (maximum drift < ~1 ‰) during
data acquisition, including samples for which temperature varied during acquisition. In addition
to replicate analysis of the NIST 951 standard, five representative samples from this project
were subjected to whole‐process replication, yielding four replicates within a range of 1.8 ‰,
and one sample rejected for less favorable replication. The overall precision is therefore
estimated to be 1‐2‰.
Sr isotope samples were prepared by dry‐down, HNO3 digestion, and subsequent
column extraction of ~3 µg Sr using Eichrom Sr‐specific resin conditioned with HNO3 and eluted
with quartz‐distilled deionized water. The resulting Sr was loaded onto single degassed Re
filaments in a Ta load solution (Birck, 1986) and analyzed by TIMS in positive ion mode at Duke
University. 87Sr/86Sr ratios are reported without normalization to the NIST 987 standard; long‐
term average 87Sr/86Sr of NIST 987 during the period preceding and including this study was
62
0.710245 ± 0.000010 (1σ, n=109). Whole‐process replicates (n=4) indicate precision within the
expected analytical uncertainty reported for NIST 987.
3.3.4 Radium isotopes
Radium was preconcentrated onto two sequential plastic columns containing 10 g (dry
weight) of manganese oxide‐impregnated acrylic fibers (Moore and Reid, 1973) at a flow rate of
~500 mL min‐1 and analyzed at Duke University within 2 d of sampling. Radium‐224 was analyzed
by alpha counting of its daughter 220Rn using a delayed coincidence counter (Moore and Arnold,
1996; Moore, 2008; Garcia‐Solsona et al., 2008) at a detector efficiency designed for hand‐
squeezed fibers. Low‐activity samples were quantified by 220Rn coincidence counts and high‐
activity samples by total alpha counts. In both methods, count rates were corrected for counts
attributable to 223Ra, 224Ra decay, and background count rate. High activity (>20 cpm samples)
were re‐analyzed ~1 week later for quantification of 223Ra interference (Moore and Arnold,
1996). For radium‐226 analysis, fibers were sealed into an evacuated glass tube for 20 days,
then analyzed for its daughter 222Rn using a Durridge RAD7 radon‐in‐air monitor (Kim et al.,
2001) calibrated using a NIST radium‐226 solution transferred onto Mn oxide fibers and
corrected for background count rate. For radium‐228 analysis, fibers (first column only) were
compressed into a wafer (5 mm tall and 65 mm in diameter) inside a metal can using a hydraulic
jack and analyzed using the 911 keV peak of 228Ac on a Canberra broad‐energy germanium
gamma detector, corrected for background count rate and the fraction of radium on the first
column, and calibrated using an average of three 232Th (equilibrated with 228Ra) standards
loaded to resemble the geometry of the compressed fibers. Stated 2σ errors represent counting
statistics only, following published methods for error calculation and propagation of error
63
(Eaton et al., 2005; Garcia‐Solsona et al., 2008) and thus do not incorporate other potential
uncertainties including moisture content (224Ra) or counting geometry within the compressed
disc (228Ra). The detection limit, defined as the activity with 95% probability of being detected at
±100% precision, is estimated by 1.96σ (Eaton et al., 2005), so the quoted 2σ errors provide an
estimate of the sample‐specific detection limit for the counting time used.
3.4 Results
3.4.1 Groundwater salinity and major element chemistry
The 16 wells in the study yielded a large range of salinity with total dissolved solids of
156‐1147 mg L‐1 (Table 8). Groundwater with the highest dissolved solids is located near the
center of Willcox Basin while groundwater with the lowest dissolved solids is in the mountain
blocks and upper elevations of the basin‐fill aquifer. However, salinity does not systematically
increase down‐gradient through the Willcox Basin groundwater system. In general, groundwater
from the mountain blocks and upland areas is characterized by a Ca‐Mg‐HCO3‐ composition, and
down‐gradient waters in the basin are composed of mix of dominant cations and anions, as well
as a distinctive Na‐HCO3‐ water containing very low concentrations of Ca and Mg (e.g. well 8;
Table 8). Among the widely varied anion concentrations include local values of nitrate
(>10 mg L‐1 as N, n=1) and fluoride (>4 mg L‐1, n=3) in excess of EPA standards. Temperature and
pH exhibit a wide range of variability within the basin‐fill aquifer (20‐33 °C and 6.8‐9.3,
respectively; Table 8). Neither pH nor temperature is clearly associated with location within the
basin. Almost all of the wells contain measurable dissolved oxygen (median 63% atmospheric
saturation, range 3‐95%).
64
3.4.2 Trace metals
Oxyanion‐forming metal concentrations (Table 9) occupy a wide range of
concentrations: V 0.6‐21.7 µg L‐1, Cr 0.8‐26.2 µg L‐1, As 0.3‐29.4 µg L‐1, Se 0.2‐10.2 µg L‐1, Mo 1.0‐
11.1 µg L‐1, and Sb 0.3‐1.9 µg L‐1. Among the oxyanion‐forming trace elements, only As exceeded
the U.S. EPA MCL for public drinking water systems (10 µg L‐1, n=4). Arsenic is also associated (ρ
> 0.50) with several major chemical parameters: SO42‐ (ρ = 0.76), F‐ (ρ = 0.74), Na (ρ = 0.74), Cl‐ (ρ
= 0.67), NO3‐ (ρ = 0.58), and Ca/Na (ρ = ‐0.67). Notably, As is not correlated with pH (ρ = 0.11).
Among trace elements, As is correlated with V (ρ = 0.87), Cr (ρ = 0.73), Se (ρ = 0.66), Mo (ρ =
0.62), B (ρ = 0.58), and Sb (ρ = 0.54).
3.4.3 Environmental isotopes (δ2H, δ18O, 3H, 87Sr/86Sr, δ11B, δ34S)
Groundwater from the Willcox Basin and its mountain blocks exhibits some variation in
hydrogen and oxygen isotope ratios (δ2H ‐85 to ‐66‰, δ18O ‐11.8 to ‐8.6‰; Figure 12). In
general, waters near the edges of the Willcox Basin exhibit similar isotopic values as the limited
data for the mountain blocks (generally < ‐10‰). A slight apparent effect was observed in which
higher values of δ18O (> ‐10‰) occur in the southern and western samples, located near the
center of the basin at relatively low elevation (Figure 13). All waters analyzed for tritium are
consistently near the detection limit (≤0.6 tritium units (TU); Table 10).
65
Figure 12: Hydrogen and oxygen isotope ratios in Willcox Basin groundwater. Dashed line represents the global meteoric water line (GMWL; δ2H = 8 × δ18O + 10‰; Craig, 1961).
A wide geographic variability was observed among Sr isotope ratios, generally
decreasing from exceedingly radiogenic values up to 0.75010 in groundwater from the
crystalline mountain blocks, to a minimum value of 0.71012 at the center of the basin (Figure
11). Waters with 87Sr/86Sr > 0.72 in the northern and southeastern parts of the basin (Figure 11)
are associated with low concentrations of oxyanion‐forming trace elements (B, As, V, Cr, Se, Sb,
Mo) and high Ca/Na (Figure 14), and are referred to as “trend 1”. Samples in the central part of
the basin exhibiting more significant oxyanion concentrations are characterized by 87Sr/86Sr <
0.72 and generally lower Ca/Na (Figure 15), and are referred to as “trend 2”. Boron isotope
ratios (as δ11B) vary significantly (‐7.0‰ to 18.2‰), and the higher values of δ11B are somewhat
associated with lower B concentrations (Figure 16a). B isotope ratios exhibit some geographic
trends also in that the samples with the highest 87Sr/86Sr and Ca/Na (trend 1) tend to exhibit
somewhat lower As and B concentrations and higher δ11B than trend 2 waters (Figure 16a‐c).
66
Figure 13: Map showing δ18O of water samples in relationship to basin topography. Elevation data obtained from U.S. Geological Survey (http://seamless.usgs.gov).
Sulfur isotope ratios (as δ34S) vary from 2.0‰ to 13.4‰. The results indicate that trend 1 waters
exhibit generally lower δ34S values than trend 2 waters (Figure 17a) and that sulfate‐dominated
waters (SO42‐:Cl‐ > 1 and Ca/SO4
2‐ < 2) exhibit δ34S in the narrow range of 8‐10‰ (Figure 17b‐c).
3.4.4 Radionuclides
With the exception of samples in the mountain blocks (sites 1‐3), Ra isotope activities
are exceedingly low (Table 10), 1‐2 orders of magnitude below the EPA maximum contaminant
level. Ra isotope activities are essentially uncorrelated with salinity, pH, and oxyanion‐forming
trace elements. 224Ra, 226Ra, and 228Ra are reasonably well correlated with Ca (ρ = 0.59‐0.76), Mg
(ρ = 0.39‐0.71), and Sr (ρ = 0.60‐0.72), but not with radium’s chemical analogue Ba (ρ = 0.05‐
0.32). U is present at low to moderate concentrations (0.3‐13 µg L‐1) throughout the basin‐fill,
and concentrations are much higher in the mountain blocks (10‐71 µg L‐1; Table 10). One
67
Figure 14: Relationships between 87Sr/86Sr and oxyanion‐forming trace element concentrations and Sr concentration.
68
mountain block water sample yielded radionuclide levels in excess of EPA drinking water
standards for U (30 µg L‐1) and Ra (185 mBq L‐1 226Ra + 228Ra).
Figure 15: Relationship between Ca/Na and 87Sr/86Sr and arsenic concentration.
Figure 16: δ11B in relation to B concentration, 87Sr/86Sr, and As concentration. Note the apparent separation between trend 1 plus the mountain blocks (low B concentration, high
δ11B, high 87Sr/86Sr, low As) and trend 2 (high B concentration, low δ11B low 87Sr/86Sr, elevated As). Black bar represents the approximate δ11B range of rhyolite and granite (Barth, 1993).
69
3.5 Discussion
3.5.1 Recharge and groundwater residence time
Hydrogen and oxygen isotope ratios (as δ2H and δ18O) of groundwater in the Willcox
Basin plot along a linear trend (Figure 12) that is along or slightly below the global meteoric
water line, consistent with groundwaters in adjacent alluvial basins in Arizona (Baillie et al.,
2007; Earman et al., 2008; Gu et al., 2008; Adkins, 2009). Specifically, the majority of δ18O values
are within the range of winter precipitation reported from adjacent mountain blocks (‐12.1 to
‐8.7‰) whereas summer precipitation exhibits higher δ18O values (‐9.5 to ‐6.8‰; Earman et al.,
2006; Adkins, 2009). Overall, δ2H and δ18O values of Willcox Basin groundwaters record waters
derived from local (basin floor) winter precipitation and/or mountain‐front recharge of higher‐
elevation precipitation, as in adjacent basins (Earman et al., 2008; Adkins, 2009). There is no
isotopic evidence of substantial evaporation within the basin (Figure 12) but some evidence of
higher‐elevation δ18O values near the basin’s northern and eastern edges (Figure 13) . Tritium
activities consistently <1 TU in the basin‐fill sediments (Table 10) indicate that precipitation is
not a source of rapid recharge on the scale of years to decades, which is consistent with the
semi‐arid nature of the basin. In places, the low tritium activities contradict other indications
from the data, such as elevated nitrate concentrations in irrigated areas, which implies
significant recharge from the surface. The most likely explanation is that recharge of tritium‐free
irrigation water derived from groundwater pumping conveys nitrate to the basin‐fill aquifer. The
slow replenishment of the basin‐fill aquifer by precipitation is indicated by large water table
declines in irrigated areas since the 1940s, although these declines have been locally reversed as
irrigation withdrawals have been reduced since the 1970s (Mann et al., 1978; Oram, 1993).
70
Figure 17: δ34SSO4 in relation to sulfate concentration, SO42‐/Cl‐ ratio, and Ca/SO4
2‐ ratio.
71
Table 8: Major element results and mineral saturation index values.
Well Elevation Depth pH Temp Diss.O2 TDS Ca Mg Na K Si F‐ Cl‐ NO3‐ SO4
2‐ Alkalinity
(m) (m) (°C) (% saturation) (mg L‐1) (mM) (meq L‐1)
Mountain blocks 1 2769 Spring 7.12 10.1 55 161 0.6 0.2 0.3 0.01 0.3 0.03 0.07 <0.01 0.04 1.68 2 1545 92 7.00 21.4 3 704 2.0 0.9 4.0 0.08 0.4 0.08 2.83 0.19 0.81 4.80 3 1512 7.12 27.0 67 586 0.9 0.8 4.2 0.05 0.4 0.04 0.49 <0.01 0.36 5.90
Willcox Basin – trend 1 4 1446 7.13 27.3 95 352 1.0 0.5 1.1 0.04 0.7 0.07 0.35 0.07 0.15 3.23 5 1409 9.12 29.5 63 347 1.2 0.4 1.2 0.04 0.5 0.07 0.57 0.09 0.24 2.94 6 1342 153 6.67 20.7 122 156 0.4 0.2 0.5 0.02 0.6 0.02 0.15 0.18 0.05 1.12 7 1332 138 8.36 27.8 20 362 0.8 0.3 1.9 0.06 0.5 0.07 0.28 0.05 0.17 3.49
Willcox Basin – trend 2 8 1288 86 9.27 23.0 639 0.03 0.002 7.8 0.05 1.2 0.42 1.39 0.19 1.41 2.96 9 1310 183 7.76 32.9 83 342 0.3 0.02 3.2 0.04 0.7 0.28 0.24 0.06 0.67 2.13
10 1303 245 7.54 28.4 71 452 0.4 0.1 5.2 0.04 0.4 0.04 2.88 0.09 0.37 2.35 11 1297 168 7.76 23.1 64 378 1.2 0.2 2.0 0.06 0.5 0.10 0.58 0.08 1.00 2.02 12 1294 245 7.94 31.7 53 950 1.0 0.1 12.6 0.06 0.4 0.06 11.77 0.06 1.05 1.19 13 1284 80 7.95 22.4 83 795 2.8 0.4 4.3 0.09 0.6 0.07 1.06 0.72 2.89 2.91 14 1274 7.08 21.3 51 600 0.8 0.3 5.0 0.07 0.7 0.26 1.03 0.30 0.61 4.61 15 1272 35 6.75 24.5 21 722 2.1 0.6 3.9 0.05 1.0 0.12 2.56 0.30 1.29 3.86 16 1266 7.11 19.8 50 1147 2.2 2.0 8.4 0.13 0.5 0.07 7.65 0.07 2.99 3.62
72
Table 9: Concentrations of oxyanion‐forming and other trace metals. See Table 10 for B and Sr concentrations.
Concentrations in µg L‐1 Proportion of As
Well V Cr Mn Fe1 As as HAsO42‐ Se Mo Sb Ba
Mountain blocks 1 1.0 1.0 0.2 10 0.3 0.62 0.2 1.3 0.3 18.32 2.0 3.8 15.7 18 1.5 2.2 11.1 2.0 81.43 0.6 0.9 0.7 ND 0.4 0.67 0.6 3.8 0.3 49.8
Willcox Basin – trend 1 4 1.6 0.8 0.6 ND 0.5 0.66 0.8 2.3 0.3 58.85 3.0 1.5 1.2 6 1.2 0.99 1.9 3.6 0.8 84.56 4.6 1.5 0.2 10 0.4 0.37 0.5 1.0 0.3 54.17 1.4 0.8 0.9 8 0.7 0.97 1.1 3.7 0.3 50.0
Willcox Basin – trend 2 8 10.0 3.6 0.3 10 29.4 1.00 1.2 6.7 0.8 33.29 5.4 3.8 0.3 ND 5.5 0.89 0.6 4.7 0.5 33.3
10 20.3 8.8 0.3 ND 9.2 0.84 1.3 1.9 0.3 187.111 3.6 3.8 1.0 ND 2.1 0.89 1.1 4.7 0.8 44.412 12.2 26.2 0.2 ND 4.3 0.94 3.0 2.4 0.4 50.813 4.7 4.0 2.6 ND 3.1 0.94 4.1 3.9 1.8 47.114 21.7 4.9 0.5 ND 21.7 0.64 7.8 8.2 0.5 63.315 9.5 9.0 9.2 49 19.0 0.48 10.2 4.5 1.9 53.316 12.4 12.4 2.1 ND 13.1 0.70 1.3 6.4 1.7 47.1
USEPA primary maximum contaminant levels 100 10 50 6 20001‐ ND indicates element not detected (detection limit 6 µg L‐1).
73
Table 10: Isotope results and related trace element concentrations.
B Proportion δ11B Sr δ18O δ2H δ34SSO43H U
Well (µg L‐1) of B as B(OH)3 (‰) (µg L‐1) 87Sr/86Sr (‰) (‰) (‰) (TU) (µg L‐1) 224Ra (mBq L‐1) 226Ra (mBq L‐1) 228Ra (mBq L‐1)
Mountain blocks 1 8.7 0.99 9.4 35 0.750097 ‐11.4 ‐79 7.7 10 1.0 ±0.4 2.6 ±0.6 ‐0.2 ± 1.1 2 77 0.99 11.4 1030 0.732472 ‐9.8 ‐72 5.0 71 220 ±11 303 ±8 351 ± 49 3 106 0.99 5.3 912 0.745016 ‐11.7 ‐85 10.6 0.0 27 19.5 ±1.0 4.5 ±0.6 33.5 ± 4.5 Willcox Basin – trend 1 4 26 0.99 1.9 164 0.733361 ‐10.9 ‐78 6.5 4.4 20.4 ±1.5 5.7 ±0.7 18.9 ± 5.1 5 44 0.53 7.8 192 0.720641 ‐9.1 ‐64 5.1 0.0 1.7 5.6 ±0.8 1.6 ±0.6 7.9 ± 4.5 6 16 1.00 18.2 68 0.728224 ‐11.1 ‐79 7.1 0.3 5.8 ±1.2 1.5 ±0.5 3.4 ± 3.3 7 50 0.87 14.8 410 0.725143 ‐9.8 ‐71 6.3 0.0 14 1.3 ±0.4 1.6 ±0.4 1.8 ± 3.6 Willcox Basin – trend 2 8 76 0.46 ‐6.7 7.4 0.715026 ‐11.2 ‐83 10.2 2.7 0.2 ±0.1 0.4 ±0.6 1.5 ± 3.0 9 201 0.96 65 0.711471 ‐9.5 ‐68 9.5 0.1 11 2.6 ±0.3 0.6 ±0.5 2.2 ± 2.3 10 99 0.98 ‐0.5 190 0.711904 ‐11.2 ‐80 11.1 0.1 1.6 3.0 ±0.7 0.6 ±0.6 3.1 ± 2.3 11 127 0.97 3.9 268 0.710778 ‐9.3 ‐66 9.2 0.1 2.0 14.5 ±0.9 2.2 ±0.4 7.6 ± 3.6 12 98 0.94 0.9 333 0.713142 ‐11.8 ‐85 10.8 0.0 1.4 4.2 ±0.6 2.7 ±0.4 2.1 ± 2.3 13 215 0.95 0.3 671 0.710594 ‐8.6 ‐62 8.8 0.4 13 21.7 ±1.0 2.9 ±0.3 9.9 ± 3.9 14 126 0.99 4.0 218 0.711677 ‐11.0 ‐81 8.3 3.7 3.9 ±0.6 2.7 ±0.6 2.6 ± 2.5 15 390 1.00 2.3 641 0.714158 ‐10.0 ‐72 2.0 0.6 11 9.5 ±1.6 1.1 ±0.5 5.2 ± 4.3 16 103 0.99 2.7 2270 0.710122 ‐9.4 ‐69 13.4 0.2 11 10.4 ±1.1 48.6 ±1.1 2.4 ± 2.7
74
3.5.2 Isotopic tracing of sediment sources using Sr
As documented in previous studies, Sr concentrations and isotope ratios in groundwater
are primarily derived from interaction with the host aquifer rocks (Franklyn et al., 1991; McNutt,
1999; Harrington and Herczeg, 2003). Given that 87Sr/86Sr ratios in groundwater are not
fractionated by the reaction processes themselves, the isotopic ratios represent the
composition of the host aquifer rocks and in a case of multiple rock types, the Sr isotope ratio
could indicate the relative mixing combinations of the different sediment sources. The Sr
isotope ratios of some contributing aquifer rocks in the Willcox Basin have been previously
reported: (1) Proterozoic granite and schist with whole‐rock 87Sr/86Sr > 0.78 (Erickson, 1981); (2)
Tertiary volcanic rocks, inferred to be contaminated by melting of Paleozoic‐Mesozoic
sedimentary rocks and exhibiting whole‐rock 87Sr/86Sr of 0.708 to 0.716 (Bryan, 1987); and (3)
relatively minor exposures of Cambrian to Triassic sedimentary rocks including conglomerate,
sandstone, carbonates, and locally present gypsum with inferred seawater‐derived strontium of
87Sr/86Sr < 0.709 (McArthur et al., 2001). From these broad ranges of isotopic ratios, it appears
that the 87Sr/86Sr ratios of Paleozoic and Tertiary rocks could overlap; however, the Proterozoic
crystalline rocks impart a distinctive, radiogenic signature.
Willcox Basin samples were divided into two major trends based on Sr isotope ratios.
Trend 1 with radiogenic isotope composition (n=4, 87Sr/86Sr 0.720‐0.733) represents granitic‐
source material in locations ranging from near the mountain front to the axis of the basin. Trend
1 waters exhibit a Sr isotope signature clearly associated with the mountain block samples in
Proterozoic crystalline rocks (n=3, 87Sr/86Sr 0.733‐0.750). Trend 1 samples exhibit exceedingly
radiogenic 87Sr/86Sr ratios for sandstone aquifers up to ~10 km from the mountain front (Figure
75
11), derived from mixing of water from Proterozoic rock units (Frost and Toner, 2004; Uliana et
al., 2007) or by weathering of geologically old K‐feldspar in the aquifer sediments (Harrington
and Herczeg, 2003). Because of the large distance from the crystalline rocks at which 87Sr/86Sr >
0.720 is observed (Figure 11), the ratios must be the result of weathering of sediments within
the basin. A radiogenic 87Sr/86Sr signature is made possible by (1) the large proportion of K‐
feldspar in the mountain block granites (Erickson, 1981); (2) the age of the crystalline mountain
block rocks (section 3.2.1); and (3) the apparently original (unweathered) nature of feldspars
deposited in the basin. Thus, the radiogenic values of 87Sr/86Sr of groundwater are dominated by
K‐feldspar weathering rather than by plagioclase weathering which, with its low Rb‐Sr ratio and
high Sr concentration, contributes nonradiogenic Sr to groundwater regardless of the minerals’
age (Franklyn et al., 1991).
Trend 2 with relatively unradiogenic isotope composition (n=9, 87Sr/86Sr 0.710‐0.715)
records the influence of Cretaceous to Tertiary volcanic rocks and/or Paleozoic sedimentary
rocks. 87Sr/86Sr values of Sr derived from these rocks are expected to fall within a large range
(0.708‐0.716) as discussed above, but are substantially less radiogenic than Proterozoic granitic
rocks contributing to trend 1. These are inferred to contain sediments derived from Tertiary
silicic volcanic rocks, but this source cannot be clearly distinguished from Paleozoic‐source
material from the groundwater tracers in this study alone. In some instances, the proximity to
dominant rock types is ambiguous; for example, sample 8 is also close to Proterozoic
metasedimentary and metavolcanic rocks that may contribute radiogenic Sr. Trend 2 waters
include three samples (8, 10, and 12) located near the trace of the Apache Pass Fault Zone
(Figure 11), a major regional fault that predates the opening of the Willcox Basin (Drewes, 1986;
Erickson, 1988), and could contribute waters to explain the large differences in temperature,
76
major element chemistry, and 87Sr/86Sr between nearby wells. Towne and Freake (2001)
suggested that the Apache Pass Fault Zone may serve as a conduit for deep, high‐As waters to
the Willcox Basin. In such a case, mixing of a deep, high‐As water into shallow groundwater
could be a significant mechanism for generating high As concentrations in addition to local
water‐rock interaction (Coes et al., 2002; Crossey et al., 2009; Scanlon et al., 2009).
Overall, the weak relationship between strontium concentration and 87Sr/86Sr ratios (ρ =
0.23 between 1/Sr and 87Sr/86Sr; Figure 14) implies that Sr in the groundwater is not derived
from simple mixing of water or strontium sources. Rather, the large Sr isotope variations can be
explained by interaction with different rock types imparting distinctive 87Sr/86Sr ratios. The
observed trends associated with these 87Sr/86Sr ratios (Figure 14) also reflect the control of the
aquifer lithology on the water composition. However, the concentration of Sr in groundwater
can be controlled by secondary reactions that reduce divalent cation concentrations (section
3.5.4) but that would not affect the original 87Sr/86Sr ratios. At the center of the Willcox Basin,
not directly sampled by this study, lake bed carbonate preserves the 87Sr/86Sr ratio of strontium
in calcite‐precipitating waters. Thus, this value reports a mixture of contributing Sr sources to
the closed basin at the time of mineral precipitation, which could be useful for paleohydrologic
inference (e.g. Hart et al., 2004). A reported 87Sr/86Sr of 0.71049 (Naiman et al., 2000) of
Pleistocene Lake Cochise lake bed carbonate is consistent with the lowest 87Sr/86Sr value
reported in this study (0.71012) observed at well 16 near the Pleistocene shoreline (Figure 11).
Thus, the radiogenic Sr derived from crystalline rocks must be balanced by Sr with 87Sr/86Sr <
0.710 to produce values ~0.710 at the center of the basin. This may include Permian marine
gypsum (section 3.5.3) with an expected 87Sr/86Sr ratio of 0.707‐0.708 (McArthur et al., 2001) or
other Paleozoic sedimentary rocks with low 87Sr/86Sr ratios (Bryan, 1987). Gypsum or anhydrite
77
with their high solubility and low 87Sr/86Sr can sharply modify groundwater Sr isotope signatures
(Jacobson and Wasserburg, 2005). Overall, it seems that groundwater 87Sr/86Sr ratios are
acquired during calcite dissolution and/or silicate hydrolysis that contribute Sr relatively early
along groundwater flowpaths prior to down‐gradient geochemical modifications that consume
divalent cations while leaving the groundwater Sr isotopic ratio unmodified.
3.5.3 Sulfate sources
Sulfur isotopes of sulfate may provide information on oxyanion–forming trace element
behavior by tracing a potential source in the sulfide minerals within volcanic rocks (Smedley et
al., 2002) and by indicating possible sulfate‐reducing conditions (Krouse and Mayer, 1999),
which may strongly influence the mobility and speciation of trace elements including As (Kirk et
al., 2004; Xie et al., 2009). No evidence of sulfate reduction has been encountered in these
primarily oxic waters except for a faint sulfide odor during pumping of well 15, a shallow well
near the center of the basin. This odor dissipated prior to collecting the sample. Other
investigators have identified sulfide in the shallow black clay beneath the lake bed (Schreiber et
al., 1972), but the primary fate for sulfate at high concentrations is gypsum precipitation in the
evaporative environment of the Willcox Playa indicated by thick buried lake bed gypsum
deposits (Vine et al., 1979), which does not significantly fractionate δ34S of residual sulfate.
Because no widespread sulfate removal mechanism is capable of fractionating δ34S in the
Willcox Basin, δ34S is primarily sensitive to sulfate sources rather than removal mechanisms.
Thus, δ34S could have some potential use in tracing As and other trace elements if a sulfur
source derived from volcanic sulfide oxidation can be identified. In general, the volcanic‐derived
material (trend 2) exhibits higher δ34S than granite‐derived material and sulfate concentrations
78
and SO42‐/Cl‐ are generally higher in the volcanic‐derived portion of the aquifer (trend 2).
However, the mode of samples with δ34S near 10‰ and high SO42‐/Cl‐ (Figure 17b), consistent
with the findings of Eastoe et al., (2004), Gu (2005), and Adkins (2009) in nearby basins, exhibits
a marine gypsum isotope signature rather than that of oxidized volcanic sulfide, which is
characterized by a lower δ34S range in the region (2‐4‰; Burtel, 1989), or meteoric‐derived
sulfate (2‐8‰; Eastoe et al., 2004). Overall, the inferred volcanic‐source material in trend 2 does
not exhibit systematically lower δ34S that might be characteristic of oxidized volcanic sulfide.
Gypsum in lower Permian carbonates in the basin‐bounding mountain blocks,
specifically the Epitaph Dolomite, is a documented source of sulfate with δ34S ~12‰ (Burtel,
1989) influencing regional groundwater (Eastoe et al., 2004; Adkins, 2009), and this sulfate can
be broadly redistributed in alluvial basins (Eastoe et al., 2004; Gu, 2005). In the mountain blocks
bounding the Willcox Basin, the Epitaph Dolomite is present in the Dragoon Mountains west of
the basin (Figure 18; Drewes, 1987). Undifferentiated Paleozoic sedimentary rocks underlie the
basin‐fill sediments at the floor of the Willcox Basin (Drewes, 1985, 1987) and could also have
contributed gypsum during the basin’s history. The sulfur isotope results are consistent with a
marine gypsum source, probably from the western mountain block and recycled in the lake bed
deposits at the center of the basin (e.g. Eastoe et al., 2004; Gu, 2005). The extent of gypsum
influence can be seen in the location of gypsum‐like δ34S values mainly within the Pleistocene
extent of Lake Cochise rather than up‐gradient portions of the basin where atmospheric values
occur (Figure 17). The anomalously low δ34S value of 2.0‰ at well 15 is consistent with
oxidation of biogenic sulfide minerals (Krouse and Mayer, 1999) such as those observed in the
shallow lake bed (Schreiber et al., 1972). Overall, application of the sulfur isotope system implies
that dissolution of marine gypsum is a significant source of sulfate in groundwater from the
79
central and western Willcox Basin. In the northern and eastern portions of the basin where
gypsum is absent, δ34S is lower (6.5‐7.7‰; Figure 14) and thus more consistent with
atmospheric sources. Paradoxically, the system obtains oxyanion‐forming trace elements from
volcanic sources, yet δ34S does not reflect oxidation of volcanic sulfide. Apparently, the large
contributions of sulfur from gypsum overwhelm the contribution from oxidized volcanic sulfide,
both in the Willcox Basin and in alluvial basins throughout the region, where δ34S values are
consistently near Permian gypsum values (Gu, 2005).
3.5.4 Significance of the Ca/Na ratio
Another paradox evident from the sulfur isotope data is that samples with gypsum‐like
δ34S values exhibit the lowest Ca/Na ratios in the study area, and many exhibit Ca/SO42‐ ratios
below the expected ratio of 1 from gypsum dissolution (Figure 17d). This apparent inconsistency
must be resolved by down‐gradient groundwater evolution that systematically removes divalent
cations. Isotopic evidence for the reactive nature of divalent cations includes the poor
correlation between Sr concentration and 87Sr/86Sr (section 3.5.2), which implies that Sr
concentrations do not simply record mixing of two main Sr sources. Shallow, recently‐recharged,
and calcite‐undersaturated groundwaters acquire a relatively high Ca/Na ratio from calcite
dissolution in both granitic‐derived (e.g. White et al., 2005) and volcanic‐derived sediments (e.g.
Earman et al., 2005). Subsequent groundwater evolution with increasing residence time and/or
distance down‐gradient may contribute Ca through dissolution of broadly distributed gypsum in
the Pleistocene lake deposits (e.g. Eastoe et al., 2004; Gu, 2005), but apparently Ca is removed
effectively by chemical processes including calcite precipitation, silicate hydrolysis linked to
smectite precipitation, and/or cation exchange (Robertson, 1991; Bowser and Jones, 2002). In
80
the Willcox Basin, Ca/Na ratios vary widely, but distinctive low values <0.2 (Figure 15) occur near
the center of the basin (samples 8, 9, 10, 12, and 14; Figure 11). Although the lowest values of
Ca/Na are seen in trend 1 waters with inferred volcanic‐source aquifer sediment, the
relationships between Ca/Na and As concentration (ρ=0.67; Figure 15), δ34S and Ca/Na (Figure
17d), 87Sr/86Sr and Ca/Na (Figure 15), and 87Sr/86Sr and Sr concentration (Figure 14) collectively
suggest that Ca/Na at least partially records hydrogeochemical processes along groundwater
flowpaths and/or surface charge of sediments rather than merely fingerprinting the geologic
sources of solutes within the basin.
Figure 18: Map indicating δ 34S in relation to the possible gypsum source area in the Dragoon Mountains and the extent of Pleistocene lake deposits.
3.5.5 Trace element and isotope behavior: Implications for oxyanion‐forming elements
3.5.5.1 Arsenic
Broadly, the correlations between arsenic and fluoride, pH, and other oxyanion‐forming
trace metals in the Willcox Basin are consistent with existing data sets for the Willcox Basin
81
(Towne and Freark, 2001; USGS National Water Information System at
http://waterdata.usgs.gov/nwis, accessed 17 Feb 2010) and other published studies of oxic,
neutral to slightly basic groundwater systems (Robertson, 1989; Del Razo et al., 1993; Smedley
et al., 2002; Armienta and Segovia, 2008; Ortega‐Guerrero, 2009; Scanlon et al., 2009; Currell et
al., 2011). The correlation between Ca/Na and As (ρ= ‐0.67) is consistent with non‐conservative
behavior of As in which the presence of Ca rather than Na increases the surface charge of As
adsorption sites, most likely Fe oxides (Wilkie and Hering, 1996; Smith et al., 2002; Masue et al.,
2007; Jia and Demopoulos, 2005; Scanlon et al., 2009). Thus, the pH value at which As can be
adsorbed effectively increases as Ca/Na increases because of the more positive surface charge
of Fe oxides interacting with Ca, and thus high Ca/Na ratios would be associated with lower As
concentrations. These positively charged sites could effectively remove arsenate species such as
H2AsO4‐ and HAsO4
2‐.
Although the Ca/Na ratio of groundwater seems to be associated with down‐gradient
geochemical evolution (section 3.5.4) that also releases As, it is possible that low Ca/Na ratios
are associated with the high‐As volcanic rock types. For example, the even stronger correlation
between As and F‐ (ρ = 0.74) is probably unrelated to surface charge effects due to the relatively
conservative nature of fluoride in waters undersaturated with respect to fluorite (e.g. Levy et al.,
1999; Rango et al., 2009), and thus it seems that this correlation is due to co‐occurrence of As
and F in volcanic‐derived sediments. In sum, although water chemistry and Ca predominance
over Na could control As abundances in oxic groundwater, As concentrations in groundwater
seem to be controlled by rock type as indicated by the association of As with 87Sr/86Sr ratios
(Figure 14) and the relatively weak association between pH and As concentration. In other
82
words, many of the waters in this study with low As concentration (and high 87Sr/86Sr and Ca/Na
ratios) seem to be in aquifer material from non‐volcanic sources with inferred low As content.
3.5.5.2 Vanadium, chromium, selenium, molybdenum, and antimony
Although the occurrence of V, Cr, Se, Mo, and Sb is of less water quality significance
than As in the Willcox Basin, the behavior of these oxyanions can be treated as analogous to As
due to their similar chemical behavior. All of these elements occur at higher concentrations in
trend 2 (Figure 14), and all exhibit correlation with each other, although correlations are
significantly weaker for Cr and Se. Vanadium is consistently associated with volcanic‐source
aquifer material and with other oxyanion‐forming trace elements (Aiuppa et al., 2003; Vivona et
al., 2007) and in the study area, V concentrations are higher in trend 2 waters (3.6‐21.7 µg L‐1)
than in trend 1 waters or the mountain blocks (0.6‐4.6 µg L‐1). Cr exists primarily as soluble Cr(VI)
in oxic waters, and has been observed in basin‐fill aquifers sourced from mafic volcanic rocks
(Ball and Izbicki, 2004; Izbicki et al., 2008) or fine‐grained weathered volcanic material
(Robertson, 1975). Cr concentrations of up to 26.2 µg L‐1 were observed in this study, but most
concentrations are much lower; the felsic nature of the volcanic source rocks may explain the
relatively low Cr concentrations. In addition, Cr is not well correlated with Ca/Na (ρ= ‐0.22) or
other oxyanion‐forming trace metals (maximum ρ=0.50 with Mo). Selenium concentrations,
while generally low throughout the study area (0.2‐10.2 µg L‐1), are highest in waters of trend 2,
consistent with other oxyanion‐forming elements (Figure 14). Molybdenum may be associated
with rhyolite rather than mafic volcanic rocks (Arnórsson and Óskarsson, 2007) and is correlated
with As in high‐As alluvial basins of the southwestern USA (Robertson, 1989). The highest Mo
concentration (11.1 µg L‐1) was observed in a mountain block water sample, although otherwise,
83
concentrations above 5 µg L‐1 were only observed in trend 2 waters consistent with other
oxyanion‐forming elements (Figure 14). Antimony concentrations are low throughout the study
area (≤ 2 µg L‐1; Table 9) and except for one mountain block sample, the highest concentrations
were observed in trend 2 waters. However, Towne and Freark (2001) reported an isolated
occurrence of elevated Sb (9 µg L‐1) above drinking water standards. Thus, overall, the behavior
of V, Cr, Se, Mo, and Sb is similar to As in the Willcox Basin.
3.5.5.3 Boron
Boron isotope data in groundwater can reflect either (1) adsorption reactions in which
tetrahedral boron (B(OH)4‐) is preferentially adsorbed onto clay minerals or (2) boron sources
with different isotopic ratios (Vengosh et al., 1994, 1999). The low‐boron groundwater of trend
1 and the mountain blocks exhibits relatively higher δ11B values (up to 18‰) than waters with
higher B concentrations in trend 2, forming two distinct groups (open vs. filled symbols in Figure
16a). The associations between δ11B and 87Sr/86Sr, and arsenic concentration (Figure 16b‐c)
suggests that the boron isotopic composition in Willcox Basin groundwater is controlled by
boron sources rather than processes in the aquifer. It seems that rain‐derived boron with higher
δ11B (Vengosh and Spivack, 1999; Rose‐Koga et al., 2006) controls the composition of the less‐
evolved, low‐B groundwater in the granite or granitic‐derived basin‐fill materials. In contrast,
the association of low δ11B with low 87Sr/86Sr (samples with δ11B < 5‰ are characterized by
87Sr/86Sr < 0.715 and relatively low Ca/Na) ratios implies that rhyolite weathering affects boron
isotope ratios.
84
3.5.5.4 Radium and uranium
The very low Ra activities are consistent with previous studies of the Willcox Basin, in
which high Ra occurs only in the crystalline mountain blocks (Towne and Freark, 2001). Also, the
radionuclide activities reported here are consistent with previously published results from
another closed basin‐fill aquifer composed of granitic‐source sediment, the Carson River Basin in
Nevada (Thomas et al., 1993). In that investigation, median 226Ra activity and 226Ra/238U activity
ratio were similarly low as the Willcox Basin (4.4 mBq L‐1 and 0.026, respectively) and the basin
exhibits a similar pH range (~6.5‐9.5) as the Willcox Basin. The elevated 228Ra/226Ra ratios
throughout the basin (median 3.4), except for sample 16 at the basin center exhibiting
228Ra/226Ra < 0.1 (Table 10), are also consistent with the higher mobility of U relative to Th or Ra
because weathering‐derived 238U may be transported toward the center of the basin and thus
not contribute by decay into the precursors of 226Ra. Long‐term transport of U to the center of
the basin is also consistent with the U deposit hosted in lake bed deposits at the center of
Willcox Basin and other alluvial basins in Arizona (Duncan and Spencer, 1993), locally
contributing high U to groundwater in the center of the basin (Towne and Freark, 2001).
In contrast to the relatively soluble nature of U, several co‐occurring mechanisms may
remove Ra from water in the aquifer: (1) adsorption to Mn and Fe oxides (Ames et al., 1983c;
Herczeg et al., 1988) and clays (Ames et al., 1983a); (2) coprecipitation of radium into barite in
waters at or above barite saturation (Table 8; Grundl and Cape, 2006) ; and/or (3) processes that
remove divalent cations including Ra2+ from solution (section 3.5.4). The pH‐dependent
adsorption‐desorption behavior of Ra is essentially the opposite of As in oxic waters; Ra is
inefficiently adsorbed at pH < 5 due to the positive surface charge of Mn oxides (e.g. Szabo et
al., 2005). In the Willcox Basin, the high pH of the aquifer maintains negative surface charge on
85
Mn and Fe oxides, which makes Ra2+ adsorption highly effective; thus, pH‐sensitive oxyanion
desorption from Fe oxides (e.g. Robertson, 1989) corresponds with increasingly efficient Ra
adsorption. The non‐overlapping, pH‐dependent nature of Ra and As(V) desorption implies that
both elements should not co‐occur at high levels in oxic waters except in mineralized areas.
Overall, the occurrence of all Ra isotopes (224Ra, 228Ra, 226Ra) in groundwater, including
224Ra/228Ra broadly on the order of 1 in the basin‐fill aquifer (Table 10), is consistent with the
concept that radium isotopes’ thorium parents are present on aquifer surfaces and radium is
released by alpha recoil with rapid subsequent adsorption (Davidson and Dickson, 1986; Tricca
et al., 2001).
3.6 Conclusions
Willcox Basin groundwater samples have been divided into two groups, trend 1 and
trend 2, based primarily on 87Sr/86Sr variations. This provides a framework for assessing the
relative contributions of low‐As (granitic) sediment sources and high‐As (volcanic) sources in
imparting high arsenic concentrations to groundwater. Strontium isotope ratios record
weathering sources from aquifer rocks with different 87Sr/86Sr ratios, but do not trace Sr (and Ca)
removal processes. Sulfur isotopes record atmospheric sources, apparently modified in the
central and western Willcox Basin by marine gypsum dissolution. Boron isotopes primarily trace
atmospheric sources and weathering of volcanic‐source material during and after recharge.
Thus, the consistent patterns between Sr, B, and S isotopes are largely attributable to water‐
rock interaction (B, S, Sr) and to some degree also meteoric sources (B, S) in which all three
isotope signatures are acquired (Figure 19). Hydrogeochemical modification associated with
high‐As, bicarbonate‐dominated waters in down‐gradient portions of the aquifer is apparently
86
not recorded by Sr or S isotope variations but is recorded by the Ca/Na ratio. However, the
Ca/Na ratio is also somewhat associated with inferred sediment source in the Willcox Basin,
which complicates its use as a tracer of subsequent hydrochemical modification. Also, the
combined use of Ra isotopes and U concentration is consistent with inferred oxyanion
desorption behavior. The results of this study demonstrate that multiple tracers of water‐rock
interaction can provide information on the relative contributions of the aquifer’s lithologic
composition and hydrogeochemical processes to concentrations of oxyanion‐forming trace
elements in groundwater in groundwater from alluvial basins. 87Sr/86Sr in particular could be a
useful tracer for evaluating many alluvial basins in the southwestern United States (Figure 19),
where highly radiogenic Proterozoic rocks, Paleozoic sedimentary rocks, and Tertiary‐
Quaternary volcanic rocks with distinct 87Sr/86Sr ratios contribute sediment to basin‐fill aquifers.
These sediments impart elevated concentrations of As and other oxyanion‐forming trace
elements and affect water quality in one of the most water‐scarce areas of the United States.
87
Figure 19: Schematic cross‐section of a generalized alluvial basin (modified from Robertson, 1991) depicting influences on observed Sr, S, B, O, and H isotope ratios and the Ca/Na ratio.
88
4. Radium and strontium isotopes in the Jordan sandstone (Minnesota, USA): Carbonate control of a sandstone aquifer
4.1 Introduction
Naturally‐occurring radionuclides in the uranium and thorium decay series can
contribute excessive levels of radioactivity to some groundwater sources, including uranium,
radium, and radon. Sandstone aquifers in particular can yield radium (Ra) isotope activities that
exceed allowable drinking water levels, although the aquifer rocks themselves may exhibit
unexceptional concentrations of U and Th (e.g. Lively et al., 1992; Szabo et al., 1997; Vengosh et
al., 2009). This is generally interpreted to be the result of alpha recoil effects in which the
thorium parents of radium are concentrated on mineral surfaces and radium is released to
groundwater by alpha recoil (Table 11). This effect is balanced by rapid removal of radium,
resulting in a distribution of radium that is retained to the solid phase under most
hydrochemical conditions (Krishnaswami et al., 1982; Davidson and Dickson, 1986; Tricca et al.,
2001). Radium removal is influenced by pH, salinity, and redox effects. In a fresh, near‐neutral
bicarbonate‐dominated aquifer (TDS < 1000 mg L‐1) with little variation in cation and anion
composition, the effects of salinity and pH are negligible but redox conditions may significantly
affect radium adsorption due to the redox‐sensitivity of Mn and Fe oxides, which adsorb radium
if present (Szabo and Zapezca, 1987; Herczeg et al., 1988).
Radium activity in groundwater in excess of the U.S. Environmental Protection Agency
maximum contaminant level of 5 pCi L‐1 (185 mBq L‐1) combined 226Ra+228Ra activity is a well‐
documented regional groundwater problem in the Cambrian‐Ordovician sandstone and
carbonate aquifer system of the midwestern United States, a major water resource for the
89
region. Numerous potential sources and controls on radium have been suggested in previous
regional studies: shale content and exchange site distribution (Weaver and Bahr, 1991), salinity
(Horick and Steinhilber, 1978), uranium‐enriched carbonate units in multi‐aquifer wells (Jansen,
1994), control by barite solubility (Grundl and Cape, 2006), and perturbation of the groundwater
system by Pleistocene recharge events related to glacial melting (Gilkeson et al., 1983).
However, these studies have not addressed the balance between radium source and removal
mechanisms (alpha recoil vs. chemical release), nor has the use of short‐lived radium‐224 been
applied to these problems.
In this study, the problem of radium in an upper Cambrian sandstone aquifer is
addressed by detailed hydrogeochemical investigation of drinking water wells and related
aquifer solids. In contrast to previous studies that draw inferences from multi‐aquifer wells open
to several Cambrian‐Ordovician carbonate and sandstone aquifers and confining units, this
study focuses on a large number of single‐aquifer wells open specifically to the Jordan
sandstone. The Jordan aquifer is a calcite‐ and dolomite‐cemented quartz sandstone with
scattered tongue‐shaped bodies containing significant K‐feldspar (Runkel, 1994). Secondary
porosity is associated with carbonate dissolution, especially in shallow bedrock settings (Runkel
et al., 2003, 2006). The Jordan sandstone overlies the St. Lawrence confining unit, composed of
shaly and dolomitic rocks, and is overlain by the Ordovician Prairie du Chien group, which
contains crystalline and sandy dolomite (Table 12). In subcrop areas, the Jordan aquifer is
overlain by up to 100 m of Pleistocene glacial clay, sand, and/or gravel deposits that locally form
areas of favorable recharge or, especially in the west where Des Moines Lobe tills (Figure 20)
constitute an effective confining unit (Runkel et al., 2003). As a highly productive, low‐salinity
aquifer generally protected from surficial sources of contamination in the areas it is heavily
90
used, the Jordan sandstone is a major water resource for southeastern Minnesota and
northeastern Iowa, including the Minneapolis‐St. Paul metropolitan area (Horick and Steinhilber,
1978; Runkel et al., 2003). Large variation of radium activities, especially radium‐226, has been
documented in several Jordan aquifer groundwater wells (Lively et al., 1992) and some wells in
the Minneapolis‐St. Paul metropolitan area exceed EPA limits for radium activity. However, the
geographic extent of elevated radium levels in this aquifer was not well understood prior to this
study.
Table 11: Simplified 238U and 232Th decay series including nuclides discussed in this paper.
Nuclide Half‐life (yr) Mode of decay 238U decay series 238U 4.47 x 109 α 234Pa 7.65 x 10‐4 β 234Th 0.091 β 234U 246,000 α 230Th 75,400 α 226Ra 1600 α 232Th decay series 232Th 1.41 x 1010 α 228Ra 5.75 β 228Ac 7.02 x 10‐4 β 228Th 1.91 α 224Ra 4.11 x 10‐4 α
This objective of this research is to improve understanding of radium sources and
removal mechanisms in a confined, anoxic quartz sandstone aquifer at near‐neutral pH in which
Ra is known to occur both above and below the US drinking water standard of 185 mBq L‐1
(4 pCi L‐1; Lively et al., 1992). In addition to these broadly applicable questions, this study
addresses the vertical and geographic extent of elevated radium in the Jordan sandstone aquifer
91
for obtaining water that complies with EPA drinking water regulations without chemical
treatment. In addition, the detailed isotopic data obtained in this study (224Ra, 226Ra, 228Ra,
87Sr/86Sr) are combined with large data sets collected by the Minnesota Department Health,
Environmental Health Division. These data include stratigraphic and well depth information
(Table 13), radium‐226 and radium‐228 activities, and tritium activities. Using this broad
combination of groundwater data and analysis of the radionuclide content of selected aquifer
solids, radium sources and dominant removal mechanisms in the aquifer are investigated.
Table 12: Simplified list of hydrostratigraphic units discussed in this paper (modified from Mossler, 1990; Runkel et al., 2003).
Series Aquifer Name Confining
unit Rock types
Thickness (m)
Abbreviation
Drift Sand, gravel Quaternary
Till Silt, clay 0‐100 Q
Galena Limestone,
dolomite, shale 55‐65 GAL
Decorah‐Platteville‐Glenwood
Shale, carbonate
10‐35 DPG
St. Peter Sandstone 40‐50 STP
Ordovician
Prairie du Chien
Dolomite and sandy dolomite
45‐95 PDC
Jordan Sandstone 20‐40 JDN
St. Lawrence Siltstone and silty dolomite
15 STL
Franconia‐Ironton‐Galesville
Sandstone 60‐65 FIG
Eau Claire Shale 25‐35 ECR
Cambrian
Mt. Simon Sandstone 45‐85 MTS
92
Figure 20: Extent of Jordan‐Prairie du Chien aquifer system, confining units, and tritium activities in the Jordan aquifer and overlying Ordovician Prairie du Chien and St. Peter aquifers. Purple line represents extent of Prairie du Chien‐Jordan aquifer system derived from map data (Mossler and Book, 1984; Mossler and Bloomgren, 1990; Runkel, 1998; Mossler and Tipping, 2000; Mossler, 2001) or inferred from well
logs. Tritium data are from published reports (Alexander and Alexander, 1989; Tipping, 1992, 1994; Wall and Regan, 1994; Zhang and Kanivetsky, 1996; Campion, 1997, 2002; Berg and Bradt, 2003; Petersen, 2005) and unpublished data from Minnesota Department of Health, Minnesota Geological Survey, and Minnesota
Pollution Control Agency. Decorah confining unit extent is from Morey and Meints (2000); Des Moines till extent is from Hobbs and Goebel (1982).
93
4.2 Methods
4.2.1 Sampling of groundwater
Water samples were collected from drinking water wells after a period of pumping to
ensure that the well’s internal volume was fully purged. In most cases, the wells sampled for this
study are high‐capacity municipal drinking water wells (~4000 L min‐1) such that the well volume
is replaced within minutes. Field chemical measurements (pH, temperature, conductivity,
dissolved oxygen concentration) were determined using meters that were calibrated daily.
Construction details and related geographic information systems data for these wells were
obtained from the Minnesota Department of Health County Well Index (Table 13).
4.2.2 Major and trace element concentrations
Analysis of Ca, Mg, Na, Si, Sr, Ba, Fe, and Mn on filtered (<0.45 µm), acidifed water
samples was performed by direct current plasma spectrometry, and K was analyzed by flame
atomic absorption spectrometry. For these elements, analysis was calibrated by a multi‐element
standard solution prepared from plasma‐grade single element standards. A subset of filtered,
acidified water samples was analyzed for uranium concentrations by inductively coupled plasma
mass spectrometry (ICP‐MS), calibrated using plasma‐grade uranium standard in a multi‐
element matrix solution. Major anion concentrations (Cl‐, NO3‐, and SO4
2‐) were analyzed by ion
chromatrography. Bicarbonate concentrations were determined by titration of unfiltered water
samples to pH 4.5.
4.2.3 Strontium isotopes and tritium
Samples were prepared for strontium isotope analysis by placing filtered water samples
containing ~3 µg of Sr in Teflon vials and drying down on a hot plate in a laminar flow hood.
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After digestion in Optima HNO3, Sr was separated from samples using Eichrom SR‐B50‐S resin,
loaded onto degassed Re filaments in a Ta solution, and analyzed for 87Sr/86Sr using a
ThermoFisher TRITON thermal ionization mass spectrometer at Duke University. The long‐term
average 87Sr/86Sr value of NIST SRM 987 was 0.710245 ± 0.000010; 1σ, n=109). Tritium analysis
of selected unfiltered water samples was performed at the University of Waterloo
Environmental Isotope Laboratory by liquid scintillation counting subsequent to enrichment by
electrolysis. This method has precision of approximately 0.8 tritium units (TU).
4.2.4 Radium isotopes
Radium isotopes were preconcentrated from unfiltered, 40 L water samples collected
into polyethylene containers by slowly pumping samples through two sequential columns each
containing 10 g (dry weight) of manganese oxide‐coated acrylic fibers. Sample containers were
kept closed until processing and samples were processed the same day that they were collected.
Quantitative extraction of Ra from these mostly anoxic water samples was monitored by
analyzing 224Ra on the first and second columns and by comparison of selected samples with
whole‐process replicates analyzed at the Minnesota Department of Health laboratory for 226Ra
using EPA‐accepted standard methodology based on BaSO4 precipitation and gross alpha
counting.
Radium‐224 was analyzed using a delayed coincidence alpha counter. For samples with
total alpha counts > 20 cpm, the majority of the samples in the study, activities were
determined using total alpha counts corrected for 223Ra interference. For the samples of lower
count rate, 220Rn coincidence counts were used to determine 224Ra activity. In all cases, 224Ra
activities were decay‐corrected to the time of collection (Moore and Arnold, 1996; Garcia‐
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Solsona et al., 2008). This system was calibrated at an efficiency matching firmly hand‐squeezed
Mn oxide fibers. 226Ra was determined using a RAD7 radon‐in‐air monitor after sealing the Mn
fibers in a glass tube, which was evacuated and incubated for three weeks prior to analysis (Kim
et al., 2001). This was calibrated using a 226Ra solution (National Institute for Standards and
Technology) that had been transferred onto Mn oxide fibers. 228Ra was quantified by gamma
spectrometry using Canberra germanium detectors at Duke University using the 911 keV peak of
228Ac analyzed on the Mn fibers compressed into a disc geometry 5 mm thick and 65 mm in
diameter. The system was calibrated using three 232Th standards: DL‐1a uranium‐thorium ore
standard (Canadian Certified Reference Materials Project), GSP‐2 (USGS), and a 232Th solution on
compressed fibers, loaded in matching geometry. For 224Ra, error calculations follow the
methods of Garcia‐Solsona et al. (2008); for 226Ra and 228Ra, standard methods of counting
statistics and propagation of error were used (Eaton et al., 2005) including background count
rate subtraction and correction for the fraction of radium on the second column of fibers. For
radium isotope analyses, 2σ errors represent counting statistics only and do not include
potential uncertainties associated with other factors such as varied counting geometry (228Ra) or
moisture content variation (224Ra).
4.2.5 Analysis of aquifer solids
Cuttings from six wells were collected from the archive of the Minnesota Geological
Survey. According to well records, these wells were constructed using cable tool methods
without the addition of drilling mud. Selected samples were subjected to three analytical
procedures. (1) 0.5 g of dried, crushed well cuttings were shaken in 25 mL of 1 N HCl for 16
hours at room temperature (Hedley et al., 1982) to dissolve calcite and dolomite, plus other HCl‐
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soluble phases such as metal oxides coating grain surfaces. These samples were centrifuged and
the extracts analyzed for U and Th concentrations by ICP‐MS. (2) 34 mg of dry, crushed well
cuttings were dissolved in a HF‐HNO3 solution (modified from Stewart et al., 2002) and
underwent bulk chemical analysis for U and Th by ICP‐MS at Duke University, calibrated using
nine internationally accepted rock standards. (3) Dry well cuttings were crushed, weighed, and
sealed into polyethylene containers containing ~60 g of sediment. These were analyzed by
gamma spectrometry for 238U using the 63 keV peak of 234Th corrected for self‐adsorption effects
due to density differences (Gilmore, 2008), 226Ra (weighted average of the 352 keV peak of 214Pb
and the 609 keV peak of 214Bi), and 228Ra (weighted average of the 911 keV peak of 228Ac and the
583 keV peak of 208Tl) after an incubation period of at least three weeks. This analysis was
calibrated with DL‐1a loaded in matching geometry and corrected for the background count
rate.
4.2.6 Supplemental data sources
The maps presented in this study contain radium, barium, and tritium data from the
Minnesota Department of Health’s Minnesota Drinking Water Information System using
samples that could be identified as being collected from sample points prior to blending or
treatment. Radium‐226 and radium‐228 data are averages of all available results for each well,
generally with a reporting limit of 1 pCi L‐1 (37 mBq L‐1). Wells in which a majority of available
results were below the reporting limit were plotted as zero. Barium data represent an average
of all available results, giving preference to more recent ICP‐MS data where available. For
tritium data, the most recent available result was used if not collected specifically for this study
(Table 15).
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4.3 Results
4.3.1 Major element chemistry and redox conditions
Major element analysis of water samples indicates that Jordan aquifer groundwater
exhibits a consistent Ca‐Mg‐bicarbonate composition with approximately a 3:2 Ca:Mg molar
ratio. Concentrations of Na are uniformly lower than Ca and Mg; elevated concentrations of Na
and Cl occur in a few settings where the Jordan aquifer is near the surface and/or heavily
pumped (Table 14). In general, total dissolved solids increase to the south where the aquifer
becomes deeper and more confined, and to the west where the aquifer subcrops beneath Des
Moines Lobe tills. In most of the aquifer, conditions are essentially anoxic, indicated by low
dissolved oxygen concentrations (<1 mg L‐1) and elevated concentrations of dissolved Fe and Mn
(Table 14). Exceptions to this pattern of redox conditions are seen in subcropping or isolated
bedrock bodies of Jordan aquifer (e.g. the Hastings and Shakopee areas of Figure 21; the
Northeast Subcrop group; Table 14; and the eastern part of the Southern group such as well
642018; Table 14). The relationship between hydrogeologic setting and groundwater residence
time is illustrated along the cross‐section from east to west across the Jordan Aquifer (Figure
21). The few oxic portions of the Jordan aquifer generally exhibit measurable tritium, as do
several areas where the aquifer subcrops below sandy glacial sediments (e.g. the Shakopee and
Hastings areas of Figure 21) and/or is heavily pumped in the Twin Cities Metropolitan Area (e.g.
the Burnsville and Apple Valley areas of Figure 21). However, subcropping portions of the Jordan
aquifer in the western part of the study area exhibit tritium activities below detection due to the
presence of thick Des Moines Lobe tills in this area (Figure 20; e.g. the Chaska area of Figure 21),
which create a barrier to modern recharge.
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4.3.2 Strontium
Strontium concentrations (median 0.14 mg L‐1, range 0.05‐0.45 mg L‐1) and 87Sr/86Sr
ratios (median 0.71136, range 0.70861‐0.71347) exhibit large variations within the study area
(Table 15, Figure 22). The lowest concentrations of Sr were observed in areas where the aquifer
subcrops and contains lower levels of dissolved solids derived from carbonate material.
Strontium concentration is very well correlated with other alkaline earth metals Mg (Spearman
rank coefficient ρ=0.75), Ca (ρ=0.79), and HCO3‐ (ρ=0.82), but not with Ba (ρ=0.34) or 226Ra (ρ =
0.29). 87Sr/86Sr ratios exhibit large variations within the study area. The minimum ratio of
0.708610 is located in the far southeastern portion of the study area, and the maximum value of
0.713466 is in the Twin Cities metropolitan area (Table 13). In the Twin Cities metropolitan area,
Sr concentration is correlated with 87Sr/86Sr to imply a mixture of two strontium sources: one
source with low Sr concentration and radiogenic 87Sr/86Sr, a second source with high Sr
concentration and nonradiogenic 87Sr/86Sr (Figure 22).
4.3.3 Barium, radium, and uranium
4.3.3.1 Water samples
Barium concentrations exhibit large variations within the study area, approximately two
orders of magnitude (median 0.10 mg L‐1, range 0.001‐0.32 mg L‐1; Table 15), and exhibits
geographic patterns that are distinct from other alkaline earth metals Mg, Ca, and Sr. Barium
concentration is generally lowest in the subcrop areas of the aquifer (e.g. well 182655) and the
western area (e.g. well 414034), and generally higher in the Twin Cities metropolitan area, but
but not in the more deep, confined portions of the aquifer (e.g well 217550). Ba is not clearly
associated with concentrations of Mg (ρ=0.29), Ca (ρ =0.33), or Sr (ρ =0.33). On the other hand,
Ba concentration is much better correlated with its chemical analogue 226Ra (ρ = 0.59).
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Figure 21: Cross‐section along line A‐A’ (Figure 20) indicating hydrogeologic setting, open intervals (box lengths), tritium activities (in tritium units, indicated by box colors), 87Sr/86Sr (black numbers), and 226Ra activity in mBq L‐1 (red numbers). Heavy green line represents land surface elevation; city names are discussed in text. See Table 12 for hydrostratigraphic unit abbreviations.
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Figure 22: 87Sr/86Sr ratios in relation to 1/Sr, Sr/Ca, and Sr/Mg ratios and K concentration. Dashed line represents late Cambrian marine carbonates (McArthur et al., 2001).
Radium‐226 activities occur over two orders of magnitude (median 85 mBq L‐1, range
0.4‐418 mBq L‐1) while 224Ra and 228Ra activities are lower and less varied (median 47 mBq L‐1for
both). High activities of 226Ra occur in a broad, continuous region that occupies the center of the
Twin Cities Basin and a narrow north‐south belt south of the Twin Cities (Figure 23), whereas
228Ra activities are far more consistent throughout the study area, so that 228Ra rarely exceeds
226Ra. This results in 228Ra/226Ra activity ratios significantly below 1 in the high‐Ra area (median
0.5 overall; Figure 24), and these ratios are lowest in the waters highest in Ra (Figure 25).
Activities of 224Ra and its great‐grandparent 228Ra are similar in most samples, considering
analytical uncertainty, resulting in 224Ra/228Ra averaging approximately 1 (median 1.0; Figure 25).
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In relationship to US Environmental Protection Agency drinking water standards, 27% of the
wells sampled in the Jordan aquifer for this study would exceed the drinking water standard for
radium of 185 mBq L‐1 (5 pCi L‐1) of combined 226Ra+228Ra activity without blending or treatment.
In contrast to radium activities, uranium concentrations (median 0.4 µg L‐1, range <0.1‐6.3 µg L‐1)
are far below the EPA drinking water standard (30 µg L‐1) throughout the primarily anoxic wells
analyzed for U (Table 15).
4.3.3.2 Aquifer solids
Based on the gamma spectrometric analysis of bulk rock samples, 226Ra and 238U
activities of aquifer solids vary within approximately one order of magnitude (226Ra 1.6‐
15.2 Bq kg‐1; 238U up to 15.2 Bq kg‐1; Table 16). Solid phase 228Ra activities are consistently low
throughout the samples analyzed (median=2.7 Bq kg‐1), resulting in solid‐phase 228Ra/226Ra < 1
(median=0.5; Figure 26a). Equivalent isotope activities calculated from ICP‐MS elemental
analysis of bulk rock samples are broadly consistent with the gamma spectrometric analysis, also
indicating low Th/U (inferred median 228Ra/226Ra activity ratio 0.3 in the subset of samples
analyzed by ICP‐MS). Overall, the 1N HCl‐extractable U and Th concentrations indicate that
approximately half of the radionuclide activities present in the bulk rock were easily extractable
using a weak extraction method intended for calcite, dolomite, and other acid‐soluble minerals
(Figure 26b).
4.4 Discussion
4.4.1 Groundwater residence time and evolution of dissolved solids
Overall, within the study area, tritium is above the detection limit in the northeast
subcrop areas, the isolated groundwater cells containing the Shakopee and Hastings wells, and
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heavily pumped areas of the Minneapolis‐St. Paul metropolitan areas where downward leakage
has occurred (Figure 20; e.g. the Burnsville and Apple Valley areas of Figure 21). Due to the
surficial nature of nitrate inputs and the more oxic conditions in near‐surface sections of the
Jordan aquifer, elevated nitrate concentrations (Table 14) also indicate recent recharge and are
broadly consistent with the pattern of tritium results. These areas of comparatively rapid
recharge are in areas where bedrock confining units are absent and the aquifer underlies
permeable glacial sediments. Beneath the Des Moines Lobe tills in the western portion of the
study area and the Decorah shale confining unit in the southern portion of the study area,
Figure 23: Map showing radium‐226 activities in Jordan Aquifer groundwater and locations of wells from which cuttings were analyzed.
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Figure 24: Map showing 228Ra/226Ra activity ratios of Jordan aquifer groundwater.
tritium activities are generally below the reporting limit of 0.8 TU (Figure 20; e.g. Chaska in
Figure 21, which indicates that little or no modern recharge later than ~1950 has recharged the
aquifer in these areas. Beneath the Decorah shale confining unit in the southern portion of the
study area (Figure 20), groundwater is substantially slower‐circulating, with widespread
Pleistocene 14C model ages (Campion, 1997, 2002; Berg and Bradt, 2003). Overall, while the
lowest radium activities are observed in subcrop areas with tritium‐bearing water, such as the
northeastern portion of the study area, and the Shakopee and Hastings areas (Figure 21), there
is not a clear relationship between radium activity and groundwater residence time because the
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Figure 25: Radium isotope activities (224Ra, 226Ra, 228Ra) plotted in relationship to each other. Line marked “MCL” indicates the EPA maximum contaminant level.
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Figure 26: Plots showing Ra and U isotope activities in aquifer solids and relationship between radium‐226 in wells analyzed for groundwater and solids. Groundwater radium‐226 data for well 151582 are from Lively et al. (1992).
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oldest waters, under the Decorah confining unit, also exhibit relatively low radium. In sum,
elevated radium occurs in waters of intermediate or mixed groundwater residence time, which
implies that radium does not significantly accumulate as water travels down hydraulic gradients.
4.4.2 Strontium isotope insights on the carbonate phase
Throughout the Jordan aquifer, fresh groundwater exhibits a consistent Ca‐Mg‐HCO3
composition in which the molar Ca/Mg ratio is approximately 1.5 in all samples. This ratio is
seen consistently in the Jordan aquifer and overlying Ordovician sandstone and carbonate
aquifers in southeastern Minnesota (Tipping, 1994) in which both calcite and dolomite are
present (Thomas, 1991). This indicates a combination of calcite and dolomite equilibria in the
groundwater system where temperatures are approximately 10°C (median 10.1°C, range 8.0‐
12.9°C), a temperature at which calcite and dolomite solubilities are similar. In such situations
where both calcite and dolomite are present, the Ca/Mg ratio of groundwater depends on the
abundance of calcite and dolomite, temperature, and groundwater flow rate because both
congruent and incongruent dissolution of both minerals are possible. However, given that
Ca/Mg > 1, there must be net dissolution of both calcite and dolomite; congruent or incongruent
dissolution of dolomite alone would result in Ca/Mg ≤ 1 (Freeze and Cherry, 1979).
In the majority of the samples in this study, including the intensively sampled Twin Cities
metropolitian area, a systematic variation in 87Sr/86Sr was observed with Sr composition (Figure
22), which suggests that mixing relationships are seen between a low‐Sr source with high
87Sr/86Sr and a high‐Sr source with low 87Sr/86Sr (e.g. Bullen et al., 1996; McNutt, 1999). In the
Jordan aquifer, the 87Sr/86Sr ratio in groundwater (up to 0.71347; Figure 22) is significantly more
radiogenic than latest Cambrian (~490 Ma) seawater (~0.7090; McArthur et al., 2001), which
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indicates that the Sr contributing to the present‐day groundwater system is terrestrially derived
and that the carbonate is not of primary marine origin. In addition to Sr concentration, 87Sr/86Sr
is well correlated with the Sr/Ca and Sr/Mg ratios (Figure 22), which implies that 87Sr/86Sr is
influenced by carbonate mineral dissolution (Banner et al., 1994), and thus that the release of Sr
from calcite and/or dolomite contributes significantly to the observed 87Sr/86Sr. However, there
is a distinctive interruption of this carbonate‐linked pattern in the southern portion of the study
area where the relationship between 87Sr/86Sr and 1/Sr, Sr/Ca, and Sr/Mg (Figure 22) diverges
into two trends. One trend is the continuation of the correlation observed in the Twin Cities
area indicating contribution of nonradiogenic Sr with high Sr/Ca and Sr/Mg, although offset
toward lower values of 87Sr/86Sr. The divergent trend ends at low‐Sr, low‐Sr/Ca, low‐Mg/Ca
waters with nonradiogenic 87Sr/86Sr near 0.709 (Figure 22).
Several possible sources could contribute Sr of high 87Sr/86Sr to the aquifer. These
include: (1) near‐surface radiogenic Sr related to the Precambian crystalline rock components of
Quaternary glacial deposits (e.g. Franklyn et al., 1991; Vinson et al., 2010); (2) Sr derived from
clay minerals interacting with the aquifer, which would exhibit significantly higher 87Sr/86Sr than
primary carbonate (Stueber et al., 1987; Winter et al., 1995); (3) Sr from diagenetic dolomite,
associated with high‐temperature fluids and exhibiting 87Sr/86Sr 0.710‐0.711 in Ordovician quartz
sandstone in the Michigan Basin (Winter et al., 1995); and/or (4) depositional K‐feldspar that
occurs in discontinuous bodies in the Jordan sandstone (Odom, 1975; Thomas, 1991; Runkel,
1994), which with its high Rb/Sr ratio and Proterozoic age of crystallization would contribute
radiogenic Sr to waters of sufficiently long residence time to weather feldspar. In the divergent
trend in Figure 22, K is positively correlated with 87Sr/86Sr, consistent with K‐feldspar hydrolysis.
However, in the majority of the data, low K concentration is associated with high 87Sr/86Sr which
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is inconsistent with contributions from K‐feldspar hydrolysis. Because the most radiogenic
values of 87Sr/86Sr in the Jordan aquifer occur not in subcrop areas but in confined bedrock
conditions in the southern part of the Twin Cities metropolitan area (Figure 21) where Paleozoic
bedrock units, Superior lobe tills, and Des Moines Lobe tills are present in a complex pattern
(Mossler, 1990; Hobbs et al., 1990) it is implied that proximity to glacial deposits alone is not the
main control on 87Sr/86Sr. Overall, regardless of the geologic sources of radiogenic Sr, the
presence of carbonate minerals, the greater solubility of carbonate minerals over feldspar,
generally short groundwater residence time (tens to hundreds of years), and the relationships
between Sr, Ca, and Mg (Figure 22) imply that the immediate reactive source of Sr to most
Jordan aquifer waters is carbonate mineral dissolution, especially in the Twin Cities metropolitan
area.
4.4.3 The significance of 228Ra/226Ra
The ratio between 228Ra and 226Ra is primarily considered to indicate Ra sources rather
than Ra‐mobilizing mechanisms, and thus should be close to the 232Th/238U activity ratio of the
aquifer rocks (Dickson, 1990; Szabo et al., 1997). In sandstone aquifer systems, some previous
studies report median 228Ra/226Ra activity ratio near 1 (Dickson et al., 1987; Szabo et al., 1997,
2005) due to the tendency of U and Th to co‐occur at similar concentrations in major silicate
minerals, resulting in 232Th/238U activity ratio ≈1. On the other hand, carbonate aquifers are
distinguished by 228Ra/226Ra < 1 (Moise et al., 2000; Sturchio et al., 2001) due to the greater
affinity for U rather than Th in carbonate minerals. In the Jordan aquifer, 228Ra/226Ra ratios of
both rocks (Figure 26d) and groundwater (Figure 25) exhibit a continuum between low Ra
activity with 228Ra/226Ra ≈ 1 and high Ra activity with 228Ra/226Ra < 0.5. This similarity of
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228Ra/226Ra ratios, and the general association of higher‐226Ra aquifer solids with higher‐226Ra
groundwater (Figure 26c), imply that the U and Th content of the aquifer solids is a major
influence on the distribution of Ra in the studied groundwater. The low, carbonate‐like
228Ra/226Ra ratios in both groundwater and aquifer solids are consistent with an additional U‐
series radionuclide source to the aquifer beyond the detrital sandstone matrix alone. This
implies that a past fluid flow event could have locally added uranium to the aquifer. Additional
evidence that the low 228Ra/226Ra ratios reflect post‐depositional modification is seen in the
relationship between 228Ra/226Ra and 87Sr/86Sr (Figure 27). This correlation (ρ= ‐0.54) implies that
the lowest 228Ra/226Ra ratios, that is, the highest levels of radium in groundwater, are associated
with radiogenic strontium sources ‐ those that are most separated from a depositional
signature. This radiogenic Sr isotope signature, derived from carbonate (section 4.4.2), is
associated with carbonate‐like 228Ra/226Ra activity ratios <1. The relationship between 87Sr/86Sr
and 228Ra/226Ra also is consistent with the assumption that 228Ra/226Ra indicates Ra sources in
this system.
Figure 27: Relationship between 228Ra/226Ra and 87Sr/86Sr in Jordan aquifer groundwater.
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Low 228Ra/226Ra ratios are also seen in four major Cambrian‐Ordovician aquifers in the
region: the Galena carbonate aquifer, the St. Peter quartz sandstone, the Prairie du Chien
carbonate aquifer, and the Jordan sandstone (Figure 28). The quartz sandstone matrix in the St.
Peter and Jordan aquifers seems to have little apparent radium contribution light of the similar
levels of 226Ra and 228Ra in all four aquifers. In contrast to the four Cambrian‐Ordovician aquifers
documented in Figure 28, the underlying Cambrian Mt. Simon aquifer exhibits 228Ra/226Ra
generally ≥ 1 (Figure 28; Gilkeson et al., 1983; Lively et al., 1992). These data may imply that the
poor correlations between 226Ra and 228Ra in previous studies of groundwater radium in Illinois
(Gilkeson et al., 1983) and Wisconsin (Weaver and Bahr, 1991; Grundl and Cape, 2006) may be
partially due to sampling from multi‐aquifer wells in those studies. The data presented here
indicate that Ra sources are distinct between the upper and lower portions of the Cambrian‐
Ordovician aquifers. Similarly, Ra‐mobilizing processes may not be the same between the upper
and lower parts of the system.
4.4.4 Radium removal mechanisms
4.4.4.1 Adsorption and redox conditions
The most effective adsorption sites for radium include clay minerals, but also include
Mn oxides, which are redox‐sensitive. Therefore, reducing conditions could increase observed
levels of radium in groundwater (Szabo and Zapezca, 1987; Herczeg et al., 1988). Several wells
exhibit measurable dissolved oxygen, all of which are low in radium. However, the large majority
of the wells in this study are essentially anoxic and encompass a large range of radium activities.
Overall, the anoxic, Fe‐ and Mn‐oxide reducing conditions in the Jordan aquifer make radium
adsorption less efficient than oxic conditions, but the large differences observed in radium
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activities between wells of similar anoxic, high‐Fe conditions is probably due to differences in
radium supply from the aquifer solids.
Figure 28: Radium‐226 versus radium‐228 in regional Cambrian‐Ordovician aquifers. Data were obtained from this study, Lively et al. (1992), and the Minnesota Drinking Water
Information System database, with other sources noted in figures. Prairie du Chien‐Jordan multi‐aquifer wells are plotted with the Prairie du Chien aquifer; some Mt. Simon aquifer
wells include underlying Proterozoic sandstone aquifers.
4.4.4.2 Coprecipitation into barite
Precipitation of barite (BaSO4) in groundwater systems may be a significant radium
removal mechanism in waters at or above barite saturation (Martin and Akber, 1999; Grundl
and Cape, 2006) unless prevented by sulfate‐reducing conditions (Bolze et al., 1974; Martin et
al., 2003). Coprecipitation into barite has been invoked as a major control on radium in the
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Cambrian‐Ordovician aquifer system (Grundl and Cape, 2006) and has been directly observed
due to mixing in the boreholes of multi‐aquifer wells in Illinois (Gilkeson et al., 1983). Conditions
in the aquifer system seem to favor barite precipitation as a regionally significant mechanism:
(1) Low groundwater temperature (~10 °C) results in very low barite solubility. Calculations
using the PHREEQC code indicate that BaSO4 has solubility product of 10‐10.27 in dilute water at
10°C; (2) On average, the Jordan aquifer is saturated with respect to barite (median saturation
index ‐0.06, range ‐2.7 to 0.8); and (3) Sulfate reduction is not effective enough to compete with
barite precipitation. The high iron concentrations in anoxic waters (Table 14) imply that sulfate
reduction is ineffective due to the strong affinity of Fe2+ for sulfide, and one δ34S analysis of
sulfate is consistent with sulfide oxidation rather than sulfate reduction (‐7.8 ‰ for sample
716426; analysis provided by B. Mayer, University of Calgary).
It has been inferred that a constant Ra/Ba molar ratio of ~1.2 x 10‐8 indicates control of
aqueous radium by barite solubility in the Cambrian‐Ordovician regional aquifer (Grundl and
Cape, 2006). However, the constancy of this ratio would only hold significance if barite
dissolution is the major source of radium to water, in which case it would be related to the
Ra/Ba ratio of the barite solids. During radium removal by barite, the aqueous Ra/Ba ratio would
also be influenced by the rate of Ra supply from recoil and would not be expected to be
constant throughout an aquifer. In a system at chemical equilibrium with respect to barite, the
Ra/Ba ratio of barite and groundwater should reach a stable value governed by the partition
coefficient in which the Ra/Ba ratio of the barite is 1.8 times the Ra/Ba ratio of the water (Zhu,
2004a). This last scenario is most realistic for an aquifer with stable chemical conditions such as
the Jordan aquifer. Because the solid‐phase Ra source varies within the Jordan aquifer according
to variations in U and Th concentrations (Figure 26), the Ra/Ba ratio at chemical equilibrium
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would be expected to vary with the U(Ra) content of the aquifer solids and thus would vary
throughout the aquifer.
Ra and Ba concentrations are reasonably well correlated (ρ = 0.59) in the Jordan aquifer
(Figure 29), which implies that they experience similar removal mechanisms. The median value
of the Ra/Ba molar ratio in the Jordan aquifer is 1.26 x 10‐8, similar to the observed ratio of
Grundl and Cape (2006). Also, it is evident that a small number of wells (labeled in Figure 29 and
including well 200664) exhibit significantly higher Ra/Ba (Figure 29). In addition to the high
aqueous Ra/Ba in well 200664, analysis of aquifer solids from that well indicates anomalously
high radium in water relative to the measured 238U and 226Ra activities in the rock (Figure 26).
Thus, this outlying group represents wells where chemical removal of radium is less effective,
similar to the suggestions of Grundl and Cape (2006) for wells with anomalously high Ra/Ba.
4.4.4.3 Sources of barium to the Jordan aquifer
As with radium, geographic variability of barium concentrations indicates regional
variation of source and/or removal mechanisms. However, unlike radium, barium undergoes no
alpha recoil so that geochemical mechanisms alone, that is, barite equilibrium, would explain
groundwater barium concentrations (Underwood et al., 2009). A distinct zone of low Ra, high Ba
(low Ra/Ba ratio) can be seen at the western edge of the Twin Cities metropolitan area, and a
small area of high Ra, low Ba (high Ra/Ba ratio) can be seen at the eastern edge (Figure 30).
Potential geologic sources of Ba could include downward leakage from Quaternary glacial
deposits, feldspar weathering, or other unidentified sources. While the highest concentrations
of barium in Minnesota are seen in Pleistocene glacial aquifers (Minnesota Pollution Control
Agency, GWMAP data set), correlations were not observed in this study that would provide clear
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support for downward leakage of high‐Ba water from Pleistocene aquifers. For example,
groundwater concentrations of barium in the bedrock aquifers are not highest near the bedrock
surface, but are vertically consistent throughout the Jordan aquifer and overlying Ordovician
formations. Primarily, variation of barium within the bedrock aquifers is geographic rather than
vertical (Figure 30). Overall, these results suggest that the barium source is independent of
radium, but in the large area where elevated Ba and Ra coincide, Ra levels are retarded relative
to the area where Ba concentrations are much lower and the barite removal process is less
significant (e.g. well 200664; Figure 26).
Figure 29: Plot of Ba vs. Ra molar concentration. Wells with anomalously high Ba/Ba are labeled by their unique numbers.
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Figure 30: Inset maps across the central Twin Cities metropolitan area, where radium concentrations are higher, showing 226Ra (top), Ba concentration (middle), and Ra/Ba ratio (bottom). Data are from samples in the Jordan sandstone aquifer and overlying Ordovician Prairie du Chien and St. Peter aquifers. Dashed lines represent inferred breaks from low to high concentrations of Ra and Ba. Note that the geographic patterns of Ra and Ba do not
exactly coincide.
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4.4.5 Relationship to alpha recoil models of Ra isotopes in groundwater
Models of Ra mobility in groundwater describe a balance between Ra sources and
removal mechanisms for individual Ra isotopes of different half‐lives (Krishnaswami et al.,
1982). As used to develop retardation and distribution coefficients, these models do not
explicitly address salinity, redox, or coprecipitation phenomena but instead use a ratio of the
inferred Ra supply rate to the Ra removal rate to describe the overall effects of hydrochemical
processes on Ra mobility (Krishnaswami et al., 1982; Copenhaver et al., 1993; Sturchio et al.,
2001). The good agreement between absolute 226Ra activity of water and rocks and 228Ra/226Ra
ratios of both water and rocks (Figure 26d) implies that recoil from the aquifer rocks is a primary
control on observed Ra levels in the groundwater. Moreover, the solid‐phase 226Ra, 228Ra, and
228Ra/226Ra ≈ 1 in the low‐Ra areas of the aquifer (well 161435; Table 15) are consistent with the
U and Th content of quartz sands (Murray and Adams, 1958; Rogers and Richardson, 1964), and
the lower values of 228Ra/226Ra in the high‐Ra areas implies that U is the major radionuclide
source and is related to a secondary enrichment and not the detrital minerals. Thus, a model of
recoil could focus on the entire rock as the general source of Ra, or on the carbonate cement
phase itself as a specific Ra source.
The detailed radionucide analyses of aquifer solids from a range of Ra levels in
groundwaters allows calculation of the dimensionless distribution coefficient K for either
situation, which is an expression of the removal effectiveness of radium relative to its supply
rate. Petrographic analysis of Jordan sandstone samples (Thomas, 1991) indicates that, on
average, 69% of the rock is primary or secondary quartz, 21% is open porosity (φ), and 10% is
occupied by carbonate cement (χcement). K is defined by the expression:
117
in which Cd is the radionuclide activity of the solids (activity/mass), C is the radionuclide activity
of groundwater (activity/volume), ρs is the density of the solid (2.5 g cm‐3), and φ is porosity
(0.21; equation from Krishnaswami et al., 1982). Solving this using the data for the Jordan
aquifer in the six wells where both solids and groundwater were analyzed, K has a median value
of 550 and ranges from 150 to 920 in the Jordan aquifer when the whole rock is treated as the
source of radioactivity (Table 16). This is consistent with anoxic sedimentary aquifers
(Copenhaver et al., 1993) but lower than oxic aquifers by perhaps one order of magnitude (e.g.
Krishnaswami et al., 1982), possibly reflecting the relative inefficiency of Ra adsorption under
reducing conditions (Copenhaver et al., 1993). If the quartz sandstone matrix is neglected as
inert and the Ra‐enriched carbonate cement is treated as the primary radioactivity source (here
φcement = φ/(φ+ χcement)=0.68), and recalculating Cd based on the radionuclides being
concentrated in only 10% of the rock (χcement), median K would increase to 2200 (range 1200‐
4300), still within the previously reported range of K values for fresh, anoxic carbonate aquifers
in the midwestern USA (Sturchio et al., 2001). However, the second model would require
carbonate‐phase U and Ra activities up to 150 Bq kg‐1, which may be unrealistic. Overall, these
calculations illustrate that the inferred location and nature of the radionuclide source within the
aquifer source affects the extent of radium retardation contemplated by calculated K values by
about a factor of 5; however, both models produce reasonable estimates of adsorption behavior
for anoxic systems.
(2)
118
4.4.6 Timing of U addition to aquifer solids
Isotopic evidence including solid‐phase 228Ra/226Ra and groundwater 87Sr/86Sr suggests
that Ra enrichment in the Jordan aquifer and overlying Ordovician aquifers is related to an input
of U to the carbonate phase rather than being driven by the U and Th present in the sandstone
detrital material. The principles of U‐series disequilibrium allow evaluation of whether this
variation in solid‐phase uranium and groundwater 226Ra occurred due to Pleistocene glaciation,
which substantially altered the physical and hydrochemical makeup of the Paleozoic aquifers of
Minnesota by inducing significant regional increases in recharge, transporting reactive minerals
to the region (e.g. metal sulfides in till deposits), and depositing overlying drift aquifers and till
confining units (Siegel, 1989, 1991; Person et al., 2007; Grasby et al., 2010). Where the Jordan
aquifer today subcrops below glacial deposits and exhibits anoxic, tritium‐free characteristics,
for example, the aquifer may have been oxic and rapidly recharged in the presence of glacial
meltwater. This might be expected to significantly affect the uranium chemistry of the aquifer,
either by introducing U from the surface or by leaching U from the aquifer.
In the case of U deposition at an expected initial 234U/238U value near 1 (e.g. Osmond
and Cowart, 1992), ingrowth of 226Ra’s parent 230Th would be governed by the 75 kyr half‐life of
230Th. Thus, if U were added during Pleistocene recharge, one would expect solid‐phase
226Ra/238U < 1 (Levinson et al., 1984). On the other hand, if U were removed by oxidizing fluids
during the Pleistocene, one would expect solid‐phase 226Ra/238U > 1 because 226Ra would be in
transient equilibrium with the relatively insoluble 230Th and its decay would be governed by the
75 kyr half‐life of 230Th rather than the 1600 yr half‐life of 226Ra. Gamma spectrometric analysis
of crushed rock samples from the Jordan aquifer indicates 226Ra/238U ≈ 1 (Figure 26e) and thus
implies that 226Ra and 238U are at or near secular equilibrium in the aquifer. Although 234U and
119
230Th, the longest‐lived intermediates between 238U and 226Ra (Table 11), were not measured for
this study, it is apparent that 238U through 226Ra are at or near secular equilibrium: (1) 226Ra and
230Th are at or near secular equilibrium, governed by the 1600 yr half‐life of 226Ra, because the
Holocene (>5 half‐lives of 226Ra) has been more hydrologically stable than the Pleistocene; (2)
Although apparent equilibrium between 226Ra and 238U could be caused by initial 234U/238U >1,
perhaps introduced by glacial recharge that had interacted with freshly crushed rock or been
forced through shale confining units (Gilkeson and Cowart, 1982; Gilkeson et al., 1983; Osmond
and Cowart, 1992), it is not apparent that uranium with initial 234U/238U > 1 was introduced to
the Jordan aquifer in Minnesota. Three lines of evidence support initial 234U/238U ≈ 1 in Jordan
aquifer solids in Minnesota: (a) aqueous 234U/238U slightly above 1 (median 2.0) in anoxic Jordan
aquifer samples from Minnesota (Lively et al., 1992) implies, considering alpha recoil effects into
anoxic waters (Osmond and Cowart, 1992), that the solids themselves exhibit 234U/238U ≈ 1; (b)
Ordovician carbonate aquifer solids from Wisconsin exhibit 234U/238U ≈ 1 (Aument, 1989); and (c)
Pleistocene speleothems in Ordovician carbonate formations of southeastern Minnesota exhibit
234U/238U ≈ 1 (Lively, 1983). The interpretation presented here may not be broadly applicable to
the entire regional aquifer system; anomalous values of 234U/238U > 20 occur in deep, confined,
anoxic groundwater in Minnesota (Lively et al., 1992), Iowa (Siegel, 1989), and Illinois (Gilkeson
et al., 1983). These high ratios are difficult to explain by recoil‐adsorption behavior alone and
could signal initial solid‐phase 234U/238U > 1, perhaps during glacially‐induced recharge (Gilkeson
et al., 1983). In Minnesota, however, these ratios are limited to the underlying Mt. Simon‐
Hinckley aquifer and were not documented in limited sampling of the Jordan aquifer, primarily
in areas where the aquifer is not overlain by shale confining units (Lively et al., 1992). 226Ra‐238U
secular equilibrium implies a minimum U enrichment age of ~350 ka (5 half‐lives of 230Th;
120
Ivanovich et al., 1992) in the areas from which solids were sampled. In addition to a lack of
evidence consistent with a Pleistocene source of radionuclides, there is also no compelling
evidence suggesting that barium from Pleistocene till deposits influences radium removal
processes within the geographic area of this study. Uranium may be have been added to the
aquifer by older fluid flow events not identified during this study (e.g. Winter and Johnson,
1995) to explain the observed order‐of‐magnitude variations of 226Ra activity and 228Ra/226Ra .
4.5 Conclusions
The findings of this study demonstrate that the carbonate cement phase can be a
significant source of radium precursors in quartz sandstones, in addition to the U and Th present
in the detrital grains themselves, and that barite coprecipitation is a significant hydrochemical
control on Ra activities in groundwater in general agreement with the findings of Grundl and
Cape (2006). Also, in the anoxic conditions of the Jordan aquifer, redox effects on Ra adsorption
sites are probably important limiting factors, facilitating high levels of groundwater radium
relative to the solid phase by making radium adsorption less effective. The low, carbonate‐like
228Ra/226Ra values <1 in four Cambrian‐Ordovician sandstone and carbonate aquifers imply that
a regional U mobilization event has affected the aquifer system; the similarity of 228Ra/226Ra
ratios between solids and water samples implies that the solid‐phase radionuclides are the
source of waterborne radium. Apparent secular equilibrium between 226Ra and 238U on the
aquifer solids suggests that the addition of U to the carbonate cement occurred before ~350 ka;
extreme disequlibrium between 226Ra and 238U in water samples is controlled by redox
conditions rather than the isotopic composition of the aquifer solids due to the different
geochemical behavior of U‐ and Th‐series nuclides. Except for this minimum age estimate for the
121
U addition to the system, no specific depositional or diagenetic control was identified to explain
the variation in radionuclide levels within the study area. Likewise, the source of the similar but
not identical geographic distribution of barium, which affects radium mobility, was not
specifically identified. These results demonstrate that cements can contain regionally significant
enrichments of naturally‐occurring trace elements of water quality concern in the Paleozoic
aquifers of the midwestern United States. Similar results have been documented in the case of
high‐arsenic sulfide cement in Ordovician sandstone aquifers in Wisconsin (Schreiber et al.,
2000; Thornburg and Sahai, 2004).
122
Table 13: Names and well construction information for the wells sampled in this study. Six‐digit well identifiers are unique identifiers in the Minnesota Department of Health County
Well Index (http://www.health.state.mn.us/divs/eh/cwi/). Samples with dates not listed are unfiltered, unpreserved supplemental water samples analyzed for anions and 87Sr/86Sr only.
Well Name Sampling Date
Unique ID
Wellhead (m)
Depth to bedrock (m)
Top of open interval (m)
Bottom of open interval (m)
Hastings Group Hastings 4 207993 264 17 96 122Hastings 7 16 May 2008 509053 239 18 63 87Northeast Subcrop Group Wilder 5 05 May 2009 182655 302 36 52 61ISD 831 208548 317 49 62 105Lakamaga 12 05 May 2009 687180 291 52 53 67Southern Group St. Clair 1 13 Jan 2009 132254 301 73 94 122Medford 2 29 Jul 2009 147951 339 33 153 176Randolph 1 12 May 2008 161402 270 15 79 109Carleton College 2 12 May 2008 171005 283 1 86 108Sakatah Lake SP 1 16 Jul 2009 213638 330 66 128 155Kingdom Hall 06 Apr 2010 216520 301 20 114 130Mantorville 2 04 May 2010 217550 369 20 204 229Zumbrota 3 29 Jul 2009 228317 315 18 111 146Kenyon 1 05 Sep 2008 228324 351 14 177 201Cannon River Scout Camp 14 May 2008 228331 304 6 139 150Cannon Falls 3 05 Sep 2008 433273 258 2 91 120Morristown 2 15 May 2008 463537 319 35 124 149Rice Co. Recycling Center 06 Apr 2010 466198 334 4 136 147Waterville 3 29 Jul 2009 516045 317 37 159 185Faribault 6 29 Jul 2009 541549 299 16 100 123Medford 4 29 Jul 2009 541775 346 22 78 144Northfield 5 12 May 2008 559422 294 13 87 112Rochester 36 601336 319 24 121 146Zumbro Falls 4 29 Jul 2009 642018 299 6 73 103Gopher Hills Golf Course 14 May 2008 697770 288 4 109 116Rochester 38 698933 323 17 114 143Faribault 7 16 Jul 2009 716426 296 10 95 127Rochester 39 733087 314 15 111 140Dundas 2 12 May 2008 743063 337 6 130 158Hwy 76 Spring 02 Aug 2009 spring Twin Cities Metropolitan Group Eagan 5 16 May 2008 112234 291 71 124 153 Burnsville 13 04 Sep 2008 112235 265 47 99 124 Burnsville 10 13 May 2008 127259 262 26 91 118 Apple Valley 6 15 Jul 2009 127263 303 15 130 155 Bloomington 4 29 Apr 2010 133389 254 86 86 115 Apple Valley 9 02 Sep 2008 151561 310 24 131 157 MNPCA 12 May 2009 151590 235 58 94 116 Inver Grove Hts 5 14 May 2008 165640 270 54 109 138 Empire Twp 2 29 Jul 2009 171018 282 32 109 140
123
Well Name Sampling Date
Unique ID
Wellhead (m)
Depth to bedrock (m)
Top of open interval (m)
Bottom of open interval (m)
UM Entymology Fisheries 01 May 2009 200160 298 49 157 183 South St. Paul 6 15 May 2008 200664 273 15 122 148 South St. Paul 3 15 May 2008 200665 228 10 74 104 South St. Paul 1 16 May 2008 200674 259 12 98 123 Edina 13 30 Apr 2009 203613 289 32 131 151 Minnetonka 3 14 May 2009 204470 278 36 120 142 Plymouth 1 14 May 2009 204617 299 43 135 154 Orono 1 09 Jul 2009 205627 297 80 96 118 ISD 623 01 May 2009 205789 291 40 141 157 Apple Valley 3 02 Sep 2008 205812 317 37 146 179 Apple Valley 2 02 Sep 2008 205825 298 22 132 162 Burnsville 8 30 Mar 2010 206172 255 34 83 109 Edina 11 30 Apr 2009 206183 253 55 98 123 Richfield 1 30 Apr 2009 206353 255 61 105 134 Long Lake 2 206933 295 48 112 137 Lakeville 3 05 Sep 2008 207727 291 37 111 141 St. Louis Park 15 23 Jul 2009 215447 284 31 122 154 Eagan 4 16 May 2008 235373 263 89 103 127 Farmington 4 235586 283 24 114 146 Lake Auburn Camp Entry 21 Jul 2009 414034 302 87 89 102 Shakopee 7 04 Sep 2008 415975 231 2 44 67 Burnsville 16 13 May 2008 420992 311 45 142 173 Lakeville 8 13 May 2008 433296 335 76 160 188 Private well 21 Jul 2009 453886 300 64 94 102 Inver Grove Hts 7 06 May 2009 463527 289 49 128 157 Brooklyn Center 10 468118 259 31 76 98 Lexington Riverside 2 16 May 2008 519954 243 1 102 125 Greenfield‐Town Hall 1 21 Jul 2009 583308 304 72 73 74 Bloomington 5 15 Jul 2009 603079 257 87 94 124 Eagan 20 04 Sep 2008 626784 270 103 117 153 Shakopee 13 04 Sep 2008 674456 273 32 78 103 Chaska 9 04 May 2009 677176 276 61 69 102 Science Museum of Minn 12 May 2009 705735 217 14 74 92 Webster 3 (west) 16 Jul 2009 725107 336 70 105 131 Medina‐Hamel 6 14 May 2009 747666 309 70 95 116 Rosemount 15 15 Jul 2009 753663 294 44 119 149 Medina‐Hamel 7 09 Jul 2009 759809 311 57 95 124
124
Table 14: Major element chemistry and saturation index values (SI) of Jordan aquifer water samples.
Unique ID pH O2 T (°C) Ca Mg Na Fe Mn K Si Cl NO3‐ as NO3
‐ SO42‐ HCO3
‐ Calcite SI Dolomite SI Barite SI
Hastings Group 207993 13.36 41.3 509053 7.65 3.3 11.4 76.5 26.6 17.4 0.02 0.01 1.2 7.4 32.8 21.81 31.4 278 0.29 0.27 ‐0.20Northeast Subcrop Group 182655 7.43 1.3 9.8 48.8 18.2 4.2 0.02 0.00 1.3 8.1 3.8 1.51 7.5 223 ‐0.19 ‐0.70 ‐1.65208548 7.44 4.0 687180 7.98 0.8 11.8 29.6 11.3 4.6 1.23 0.03 1.8 5.3 1.0 0.41 0.1 168 0.07 ‐0.12 ‐2.73Southern Group 147951 7.02 0.2 9.6 117.4 33.5 27.7 0.80 0.05 3.4 5.3 0.8 1.77 60.7 520 0.04 ‐0.34 0.11161402 7.41 9.7 56.3 22.7 2.9 0.28 0.04 1.3 5.1 0.9 0.04 21.6 259 ‐0.11 ‐0.50 ‐0.43171005 7.33 10.1 92.5 28.4 6.1 1.09 0.06 2.3 5.3 1.3 0.06 17.8 424 0.21 0.02 ‐0.07213638 7.12 0.1 12.2 109.2 35.1 13.5 2.60 0.05 3.1 6.4 1.0 0.25 51.3 503 0.14 ‐0.04 0.06216520 7.06 0.2 9.5 87.7 30.9 10.5 1.05 0.05 3.6 6.9 1.5 0.06 10.3 466 ‐0.05 ‐0.45 0.02217550 7.19 0.4 10.4 68.2 23.9 3.4 0.29 0.01 1.6 4.0 0.8 0.07 27.8 322 ‐0.16 ‐0.64 ‐0.10228317 7.16 0.2 9.4 78.8 23.6 4.5 0.57 0.05 1.2 5.6 5.5 0.06 30.5 324 ‐0.14 ‐0.70 0.06228324 7.30 <0.1 11.3 68.9 21.2 4.0 0.23 0.03 3.2 3.3 4.1 1.29 58.0 270 ‐0.11 ‐0.59 ‐0.02228331 7.77 0.7 8.3 76.3 25.7 3.7 0.35 0.08 1.1 5.2 0.8 0.11 30.0 338 0.45 0.51 ‐0.14433273 7.56 0.0 10.5 49.4 18.5 2.9 0.24 0.03 1.2 4.9 4.5 1.29 25.7 227 ‐0.05 ‐0.40 ‐0.44463537 7.21 1.6 9.9 105.1 30.7 22.8 1.09 0.06 3.2 5.8 1.4 <0.01 43.3 475 0.17 ‐0.08 0.35466198 7.68 1.9 10.6 80.8 26.0 3.9 0.05 0.07 1.4 6.2 1.6 0.04 15.3 357 0.44 0.53 ‐0.01516045 7.26 0.2 9.8 92.5 28.5 21.2 0.61 0.04 2.8 6.1 1.1 0.08 46.0 422 0.12 ‐0.16 0.09541549 7.07 0.7 10.2 119.5 34.4 14.1 0.86 0.05 2.7 5.5 5.2 0.10 54.6 433 0.04 ‐0.33 0.55541775 7.12 0.3 10.8 111.1 31.6 20.5 1.39 0.05 4.2 5.2 1.0 0.10 24.1 521 0.15 ‐0.11 ‐0.06559422 7.43 9.5 82.5 31.0 3.6 0.07 0.08 1.3 5.9 5.8 0.72 35.6 360 0.18 0.05 ‐0.28601336 0.02 18.6 642018 7.42 9.3 11.0 63.0 24.9 2.7 0.03 <0.01 0.8 7.1 3.4 7.81 8.6 318 0.05 ‐0.16 ‐0.73697770 7.89 0.5 10.4 44.0 17.1 2.4 1.22 0.05 0.7 5.3 0.7 <0.01 16.2 207 0.19 0.10 ‐0.70698933 0.02 29.3 716426 7.15 0.2 10.9 96.6 29.0 17.7 0.79 0.05 2.7 5.6 13.6 0.15 48.0 405 0.10 ‐0.22 0.81733087 0.86 33.2 743063 7.24 9.5 71.5 24.5 2.4 0.92 0.04 1.2 5.9 1.0 0.07 21.5 347 ‐0.07 ‐0.49 ‐0.25
125
Unique ID pH O2 T (°C) Ca Mg Na Fe Mn K Si Cl NO3‐ as NO3
‐ SO42‐ HCO3
‐ Calcite SI Dolomite SI Barite SISpring 11.2 9.2 60.0 27.1 4.0 0.03 <0.01 0.7 6.8 7.6 18.69 6.4 301 Twin Cities Metropolitan Group 112234 7.45 0.4 10.0 73.8 28.5 3.5 0.08 0.10 1.1 6.3 5.9 0.06 22.1 359 0.17 0.04 ‐0.29112235 7.32 <0.1 10.1 80.3 29.1 5.3 0.31 0.06 2.0 5.8 15.0 1.30 21.6 369 0.08 ‐0.15 0.34127259 7.01 10.2 70.3 25.2 3.5 0.26 0.08 2.2 5.3 5.0 <0.01 18.7 346 ‐0.30 ‐0.91 0.03127263 7.33 0.4 10.1 73.4 24.2 6.7 0.41 0.07 1.8 6.3 13.4 0.05 20.7 323 0.01 ‐0.34 0.26133389 7.26 0.1 9.9 80.4 31.9 6.4 2.36 0.07 2.8 9.3 17.7 0.05 9.4 406 0.06 ‐0.17 0.08151561 7.33 <0.1 10.2 85.9 30.1 9.8 0.53 0.07 1.7 6.6 28.5 <0.01 37.5 345 0.09 ‐0.16 0.40151590 7.62 0.1 12.8 50.8 18.9 4.2 2.58 0.18 2.1 4.8 1.5 0.28 4.2 281 0.15 0.03 ‐1.12165640 7.56 0.8 10.0 81.7 29.3 6.6 0.06 0.25 1.5 6.3 17.9 1.68 21.1 354 0.31 0.30 0.42171018 7.52 0.1 9.2 62.1 21.7 3.2 0.53 0.06 1.7 5.9 0.9 0.09 8.9 304 0.10 ‐0.15 ‐0.23200160 7.36 0.1 10.5 51.3 18.3 5.6 0.29 0.05 1.9 6.7 2.1 0.45 1.9 282 ‐0.14 ‐0.59 ‐1.12200664 7.62 1.3 10.2 61.4 22.6 4.1 0.29 0.04 1.6 5.3 1.0 <0.01 8.9 323 0.24 0.16 ‐0.36200665 7.53 1.2 10.6 64.2 25.6 4.7 0.27 0.05 2.9 4.9 2.1 0.11 11.8 300 0.14 0.01 ‐0.23200674 7.61 0.4 10.5 72.2 25.9 5.6 0.35 0.06 2.1 5.9 6.8 0.07 16.0 335 0.30 0.29 ‐0.06203613 7.17 0.2 10.1 82.8 30.7 15.5 0.71 0.05 2.1 7.7 37.6 0.27 19.2 361 ‐0.07 ‐0.45 0.14204470 7.30 0.2 9.7 68.3 27.2 12.5 0.29 0.18 2.4 8.0 2.5 0.32 6.8 393 0.03 ‐0.24 ‐0.42204617 7.39 0.1 9.7 65.6 21.9 4.1 0.49 0.13 1.6 8.9 1.8 0.33 3.2 352 0.06 ‐0.23 ‐0.62205627 7.45 0.2 9.7 88.0 41.4 23.2 0.54 0.20 4.2 7.7 1.4 0.05 21.9 509 0.36 0.50 ‐0.03205789 7.61 10.6 45.1 16.0 5.2 1.29 0.05 2.0 5.8 1.7 0.39 1.5 254 0.02 ‐0.27 ‐1.44205812 7.44 <0.1 10.1 66.1 25.9 3.3 0.28 0.07 1.1 5.8 5.7 <0.01 23.3 320 0.07 ‐0.14 0.04205825 7.42 <0.1 11.0 71.3 23.7 4.9 0.33 0.05 1.7 6.5 3.9 1.36 5.8 344 0.13 ‐0.07 ‐0.19206172 7.14 0.1 9.9 77.5 30.3 5.6 0.51 0.07 2.2 6.4 16.1 0.16 23.5 373 ‐0.11 ‐0.51 0.48206183 7.19 0.1 10.2 83.4 32.0 15.3 0.59 0.06 2.0 7.6 51.4 0.76 26.9 358 ‐0.05 ‐0.39 0.35206353 7.24 0.1 10.4 80.2 30.0 15.9 0.41 0.08 2.3 7.9 54.2 0.28 27.1 344 ‐0.03 ‐0.36 0.17207727 7.45 <0.1 10.9 68.4 23.2 3.2 0.58 0.08 1.3 6.7 4.8 1.33 11.6 313 0.11 ‐0.12 ‐0.01215447 7.25 0.3 10.0 93.7 31.7 11.2 0.81 0.10 2.0 7.1 25.1 0.08 33.5 387 0.08 ‐0.18 0.43235373 7.71 0.9 10.0 78.6 27.8 4.4 0.28 0.25 1.5 8.5 6.6 0.10 18.6 375 0.47 0.61 0.23235586 0.02 20.6 414034 7.30 0.3 9.7 83.2 39.3 11.7 0.45 0.03 3.6 5.5 1.3 0.18 14.4 478 0.17 0.12 ‐2.19415975 7.30 3.8 11.3 91.1 34.7 45.0 0.01 0.01 4.3 8.2 92.6 22.08 13.6 333 0.07 ‐0.13 ‐0.21420992 7.24 9.9 79.9 29.9 5.8 0.36 0.15 1.6 6.3 19.3 <0.01 20.1 356 ‐0.02 ‐0.34 0.30433296 7.53 9.8 73.8 24.0 5.6 0.36 0.07 1.6 5.4 1.6 0.36 7.9 369 0.27 0.16 ‐0.15
126
Unique ID pH O2 T (°C) Ca Mg Na Fe Mn K Si Cl NO3‐ as NO3
‐ SO42‐ HCO3
‐ Calcite SI Dolomite SI Barite SI453886 7.17 0.3 12.9 91.4 36.4 14.6 0.20 1.71 4.9 6.1 1.5 0.13 12.7 492 0.14 0.05 ‐0.05463527 7.23 <0.1 10.0 69.8 24.7 4.0 0.11 0.25 1.8 7.9 1.7 0.03 11.0 338 ‐0.09 ‐0.51 0.11519954 7.69 2.3 12.0 65.3 24.6 4.3 0.28 0.10 1.6 6.0 5.0 0.05 19.1 325 0.35 0.44 ‐0.04583308 7.33 0.2 12.9 93.2 34.8 15.4 0.72 0.07 2.8 8.6 1.4 0.02 10.8 488 0.30 0.35 ‐0.23603079 7.42 0.3 10.7 76.6 29.6 5.1 0.76 0.08 2.2 8.4 14.6 1.51 20.0 338 0.14 0.00 0.43626784 7.50 <0.1 10.1 67.3 24.1 3.7 0.17 0.08 1.4 8.0 4.1 1.28 11.6 327 0.15 ‐0.02 ‐0.05674456 7.23 1.4 10.0 89.2 36.4 9.4 0.01 0.07 4.0 8.3 16.3 4.75 17.7 428 0.09 ‐0.10 0.13677176 6.95 0.1 9.2 81.1 33.9 20.8 0.11 0.71 4.1 7.6 1.5 0.27 20.1 481 ‐0.20 ‐0.67 ‐0.23705735 7.37 0.1 10.7 57.0 21.1 4.8 0.32 0.11 2.0 6.8 4.7 0.34 17.8 299 ‐0.07 ‐0.44 0.02725107 7.26 0.1 9.9 100.5 41.1 25.8 2.45 0.15 2.0 6.9 1.1 0.10 32.1 551 0.25 0.23 0.10747666 7.30 0.2 9.0 81.4 30.0 5.9 0.48 0.31 2.6 10.5 1.3 0.27 5.0 452 0.14 ‐0.06 ‐0.58753663 7.30 0.4 9.8 73.8 26.9 3.8 0.09 0.08 1.8 5.5 5.6 4.02 24.4 307 ‐0.05 ‐0.41 ‐0.12759809 7.24 0.1 10.3 90.4 31.6 6.9 1.21 0.27 1.9 8.1 1.1 0.02 1.4 455 0.14 ‐0.05 ‐1.24
127
Table 15: Trace metal concentrations, Sr isotope ratios, and tritium and radium activities of Jordan aquifer groundwater. Italized tritium activites are from samples collected by the Minnesota Department of Health on different dates than indicated in Table 13.
Unique ID Ba (mg L‐1) Sr (mg L‐1) 87Sr/86Sr Tritium (TU) 224Ra (mBq L‐1) 226Ra (mBq L‐1) 228Ra (mBq L‐1) U (µg L‐1)
Hastings Group 207993 0.710870 7.2 509053 0.05 0.10 0.710985 10.2 7.3 ± 1.2 8.2 ± 2.7 3.7 ± 8.9 Northeast Subcrop Group 182655 0.00 0.06 0.712407 6.3 0.6 ± 0.3 0.5 ± 0.8 <1 0.7208548 0.712073 687180 0.04 0.07 0.710953 2.9 4.9 ± 0.7 9.0 ± 1.3 3.9 ± 2.5 Southern Group 132254 0.709965 147951 0.06 0.36 0.709354 <0.8 112 ± 7 92.8 ± 4.4 84.8 ± 18.4 <0.1161402 0.03 0.09 0.710276 <0.8 23.3 ± 1.8 59.7 ± 3.8 33.5 ± 7.3 171005 0.11 0.21 0.710026 <0.8 54.1 ± 2.9 94.5 ± 4.0 56.5 ± 10.5 0.3213638 0.06 0.37 0.709732 <0.8 69.6 ± 3.1 61.0 ± 3.1 86.4 ± 15.8 0.8216520 0.22 0.33 0.709933 <0.8 147 ± 6 214 ± 7 127 ± 20 217550 0.06 0.12 0.710631 <0.8 54.7 ± 3.8 31.1 ± 1.9 47.4 ± 12.3 228317 0.08 0.11 0.710465 2.3 43.5 ± 2.5 46.2 ± 2.6 43.8 ± 10.0 228324 0.04 0.25 0.709802 <0.8 83.2 ± 5.6 79.1 ± 3.9 63.0 ± 11.8 228331 0.05 0.13 0.710362 <0.8 40.8 ± 2.5 48.1 ± 3.6 46.9 ± 9.2 433273 0.03 0.09 0.710234 <0.8 30.0 ± 3.0 63.3 ± 3.9 21.6 ± 5.2 463537 0.13 0.26 0.709761 <0.8 56.4 ± 3.8 57.4 ± 3.8 57.7 ± 13.8 466198 0.06 0.14 <0.8 59.0 ± 3.9 75.1 ± 3.4 70.6 ± 19.4 0.4516045 0.06 0.25 0.710258 <0.8 88.5 ± 3.8 64.2 ± 2.9 108 ± 16 0.4541549 0.17 0.27 0.709793 1.2 77.5 ± 3.7 64.4 ± 3.4 74.0 ± 12.0 541775 0.09 0.45 <0.8 34.6 ± 2.5 140 ± 6 47.7 ± 13.5 559422 0.03 0.10 0.710185 3.3 38.9 ± 2.6 22.1 ± 4.1 26.2 ± 7.0 1.4601336 0.709943 1.2 642018 0.04 0.06 0.709451 3.3 17.8 ± 1.7 4.7 ± 1.0 16.8 ± 5.5 697770 0.02 0.05 0.709595 <0.8 13.4 ± 1.3 37.4 ± 3.1 11.3 ± 3.4
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Unique ID Ba (mg L‐1) Sr (mg L‐1) 87Sr/86Sr Tritium (TU) 224Ra (mBq L‐1) 226Ra (mBq L‐1) 228Ra (mBq L‐1) U (µg L‐1)
698933 0.710147 1.9 716426 0.32 0.26 <0.8 89.7 ± 3.9 145 ± 7 91.8 ± 18.9 733087 0.709113 3.7 743063 0.05 0.09 0.710660 <0.8 18.8 ± 1.4 21.2 ± 3.0 34.5 ± 7.7 <0.1 0.04 0.05 0.708610 10.6 ± 0.9 3.8 ± 0.5 8.5 ± 4.9 Twin Cities Metropolitan Group 112234 0.05 0.09 0.713091 47.0 ± 2.4 87.3 ± 4.3 47.1 ± 9.5 0.4112235 0.22 0.14 0.712394 5.0 47.8 ± 3.0 119 ± 4 54.6 ± 13.4 127259 0.12 0.14 0.711415 2.7 38.2 ± 3.1 84.6 ± 4.3 39.5 ± 8.8 0.5127263 0.18 0.14 0.711684 2.2 43.1 ± 3.2 92.2 ± 4.6 42.5 ± 10.9 0.2133389 0.27 0.23 0.711352 55.9 ± 3.7 94.6 ± 3.9 43.4 ± 13.2 151561 0.15 0.13 0.712413 6.1 61.2 ± 5.7 137 ± 6 47.7 ± 9.3 151590 0.04 0.10 0.711970 <0.8 6.3 ± 1.3 30.8 ± 2.1 11.9 ± 2.5 165640 0.27 0.12 0.712516 5.2 61.8 ± 8.0 233 ± 8 60.6 ± 21.2 171018 0.12 0.09 0.711885 <0.8 26.9 ± 2.4 73.7 ± 3.8 29.0 ± 9.7 200160 0.07 0.14 0.712201 <0.8 34.7 ± 2.1 82.3 ± 3.8 36.8 ± 5.3 200664 0.09 0.12 0.712034 <0.8 56.4 ± 9.9 418 ± 12 58.3 ± 37.6 0.2200665 0.10 0.18 0.710503 <0.8 23.1 ± 2.1 86.9 ± 5.9 39.7 ± 9.3 200674 0.11 0.14 0.711780 32.1 ± 2.1 61.7 ± 3.8 39.8 ± 7.0 0.2203613 0.16 0.19 0.712009 2.4 78.4 ± 5.1 156 ± 6 82.7 ± 9.2 204470 0.11 0.25 0.711185 <0.8 57.4 ± 3.8 149 ± 5 41.7 ± 5.7 1.5204617 0.14 0.12 0.712333 <0.8 20.5 ± 1.7 107 ± 4 19.4 ± 2.9 205627 0.10 0.29 0.711360 <0.8 145 ± 8 213 ± 7 130 ± 27 205789 0.04 0.10 0.712467 <0.8 20.4 ± 0.9 49.6 ± 2.3 16.7 ± 3.1 <0.1205812 0.09 0.09 0.713123 2.8 29.8 ± 2.1 51.8 ± 3.9 30.7 ± 5.7 205825 0.23 0.14 0.711540 <0.8 69.5 ± 5.7 148 ± 6 61.8 ± 11.4 206172 0.27 0.16 0.712159 40.6 ± 2.1 97.6 ± 4.1 44.1 ± 7.6 206183 0.19 0.19 0.712253 6.1 81.7 ± 4.1 125 ± 5 68.6 ± 8.9 206353 0.12 0.20 0.712093 5.8 100 ± 6 128 ± 5 75.9 ± 7.4 206933 0.710852 207727 0.17 0.10 0.712664 1.6 61.2 ± 5.7 129 ± 6 55.4 ± 12.4 0.5215447 0.19 0.18 0.712199 60.4 ± 4.9 167 ± 5 72.4 ± 17.4 0.2
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Unique ID Ba (mg L‐1) Sr (mg L‐1) 87Sr/86Sr Tritium (TU) 224Ra (mBq L‐1) 226Ra (mBq L‐1) 228Ra (mBq L‐1) U (µg L‐1)
235373 0.19 0.12 0.712393 3.7 40.2 ± 2.8 110 ± 5 67.3 ± 15.7 0.5235586 0.710743 2.9 414034 0.001 0.28 0.711670 2.5 ± 0.6 1.9 ± 0.7 2.6 ± 2.9 415975 0.11 0.18 0.711653 6.0 33.8 ± 2.0 10.0 ± 2.9 10.9 ± 2.4 420992 0.21 0.13 0.712833 35.6 ± 2.2 63.5 ± 3.9 37.8 ± 9.1 433296 0.17 0.14 0.711519 1.1 46.6 ± 2.7 117 ± 5 56.0 ± 14.7 453886 0.19 0.41 0.710350 71.7 ± 4.9 21.9 ± 1.8 51.8 ± 11.3 463527 0.23 0.11 0.712545 <0.8 58.1 ± 4.8 250 ± 7 69.7 ± 9.2 468118 0.712977 519954 0.10 0.12 0.711983 1.6 52.1 ± 3.8 298 ± 7 48.8 ± 13.7 0.2583308 0.15 0.26 0.711204 <0.8 55.4 ± 4.8 83.4 ± 3.5 40.5 ± 13.3 603079 0.29 0.17 0.711923 3.8 36.2 ± 2.6 113 ± 5 37.3 ± 9.2 0.1626784 0.15 0.11 0.712135 <0.8 45.9 ± 3.4 105 ± 5 37.2 ± 6.6 674456 0.17 0.29 0.710698 1.9 18.6 ± 1.5 6.6 ± 3.6 9.3 ± 2.3 677176 0.06 0.40 0.710786 <0.8 85.1 ± 3.4 30.4 ± 2.1 47.4 ± 7.0 705735 0.11 0.12 0.712582 1.6 102 ± 6 136 ± 5 79.7 ± 9.9 0.5725107 0.10 0.26 0.711361 <0.8 72.6 ± 4.4 116 ± 5 70.5 ± 15.4 0.9747666 0.10 0.25 0.711160 <0.8 129 ± 6 188 ± 7 127 ± 26 753663 0.06 0.09 0.713466 2.7 41.8 ± 3.6 63.5 ± 2.7 31.5 ± 9.2 6.3759809 0.09 0.23 0.711239 <0.8 83.8 ± 5.9 211 ± 7 84.4 ± 20.0 1.3
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Table 16: Results of analysis of aquifer solids from Jordan aquifer wells.
1N‐HCl extractable (mg kg‐1)
Equivalent activities (Bq kg‐1)
Bulk ICP‐MS (mg kg‐1)
Equivalent activities (Bq kg‐1) Bulk gamma analysis (Bq kg‐1)
UniqueID
Depth (m) U Th 238U 232Th232Th/238U/
U Th 238U 232Th232Th/238U
226Ra 228Ra 228Ra/226Ra 238U K
161435 70 ‐ 78 0.3 0.6 0.0 0.0 0.6 4.5 ± 1.1 3.8 ± 3.0 0.8 ± 2.7 (1) 78 ‐ 83 0.1 0.2 0.0 0.0 0.6 2.4 ± 0.7 2.6 ± 1.5 1.1 ± 1.8 83 ‐ 92 0.1 0.2 0.0 0.0 0.7 3.1 ± 1.0 3.9 ± 2.0 1.3 ± 2.9 92 ‐ 98 0.2 0.7 0.0 0.0 1.4 1.9 ± 1.1 3.4 ± 2.6 1.8 ± 5.1
920
151582 87 ‐ 96 0.6 0.5 0.0 0.0 0.3 1.1 0.7 13.4 2.7 0.2 15.2 ± 0.8 3.9 ± 1.6 0.3 ± 0.4 15.2 ± 3.3 96 ‐ 106 0.7 0.5 0.0 0.0 0.3 0.9 0.9 11.6 3.5 0.3 6.2 ± 1.0 3.1 ± 2.4 0.5 ± 1.3 106 ‐ 115 0.2 0.2 0.0 0.0 0.3 0.4 0.4 4.7 1.5 0.3 5.3 ± 0.5 2.4 ± 0.8 0.4 ± 0.4 8.1 ± 2.5 115 ‐ 119 0.2 0.9 0.0 0.0 1.3 0.5 1.1 6.8 4.4 0.6 6.7 ± 0.7 3.5 ± 1.3 0.5 ± 0.8 4.3 ± 3.5
350
133389 81 ‐ 90 1.6 ± 0.7 2.8 ± 1.7 1.8 ± 3.3 92 ‐ 96 0.1 0.2 0.0 0.0 0.5 3.2 ± 0.9 1.7 ± 1.9 0.5 ± 1.2 107 ‐ 112 0.4 0.4 0.0 0.0 0.4 0.9 0.8 11.1 3.2 0.3 8.1 ± 0.8 3.9 ± 1.5 0.5 ± 0.8
430
151561 138 ‐ 141 0.4 0.7 0.0 0.0 0.6 0.8 1.1 10.1 4.3 0.4 12.4 ± 0.5 4.0 ± 0.7 0.3 ± 0.3 9.0 ± 1.9 144 ‐ 147 0.4 0.6 0.0 0.0 0.6 0.9 1.4 10.9 5.8 0.5 9.7 ± 0.6 6.2 ± 1.1 0.6 ± 0.8 11.9 ± 2.6
760
200664 116 ‐ 125 0.5 0.3 0.0 0.0 0.2 1.0 0.5 12.4 2.1 0.2 9.1 ± 0.9 2.3 ± 1.9 0.3 ± 0.5 125 ‐ 135 0.5 0.2 0.0 0.0 0.2 0.8 0.7 10.0 2.6 0.3 7.1 ± 1.0 2.0 ± 2.3 0.3 ± 0.7 135 ‐ 142 0.3 0.2 0.0 0.0 0.2 0.6 0.4 7.1 1.5 0.2 3.8 ± 0.9 2.3 ± 1.9 0.6 ± 1.3
150
753663 116 ‐ 118 0.3 0.4 0.0 0.0 0.4 5.9 ± 0.6 2.0 ± 0.9 0.3 ± 0.4 5.6 ± 1.8 118 ‐ 119 0.1 0.3 0.0 0.0 0.7 2.8 ± 0.6 2.4 ± 1.1 0.9 ± 1.1 4.1 ± 2.5 131 ‐ 133 0.3 0.2 0.0 0.0 0.2 3.6 ± 0.6 2.2 ± 1.1 0.6 ± 0.8 6.6 ± 2.2 145 ‐ 147 0.3 0.5 0.0 0.0 0.7 6.1 ± 0.5 3.2 ± 0.7 0.5 ± 0.5 7.2 ± 2.0
680
1‐ Well 161435 is adjacent to Jordan aquifer well 677176 sampled for this study.
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5. Geochemical evaluation of salinity sources and radium response to salinization of an unconfined coastal aquifer in Morocco
5.1 Introduction
Coastal aquifer salinization is a global phenomenon, primarily the result of the high level
of groundwater demands in densely‐populated and groundwater‐dependent coastal areas, but
also exacerbated by sea level rise (Vengosh, 2003). Salinization by modern seawater, as opposed
to mixing with other natural saline groundwater sources, is often characterized by a rapid
salinity increase as a relatively sharp freshwater‐saltwater interface is drawn across the open
interval of a well. In contrast, the mixing of natural saline groundwaters into fresh waters results
in a more gradual salinity increase due to the more diffuse nature of salinity interfaces with old
saline waters such as old unflushed seawater (Jones et al., 1999; Vengosh, 2003). In practice,
heavily pumped urban coastal aquifers may be impacted by a mixture of seawater, other natural
saline waters, and high‐nitrate wastewater (e.g. Vengosh et al., 1999, 2005; Fakir et al., 2002;
Land et al., 2004; Bouchaou et al., 2008), which seriously complicates prediction of future
salinity trends and planning of adaptation mechanisms. Modern seawater intrusion can be
diagnosed and distinguished from other salinization mechanisms using a combination of ion
ratios and environmental isotopes including Br/Cl‐, Na/Cl‐, Mg/Ca, δ2H, δ18O, δ34SSO4, and δ11B
(Jones et al., 1999; Vengosh, 2003). Numerous case studies in North Africa and the
Mediterranean region have documented the extent of coastal aquifer salinization, especially in
heavily exploited urban aquifers (Giménez and Morell, 1997; Ben Kabbour and Toto, 2001;
Ekwurzel et al., 2001; Yechieli et al., 2001; Fakir et al., 2002; Vengosh, 2003; Sivan et al., 2005;
Vengosh et al., 2005; Ben Hamouda et al., 2008; Bouchaou et al., 2008; Kouzana et al., 2009).
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Reconnaissance isotopic investigations in the Souss‐Massa region of Morocco, near the city of
Agadir (Figure 31), have indicated that intrusion of modern seawater is occurring, but that
salinity is also derived from other natural and anthropogenic sources (Ekwurzel et al., 2001;
Bouchaou et al., 2008).
In addition to the increase of total dissolved solids, which alone can render a water
resource unusable, salinization stimulates water‐rock interaction that could cause mobilization
of naturally‐occurring trace elements and further impair water quality. Examples of naturally‐
occurring trace elements that may respond to salinization events include (1) arsenic, which
more typically responds to pH and redox conditions than salinity but also is adsorbed less
effectively in waters of high Ca/Na subsequent to mixing of fresh and saline inland waters
(Scanlon et al., 2009); (2) boron, which tends to be lost from seawater by adsorption onto
marine clays during seawater intrusion (Vengosh and Spivack, 1999), but can be released from
adsorption sites during freshening reactions (Ravenscroft and McArthur, 2004; Faye et al., 2005;
Halim et al., 2010); and (3) radium, which is less efficiently adsorbed as salinity increases
(Kraemer and Reid, 1984; Moise et al., 2000; Sturchio et al., 2001; Tomita et al., 2010). In some
cases (e.g. B(OH)30, H3AsO3
0), natural contaminants may even remain in saline waters treated by
reverse osmosis due to their uncharged nature at near‐neutral pH (Prats et al., 2000; Walker et
al., 2008; Vinson et al., in press). Finally, over‐exploited urban coastal aquifers are also subject
to anthropogenic contamination associated with extensive development and human activities
(Vengosh et al., 2005; Bouchaou et al., 2008). The water quality deterioration of coastal aquifers
is therefore often a result of multiple processes and not typically simple intrusion of seawater.
In general, the increase of aqueous radium with increasing salinity at near‐neutral pH
has been consistently observed along salinity gradients, but a more detailed understanding is
133
required due to significant uncertainties associated with the composition of saline waters.
Although desorption or cation exchange can be the dominant means of Ra release in extreme
salinization events (Wood et al., 2004), it is assumed that in typical groundwater settings, the
primary mechanism of Ra release in fresh to saline waters is alpha recoil (Krishnaswami et al.,
1982). The mechanism by which radium activities increase with salinity is that radium removal
becomes less efficient, rather than the source term becoming larger. Within this basic
understanding of radium behavior, unresolved issues remain. For example, the possible effects
imparted by relative cation or ion compositions in inhibiting radium removal in saline waters are
not as well documented. As an example of effects specific to major ion composition, it has been
documented that Ca, where present at high concentrations, successfully competes with Ra for
adsorption sites (Nathwani and Phillips, 1979b). In addition, because many saline aquifers are
also anoxic, and specifically seawater intrusion is often associated with sulfate reduction
(Magaritz and Luzier, 1985; Appelo and Postma, 1993; Jones et al., 1999; Sivan et al., 2005; de
Montety et al., 2008), the concurrent role of redox conditions in impairing radium adsorption
sites remains unresolved.
The suite of tracers utilized in this study assess not only the sources of salinity, which is
essential for anticipating future trends and mitigation measures, but also provide constraints on
specific mechanisms that relate to radium’s behavior in the groundwater system undergoing
salinization: (1) Multi‐isotopic radium analysis (224Ra, 226Ra, 228Ra) assesses Ra from U‐ vs. Th‐
derived sources and may also identify critical radium‐removing processes due to their half‐life
differences; (2) Strontium isotopes identify sources and mechanisms of divalent cation mobility
such as cation exchange; (3) Boron concentrations and B/Cl‐ ratios complement Sr isotopes
because B is commonly released during freshening events associated with cation exchange in
134
which sodium is released, which would not be recorded by Sr isotopes because Sr is removed
from the groundwater; and (4) Sulfur and oxygen isotopes of sulfate evaluate the sources of
sulfate, sulfate reduction processes, and the stability of barite ‐ which removes radium by
coprecipitation ‐ in sulfate‐reducing conditions. In addition to these isotopic tracers, additional
insights on the salinization process and salinity sources are gained through the δ2H, δ18O, and Br‐
/Cl‐, Na/Cl‐, Ca/Na, and Ca/Mg ratios.
5.2 Hydrogeology of the study area
In the study area including the city of Agadir in western Morocco (Figure 31), the main
aquifer sediments range from mid‐Pliocene to early Pleistocene in age and represent
transgressive marine sediments. The lower portion of the Pliocene deposits outcrop in the
structurally complex area near well M9‐41, at the northern edge of the study area, but in most
of the study area, the ~100 m thick older Pliocene (Plaisancian‐Astian) deposits are deeper than
the typical depth of water wells. These marine rocks include limestone, sandstone, coquina, and
marl. The primary aquifer for most of the area is the overlying, Pliocene to Pleistocene
(Villafranchian) transgressive deposits. This unit is composed of ~100 m of marine clay, sand
gravel, and limestone, overlain by a wedge‐shaped unit composed of conglomerate, sandy
limestone, and shelly sandstone. Further east (inland) these deposits are of fluvial‐lacustrine
origin. This upper unit thickens north to south from ~15 to ~75 m approximately along a line
from Agadir to the vicinity of sample AC‐50 in Figure 31 (Dijon, 1969). The known well depths
(Table 17) correspond mostly to the Villafranchian aquifer, which is interpreted as an
unconfined aquifer. The sandstone and conglomerate layers in the shallow aquifer are especially
important as water‐yielding units (Dijon, 1969).
135
5.3 Methods
5.3.1 Sample collection
Water samples were collected from drilled wells, primarily those delivering large
quantities of water for landscaping, public supply, and commercial uses (Figure 31). Samples
were collected as close to the wellhead as possible and while pumps were operating. However,
in some cases it was only possible to collect water samples after storage tanks. pH, temperature,
specific conductance, and dissolved oxygen were measured in the field, minimizing aeration
during dissolved oxygen measurement whenever possible. Filtered samples were collected for
trace metals, major anions, and isotopes of boron, sulfur, and strontium using syringe‐tip 0.45
µm filters. The trace metal samples were collected and acid‐preserved in new polythelene
bottles that had been acid‐washed using trace metal grade acids. Unfiltered samples were
collected without headspace for alkalinity and stable isotopes.
5.3.2 Major and trace element analysis
Major cations (Ca, Mg, Na, Si, Sr) were analyzed by direct current plasma spectrometry.
Potassium analysis was performed by flame atomic absorption spectrometry. All were calibrated
using a solution prepared from commercially obtained single‐element standards. Major anions
(Cl‐, Br‐, NO3‐, SO4
2‐) were analyzed by ion chromatography at favorable dilution factors for each
anion. Boron, barium, and uranium concentrations were analyzed by inductively coupled plasma
mass spectrometry with B and Ba calibrated using NIST 1643e solution and U calibrated with a
plasma‐grade single‐element standard.
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Figure 31: Map showing sampling locations and chloride concentrations in Agadir area, western Morocco. Unlabeled square data points are from Tagma (2005). Base map content obtained from ESRI ArcGIS Online (World Topographic Map, World Terrain Base, and World
Reference Overlay).
5.3.3 Environmental isotopes (δ2H, δ18O, δ34SSO4, δ18OSO4,
87Sr/86Sr)
Hydrogen and oxygen isotopes of water samples were analyzed at the Duke University
Environmental Isotope Laboratory with a ThermoFinnigan thermochemical analyzer and
Delta+XL isotope ratio mass spectrometer on the peaks of H2 and CO gas, normalized to the
VSMOW and VSLAP standards and reported as δ2H and δ18O with precision of 0.3‰ and 0.1‰,
respectively. Sulfur and oxygen isotope ratios of sulfate (δ34SSO4, δ18OSO4) were analyzed from
BaSO4 precipitates which were prepared by extracting ~7 mg of sulfate from filtered water
samples using BaCl and filtering the resulting precipitate onto glass fiber filters. The sulfate was
137
analyzed at the University of Calgary Isotope Science Laboratory. Sulfur isotope ratios were
analyzed by continuous flow isotope ratio mass spectrometry, normalized to VCDT using NIST
and IAEA sulfate standards, and reported as δ34S with precision of 0.25‰. Oxygen isotopes of
sulfate were analyzed by reacting BaSO4 in a thermochemical analyzer and determining the δ18O
of the resulting CO gas using a Delta+XL mass spectrometer. As a check on the purity of BaSO4
precipitate, δ18O data are reported only from samples containing 27±2% oxygen. These ratios
were normalized to VSMOW, also using NIST and IAEA sulfate standards, and reported as δ18O
with precision of 0.3‰.
Strontium isotope ratios were analyzed by the thermal ionization mass spectrometry
(TIMS) technique at Duke University. Samples were prepared for Sr isotope measurement by
drying down ~3 µg of Sr in a Teflon vial in a laminar flow hood, digesting the sample in high‐
purity (Optima) HNO3, and extracting Sr using Eichrom SR‐B50‐S resin. Samples were then
loaded onto degassed Re filaments in a Ta load solution. Analysis was performed in positive ion
mode and all sample runs included replicates of the NIST 987 standard (long‐term average
0.710245 ± 0.000010; 1σ, n=109).
5.3.4 Radium
Large‐volume, unfiltered water samples collected into rinsed plastic containers (25‐50 L)
were concentrated onto manganese oxide‐coated acrylic fibers on the same day of collection.
17 of the 19 radium samples were collected onto two sequential columns of Mn oxide fibers to
confirm the efficient extraction of Ra. Fibers were analyzed for the short‐lived radium‐224 upon
arrival at Duke University within 3‐8 d of collection. Analysis methods for radium include: (1)
delayed coincidence counting for 224Ra, in which count rates were corrected for decay since
138
collection and alpha interference from 223Ra (Moore and Arnold, 1996; Garcia‐Solsona et al.,
2008) and calibrated using an efficiency appropriate for hand‐squeezed fibers; (2) alpha
counting of 222Rn at equilibrium with 226Ra incubated in a sealed glass tube using a RAD7 radon‐
in‐air monitor (Kim et al., 2001); and (3) gamma spectrometry for 228Ra of fibers compressed
into a wafer 65 mm in diameter and 5 mm tall using the 911 peak of 228Ac. In all cases, samples
were corrected for background count rate and standard propagation of error equations were
applied to background subtraction and the division of radium between the first and second
columns of Mn oxide (Eaton et al., 2005). The reported 2σ errors represent random error from
counting statistics only and do not consider other factors such as moisture content variation
(224Ra) or geometric variation (228Ra). The sample‐specific 2σ counting errors provide an estimate
of detection limit, defined as the activity with ±100% precision and 95% probability of being
detected (1.96σ; Eaton et al., 2005).
5.3.5 Speciation calculations
Saturation index and speciation were calculated using the PHREEQC version 2
geochemical code (Parkhurst and Appelo, 1999). Thermodynamic constants were added to the
default database for the species Ra2+, RaSO40, RaCl+, and RaHCO3
‐ (Langmuir and Riese, 1985).
5.4 Results
5.4.1 Major element chemistry
Total dissolved solids (TDS) vary across one order of magnitude within the study area
(e.g. Na 96‐1050 mg L‐1; Cl‐ 242‐2120 mg L‐1; SO42‐ 63‐816 mg L‐1; Table 17). The spatial
distribution of salinity is irregular. While the highest‐chloride waters are, in general, seen in
wells closest to the Atlantic Ocean, there is no systematic trend in salinity with distance relative
139
to the coastline (Figure 31). Nitrate concentrations in these waters (median 27.6 mg L‐1 as NO3‐,
range 0.1‐90.3 mg L‐1) in some cases exceed Moroccan allowable limits and World Health
Organization guidelines for drinking water (50 mg L‐1 as NO3‐; World Health Organization, 2008;
Tagma et al., 2009). Where dissolved oxygen data are available, waters are generally oxic (Table
17), although these are maximum values due to the possibility of aeration prior to sampling of
some wells.
Overall, three general groups of waters were identified; (1) relatively fresh (TDS
<1800 mg L‐1) waters in which bicarbonate exceeds 25% of anions (as equivalents; samples M9‐
42, M9‐43, M9‐45 AC‐50, and M10‐50); (2) Cl‐dominated waters (51‐81% of anions) with SO42‐
/Cl‐ < 0.25 and TDS > 2000 mg L‐1; and (3) a distinctive high‐sulfate (43‐45% of anions), low‐
chloride water with SO42‐/Cl‐ > 0.5 and TDS > 1000 mg L‐1 (samples M9‐46 and M9‐47) located
near the Oued (Wadi) Souss (Figure 31; Table 17). The cation composition of samples overlaps
significantly between these three salinity groups. The proportion of sodium among the major
cations ranges from 26 to 63% (median 43%). The three lowest sodium proportions (26‐28%) are
observed in the three wells near the Oued Souss, which comprise the group 3 waters plus well
AC‐50 (Figure 31). Only in five group 2 wells closest to the ocean does Na (or Na+K) exceed 50%
of total cations present in the waters. In the rest of the waters, including all of group 1 and
group 3 and approximately half of group 2, the combined Ca+Mg ions are more abundant than
Na. Although the most sodium‐dominated waters are located near the Atlantic Ocean, molar ion
ratios directly associated with seawater exhibit significant overlap within the waters in this study
(Table 18) and exhibit large deviations from expected seawater values, even in group 2 waters
where seawater influence is most plausible: (1) SO42‐/Cl‐ ratios is above the seawater value of
0.05 (group 2 median 0.11, range 0.04‐0.24); (2) Na/Cl‐ is below the seawater value of 0.86
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(group 2 median 0.69, range 0.44‐0.87); Ca/Na is substantially above the seawater value of and
0.02 (group 2 range 0.13‐0.41); and (4) Ca/Mg is far above the seawater value of 0.20 (group 2
median 0.88, range 0.73‐1.07). Another ion ratio, Br‐/Cl‐, exhibits distinctive values for seawater
and other salinity sources including brine, evaporite dissolution, and wastewater. Br‐/Cl‐ ratios
were observed throughout the study area near and slightly below the marine value of 1.5 x 10‐3
(median 1.4 x 10‐3; Figure 32; Table 18). The lowest values of Br‐/Cl‐ were observed in low‐Cl‐,
high‐SO4‐/Cl‐ waters of group 3, and the highest values of Br‐/Cl‐ (near the seawater value) were
observed in the coastal wells of group 2 (Figure 32).
Boron concentrations (Table 19) and the B/Cl‐ ratio (Table 18) also vary significantly
within the data set (Table 19) and some correlations between indicators of salinity were
observed. In general, the B/Cl‐ ratio exceeds the seawater value of 7.8 x 10‐4. Strong correlations
were observed between boron and sodium and chloride concentrations (Spearman rank
coefficients ρ=0.94 and 0.85, respectively; Figure 33). Although boron concentration was not
correlated with bicarbonate concentration, the B/Cl‐ ratio is somewhat correlated with
bicarbonate concentration (ρ=0.48; Figure 33), with the exception of M9‐42, an isolated fresh
water sample surrounded by saline wells. B/Cl‐ is also moderately correlated with the Na/Cl‐
ratio (ρ=0.57; Figure 33).
5.4.2 Environmental isotope ratios
δ18O and δ2H values vary within a broad range in the coastal aquifer of Agadir (‐40.6‰
to ‐16.0‰ and ‐6.5‰ to ‐3.2‰ respectively; Table 19). The data plot roughly along the global
meteoric water line (Figure 34) although with apparent deviations. The slight deviations from
the global meteoric water line are not supported by any relationships between δ18O and
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Table 17: Major element concentrations and saturation index values of groundwater samples.
Depth Concentrations in mg L‐1 Saturation index
Sample (m) Date T (°C) pH O2 Ca Mg Na K Si Cl‐ Br‐ NO3‐ SO4
2‐ HCO3‐ Calcite Barite
Group 1 AC‐50 11 Mar 2008 24.3 6.9 1.7 125 107 123 7.8 15.0 257 0.46 90.3 139 0.09 ‐0.04M9‐42 19 May 2009 22.2 7.4 8.2 101 47.9 126 5.1 7.1 242 0.62 9.5 155 232 0.13 0.21M9‐43 19 May 2009 25.1 7.2 5.6 127 103 160 6.2 11.0 396 1.35 0.4 149 465 0.33 0.05M9‐45 19 May 2009 25.3 6.8 5.9 179 112 225 6.2 14.2 398 1.05 13.2 358 474 0.01 ‐0.12M10‐50 130 18 Mar 2010 26.1 7.5 87.2 48.7 96.5 4.4 14.7 163 0.53 0.1 63.2 468 0.54 ‐0.08Group 2 AC‐43 11 Mar 2008 24.4 6.8 4.1 230 183 1050 9.3 13.9 1860 5.82 37.2 482 0.09 0.07M8‐11 18 Dec 2008 24.0 7.0 3.4 191 148 411 6.7 12.0 843 2.71 8.6 370 408 0.19 0.00M8‐12 18 Dec 2008 23.4 7.1 3.9 243 201 663 13.5 11.7 1530 4.81 27.6 444 413 0.26 0.35M9‐41 70 19 May 2009 24.0 6.8 6.5 259 138 376 6.7 14.8 1020 3.38 55.7 166 474 0.20 0.17M9‐44 19 May 2009 24.8 6.9 7.2 400 291 627 11.9 12.0 1870 6.21 43.2 605 402 0.27 0.14M9‐48 27 May 2009 24.8 6.9 4.2 254 165 418 7.2 13.6 1160 3.94 15.1 332 351 0.07 0.05M9‐49 80 29 May 2009 23.1 7.0 0.2 251 207 687 16.1 11.1 1480 4.57 28.9 423 457 0.23 0.29M10‐51 70 18 Mar 2010 23.9 7.1 236 192 966 9.8 12.9 1880 6.42 59.5 555 585 0.35 0.02M10‐52A 18 Mar 2010 24.2 7.0 169 111 428 5.4 14.1 703 1.97 45.3 450 560 0.22 ‐0.10M10‐52B 18 Mar 2010 22.8 7.1 182 117 412 6.2 13.8 739 2.09 33.3 483 534 0.34 ‐0.04M10‐53 18 Mar 2010 23.6 7.0 363 230 777 9.1 14.0 2120 7.10 20.9 500 455 0.34 0.14M10‐54 18 Mar 2010 25.6 6.9 341 192 478 5.8 14.9 1680 5.50 45.8 188 370 0.25 0.25M10‐55 90 18 Mar 2010 25.3 7.1 280 167 511 6.8 14.5 1440 4.71 71.0 262 357 0.35 0.29Group 3 M9‐46 20 May 2009 23.3 6.9 0.5 242 195 235 6.0 12.3 433 0.58 25.2 816 490 0.16 0.30M9‐47 61 20 May 2009 23.7 7.2 4.1 165 112 155 3.8 11.6 236 0.46 1.2 522 481 0.35 0.27
Seawater (Millero and Sohn, 1991) 8.2 412 1280 10800 399 19400 67 2710 114
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Table 18: Element and ion ratios diagnostic of salinity sources in the coastal aquifer.
Na/Cl‐ Ca/Mg Ca/Na B/Cl‐ Br‐/Cl‐ SO42‐/Cl‐
Group 1 AC‐50 0.79 0.7 0.14 1.5 x 10‐3 8.0 x 10‐4 0.11 M9‐42 0.80 1.3 0.46 2.7 x 10‐3 1.2 x 10‐3 0.24 M9‐43 0.63 0.8 0.46 1.6 x 10‐3 1.5 x 10‐3 0.14 M9‐45 0.91 1.1 0.52 1.8 x 10‐3 1.2 x 10‐3 0.14 M10‐50 0.94 0.9 0.23 2.5 x 10‐3 1.5 x 10‐3 0.24 Group 2 AC‐43 0.56 0.9 0.35 4.2 x 10‐3 1.4 x 10‐3 0.11 M8‐11 0.52 0.8 0.37 1.4 x 10‐3 1.4 x 10‐3 0.12 M8‐12 0.72 0.7 0.21 1.0 x 10‐3 1.4 x 10‐3 0.11 M9‐41 0.74 0.7 0.58 1.2 x 10‐3 1.5 x 10‐3 0.20 M9‐44 0.87 1.0 0.46 7.0 x 10‐4 1.5 x 10‐3 0.33 M9‐48 0.67 0.7 0.21 6.2 x 10‐4 1.5 x 10‐3 0.11 M9‐49 0.57 1.1 0.40 1.3 x 10‐3 1.4 x 10‐3 0.06 M10‐51 0.86 0.9 0.25 3.7 x 10‐3 1.5 x 10‐3 0.24 M10‐52A 0.57 1.0 0.27 2.2 x 10‐3 1.3 x 10‐3 0.09 M10‐52B 0.44 1.1 0.41 2.0 x 10‐3 1.3 x 10‐3 0.04 M10‐53 0.55 1.0 0.31 2.1 x 10‐3 1.5 x 10‐3 0.07 M10‐54 0.84 0.8 0.59 8.2 x 10‐4 1.5 x 10‐3 0.70 M10‐55 1.02 0.9 0.61 1.2 x 10‐3 1.5 x 10‐3 0.82 Group 3 M9‐46 0.87 0.8 0.13 1.1 x 10‐3 6.0 x 10‐4 0.10 M9‐47 0.75 0.8 0.27 1.9 x 10‐3 8.7 x 10‐4 0.16 Seawater 0.86 0.2 0.02 7.8 x 10‐4 1.5 x 10‐3 0.05
Figure 32: Plot showing relationship between bromide and chloride concentration. Dashed line represents seawater ratio.
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Figure 33: Boron concentrations in relationship to sodium and chloride concentrations. Dashed lines represent hypothetical dilution of seawater.
chloride, nitrate, or total dissolved solids concentrations in the data, which might indicate
seawater and/or evaporated wastewater, nor is δ18O associated with the three salinity groups.
Sulfur and oxygen isotope ratios are somewhat bimodal, with a group of samples exhibiting a
narrow range of positive values (11.6‐17.3 ‰), and the two high‐sulfate samples (group 3
above) exhibit distinctive, negative δ34SSO4 values (‐9.1 to ‐8.0 ‰). The higher values of δ34SSO4
144
generally correspond with higher δ18OSO4 (6.5‐11.6 ‰), and the lowest value of δ34SSO4
corresponds with the lowest value of δ18OSO4 (Figure 35). Strontium isotope ratios fall within a
narrow distribution (0.70835‐0.70887; Figure 36) for all but three samples in the southeastern
part of the study area (0.70929‐0.70962, representing Group 3 plus sample AC‐50) exhibiting
more radiogenic values than other samples and modern seawater (0.70924; Table 19).
Strontium isotope ratios are not correlated with strontium concentration.
Figure 34: δ18O plotted against δ2H of Agadir area groundwater. GMWL denotes the global meteoric water line, δ2H=8 x δ18O + 10‰ (Craig, 1961).
5.4.3 Radium, barium, and uranium
Overall, radium activities are significantly below WHO guidance levels for drinking water
(100 mBq L‐1 for 228Ra, 1000 mBq L‐1 for 224Ra and 226Ra; World Health Organization, 2008)
throughout the study area and thus below levels of concern for water quality (224Ra, 226Ra, and
228Ra are all below 40 mBq L‐1; Figure 37). Several significant correlations were observed among
the radium isotope data: (1) radium isotopes are generally correlated with each other: 224Ra and
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Figure 35: Sulfur and oxygen isotope ratios of the sulfate ion for selected water samples. Lines indicate hypothetical two‐component mixing curves among possible sulfate sources. Box marked “gypsum” represents the range of δ34SSO4 and δ
18OSO4 in Mesozoic marine sulfate minerals (Claypool et al., 1980; Strauss, 1997).
Figure 36: Histogram showing 87Sr/86Sr values of groundwater in relation to potential Sr sources. Bars represent ranges of marine 87Sr/86Sr values (McArthur et al., 2001) of rocks
encountered in the study area or the adjacent Atlas Mountains.
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Figure 37: Plots depicting radium‐224, ‐226, and ‐228 activities in relation to chloride, sulfate, and each other. Circled data points represent samples with 87Sr/86Sr > 0.709.
147
Table 19: Environmental isotope ratios, trace metal concentrations, and radium activities.
Radium activities in mBq L‐1
Sample δ2H (‰)
δ18O (‰)
δ34SSO4 (‰)
δ18OSO4 (‰) Sr (mg L‐1) 87Sr/86Sr
224Ra 226Ra 228Ra B
(µg L‐1)Ba
(µg L‐1)
% aBaas
BaSO40
% aRa‐226as
RaSO40
U (µg L‐1)
Group 1 AC‐50 16.7 8.5 1.04 0.709622 12.4 ± 1.3 9.8 ± 1.5 4.9 ± 2.2 117 53 20 23 3.7M9‐42 ‐22.2 ‐3.5 11.6 10.4 0.89 0.708676 5.1 ± 0.4 3.9 ± 0.7 5.8 ± 2.1 197 61 26 29 2.0M9‐43 ‐16.0 ‐3.4 ‐2.2 11.6 1.67 0.708469 16.7 ± 1.3 25.5 ± 1.6 35.8 ± 8.8 198 63 21 24 1.3M9‐45 ‐30.6 ‐4.9 16.0 2.10 0.708348 6.4 ± 0.6 2.8 ± 0.6 8.8 ± 2.8 220 23 38 42 10.6M10‐50 ‐25.9 ‐5.5 0.93 0.708684 18.1 ± 1.2 27.7 ± 1.5 20.1 ± 4.7 126 80 12 14 1.1Group 2 AC‐43 15.2 6.1 2.81 0.708691 19.9 ± 1.3 4.2 ± 1.2 6.9 ± 1.8 2350 41 39 42 9.3M8‐11 2.39 0.708484 24.0 ± 1.7 8.5 ± 1.1 12.6 ± 3.6 356 33 37 40 6.5M8‐12 4.42 0.708770 26.8 ± 1.8 8.2 ± 0.9 20.9 ± 5.0 472 75 38 41 7.3M9‐41 ‐19.3 ‐3.2 13.0 6.5 1.73 0.708671 25.3 ± 1.8 20.4 ± 2.1 18.2 ± 5.1 358 102 20 22 5.0M9‐44 ‐29.0 ‐4.8 16.4 9.1 5.78 0.708524 34.5 ± 2.4 7.7 ± 0.9 34.8 ± 8.4 401 46 41 44 12.8M9‐48 ‐38.9 ‐5.9 17.3 2.95 0.708396 23.6 ± 1.6 5.8 ± 1.0 13.4 ± 3.8 219 46 32 35 6.4M9‐49 ‐31.6 ‐4.9 17.3 9.2 4.08 0.708727 32.4 ± 2.1 8.0 ± 1.0 18.5 ± 5.0 565 69 36 39 6.9M10‐51 ‐24.5 ‐3.9 3.09 0.708652 14.8 ± 1.1 6.3 ± 0.9 9.0 ± 2.5 2120 32 42 46 10.5M10‐52A ‐40.6 ‐6.5 2.00 0.708576 17.2 ± 1.3 5.0 ± 1.0 9.6 ± 2.5 462 21 43 47 9.0M10‐52B ‐33.5 ‐5.1 2.14 0.708579 455 22 44 10.4M10‐53 ‐35.7 ‐5.6 4.84 0.708602 34.2 ± 2.3 10.6 ± 1.0 16.4 ± 3.9 1380 49 38 41 10.3M10‐54 ‐20.9 ‐3.8 2.13 0.708865 41.2 ± 2.6 17.3 ± 1.3 32.8 ± 6.8 420 145 19 21 4.4M10‐55 ‐26.5 ‐4.2 2.23 0.708726 30.5 ± 2.3 17.0 ± 1.3 35.9 ± 7.6 505 105 26 29 5.7Group 3 M9‐46 ‐27.8 ‐4.5 ‐8.0 2.10 0.709516 6.5 ± 0.5 1.8 ± 0.6 6.0 ± 2.2 147 37 55 59 11.2M9‐47 ‐34.3 ‐6.0 ‐9.1 2.8 1.53 0.709287 8.2 ± 0.5 3.6 ± 0.6 12.1 ± 3.1 137 39 49 53 5.7Seawater 0.0 0.0 21.0 1 9.5 1 7.9 2 0.70924 1 4600 3 14 2 1‐ Clark and Fritz, 1997 2‐ Millero and Sohn, 1991 3‐ Vengosh and Spivack, 1999
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228Ra (ρ=0.71), 228Ra and 226Ra (ρ=0.67), and 224Ra and 226Ra (ρ=0.55). The relationship between
radium and salinity is more complex. 224Ra, the shortest‐lived Ra isotope, is well correlated with
Mg, Ca, Na, K, and Cl‐ but not with sulfate concentration. In contrast, the longest‐lived radium
isotope 226Ra is essentially uncorrelated with salinity, except that it is significantly negatively
correlated with sulfate concentration (ρ = ‐0.60). Radium‐228, of intermediate half‐life, exhibits
correlations between those of 224Ra and 226Ra (Table 20). Like radium isotopes and salinity
overall, barium concentrations vary within approximately one order of magnitude
(21‐145 µg L‐1). Barium concentration is well correlated with 226Ra activity (ρ=0.72) and exhibits
progressively weaker correlations with 228Ra (ρ=0.61) and 224Ra (ρ=0.51) activities. Barium’s
relationship to salinity is essentially identical to that of 226Ra, but the opposite of 224Ra and 228Ra
(Table 20). Moderate uranium concentrations were detected throughout the study area,
(1.1‐12.8 µg L‐1), below the international recommendation of 15 µg L‐1 (World Health
Organization, 2008). Uranium was not observed to follow clear spatial trends, but uranium
concentration is moderately inversely correlated with 226Ra activity and strongly positively
correlated with sulfate concentration (ρ= ‐0.59 and 0.86, respectively; Figure 38).
5 Discussion
5.5.1 Groundwater evolution and salinity source identification
The fresh water end member (represented approximately by Group 1 waters) exhibits a
Ca‐Mg‐HCO3‐ composition ((Ca+Mg/HCO3
‐ ≈ 0.5; Figure 39) that is characteristic of carbonate
mineral dissolution associated with initial equilibration of fresh recharging waters with the
carbonate‐rich aquifer solids. The ratio of 0.5 is based on the combination of 1 mole of soil or
microbal CO2 decreasing pH to stimulate 1 mole of Ca(Mg)CO3 dissolution (buffering), yielding 2
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moles of dissolved inorganic carbon (DIC) per mole Ca+Mg. Based on relationships between
limited δ13C and carbon‐14 activities, the evolution of DIC in the lower Souss plain, including the
study area, is defined by dissolution of carbonate minerals into the fresh water end member
(Bouchaou et al., 2008). The fresh, Ca‐Mg‐HCO3‐ waters at or near calcite saturation are
modified by inputs of saline water. As salinity increases, the Ca+Mg/HCO3‐ ratio increases
significantly (Figure 39), indicating both the Ca and Mg content of saline waters and the effects
of cation exchange. In contrast to increasing cation concentrations, bicarbonate concentrations
do not increase with salinity. Although conservative mixing of seawater and fresh, calcite‐
saturated groundwater is associated with calcite undersaturation (Sanford and Konikow, 1989),
inputs of Ca, including by ion exchange, may drive supersaturation with respect to calcite
(median saturation index 0.25; Table 17), which implies calcite precipitation in the saline waters
(de Montety et al., 2008) and would also support the interpretation that DIC is contributed to
waters primarily in the fresh water end member. Another possible component of DIC evolution
is the composition of wastewater, which could enable organic matter metabolism yielding
additional CO2 and thus higher HCO3‐ concentrations. However, without extensive δ13C or
organic carbon analysis, the possible role of microbial respiration cannot be evaluated.
Table 20: Spearman rank coefficients between major ion compositions, and radium and barium. Note (1) the decreasing values of ρ as radium half‐life increases (left to right); (2) the opposite trend associated with sulfate concentration; and (3) the close similarity between the
behavior of Ba and 226Ra, but not other Ra isotopes.
224Ra 228Ra 226Ra Ba Ca 0.79 0.45 0.14 0.27 Mg 0.66 0.27 ‐0.07 0.07 Na 0.62 0.22 ‐0.06 ‐0.04 K 0.50 0.16 0.08 0.10 Cl‐ 0.71 0.28 0.09 0.08 SO4
2‐ 0.06 ‐0.18 ‐0.60 ‐0.59
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Figure 38: Relationship between dissolved uranium and radium‐226 and sulfate concentration.
Figure 39: Relationship between chloride concentration and the (Ca+Mg)/HCO3‐ ratio. Note
that the trend extrapolates to (Ca+Mg)/HCO3‐ = 0.5 at zero salinity.
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Mixing of saline waters, such as seawater and wastewater, with the Ca‐Mg‐HCO3‐ fresh
water is the simplest possible model for the observed distribution of salinity. While δ18O, δ2H,
and chloride concentration should behave conservatively during mixing, both seawater and
evaporated wastewater would exhibit higher values of δ18O (closer to 0‰) and lower deuterium
excess than precipitation, making these two sources indistinguishable by these tracers alone.
Because δ18O and δ2H record mixing of waters rather than solutes, small volumetric
contributions of water, difficult to detect isotopically, could explain large variations in salinity.
Overall, the distribution of δ18O and δ2H is broadly consistent with regional groundwater of the
lower Souss plain (Bouchaou et al., 2008). δ18O and δ2H are also consistent with the global
meteoric water line with deuterium excess of 10, and consistent with Atlantic‐derived Moroccan
precipitation with deuterium excess near 10 (Ouda et al., 2008), rather than converging to
seawater with deuterium excess of 0, although deviations from the global meteoric water line
are present within the data (Figure 34). Overall, these results do not demonstrate clear mixing
relationships between two water sources.
Beyond the inadequate model of mixing of water sources, other environmental tracers
of salinity sources may explain the evolution of the coastal groundwater. The lack of distinct
geographic trends in salinity (Figure 31) and ion ratios (Table 18) implies that the coastal aquifer
receives a combination of salinity sources, possibly including modern seawater, sewage, and
Triassic (Atlas) evaporites (Hsissou et al., 2002; Bouchaou et al., 2008). Thus the salinity cannot
be described in terms of a single seawater intrusion front, but instead salinity is contributed by a
combination of natural and possibly anthropogenic sources, as is typical of heavily‐exploited
urban coastal aquifers (Giménez and Morell, 1997; Vengosh et al., 1999; Fakir et al., 2002;
Vengosh et al., 2005; Bouchaou et al., 2008). The general geographic areas of high chloride
152
concentrations (group 2) and sulfate concentrations (group 3) have been identified in previous
studies (Dijon, 1969; Hsissou et al., 2002; Tagma, 2005; Bouchaou et al., 2008); at a minimum,
the chloride concentrations and dissolved solids in group 2 seem to be increasing with time,
whereas in other areas there may be a basis for naturally‐occurring salinity in the shallow
aquifer, such as the high sulfate concentrations in group 3. The consistently high nitrate
concentrations (Table 17) clearly indicate recharge of wastewater to the water table in a
broadly‐distributed pattern across the aquifer. Additional evidence of non‐seawater salinity is
seen in the Br‐/Cl‐ ratios of Group 3 waters which, consistent with the findings of Hsissou et al.
(2002), are lower than the marine ratio (Figure 32) and could indicate the presence of
wastewater or evaporite‐derived chloride (Davis et al., 1998). As an additional complication to
the geographic distribution of salinity, waters may be diluted locally by recharge of waters such
as potable municipal water. Overall, although progressive salinization seems to be underway in
the Agadir urban groundwater system, it is not straightforward to clearly separate recent salinity
increases from natural saline water that has been present in the aquifer since pre‐development.
5.5.2 Cation exchange and 87Sr/86Sr
One form of water‐rock interaction, well documented in seawater intrusion settings, is
Ca/Na exchange resulting in Ca enrichment in the salinized groundwater relative to seawater
abundance (Nadler et al., 1980; Mercado, 1985; Jones et al., 1999; Vengosh et al., 2002; Sivan et
al., 2005). These source and removal mechanisms for alkaline earth metals directly affect
radium because the cation exchange environment in which Mg, Ca, and Sr are being released is
unfavorable for Ra removal by adsorption or ion exchange. For example, Ca successfully
competes with Ra for adsorption sites in high‐Ca waters, maintaining high Ra levels (Nathwani
153
and Phillips, 1979b). Cation exchange is here indicated by a notation such as ΔNa that
represents the difference of cation concentrations relative to seawater diluted to the same
chloride concentration as the groundwater. Assuming seawater to be the salinity source, this
notation indicates release or uptake of cations from exchange sites (Vengosh et al., 2002):
When plotted against ΔNa, ΔCa is close to the ‐1:1 line (equivalents ratio; Figure 40) indicating
direct cation exchange, but slightly higher in the divalent cations than can be accounted for by
cation exchange alone (Figure 40). This implies that some Ca release has occurred in addition to
that induced by base‐exchange reactions. Excess Ca (ΔCa+ΔMg+ΔNa+ΔK) is plotted against
ΔHCO3‐ to indicate the presence of cations that cannot be explained by saline inputs and ion
exchange alone (Figure 40). The probable source of additional Ca is net dissolution of carbonate
minerals and the effect of excess Ca is most apparent at samples with ΔNa <10 meq L‐1, that is,
those least affected by salinization (Figure 40), consistent with the Ca‐Mg‐HCO3‐ composition of
the fresh end member.
Although Sr isotopes can indicate conservative mixing relationships, including a few
seawater intrusion studies (Kim et al., 2003; Jørgensen et al., 2008), seawater strontium is often
not clearly traceable using 87Sr/86Sr of groundwater because base‐exchange reactions and/or
carbonate mineral dissolution release Sr from sources other than modern seawater with its
distinctive 87Sr/86Sr ratio of 0.70924 (Vengosh et al., 2002; Bouchaou et al., 2008; Lara and
Weber, 2010). Given that the 87Sr/86Sr of seawater recorded in primary marine carbonate and
evaporites has varied substantially during geologic time, these values may provide some
indication of geologic Sr sources although the immediate source of Sr to an aquifer is more
(3)
154
typically clays whose exchange sites record prior hydrogeochemical events. The observed values
of 87Sr/86Sr are both above and below the modern seawater ratio, but most waters exhibit less
radiogenic values, indicating additions of older marine‐derived strontium (Figure 36). The
majority of 87Sr/86Sr values in this study are less radiogenic than the inferred Pliocene age of the
aquifer sediments (0.7090‐0.7091) but significantly more radiogenic than the Triassic‐
Cretaceous (0.7072‐0.7083; McArthur et al., 2001) age of marine sedimentary rocks of the
western Atlas Mountains and other rocks adjacent to the study area (Bouchaou et al., 2008).
Also, in the waters exhibiting 87Sr/86Sr < 0.709, the ratio generally decreases progressively inland
(Figure 41). This may indicate a mixing relationship between modern seawater and Mesozoic
marine Sr, but could also indicate a mixture of Pliocene Sr with Mesozoic Sr, presumably
reworked from the nearby Atlas foothills into the Pliocene formations. However, the observed
87Sr/86Sr ratios are inconsistent with a mixture of Sr from modern seawater and the Pliocene
aquifer solids alone. Thus, 87Sr/86Sr is consistent with Sr release to the aquifer from marine‐
derived solids of pre‐Pliocene age but does not adequately test the presence or absence of
modern seawater. The three samples with 87Sr/86Sr > 0.709 are located along the Oued Souss
river and more inland than the other sampling sites (Figure 41). These values apparently
represent nonmarine Sr sources, which are abundant in the sediment source area of the
southern Souss valley including crystalline rocks (Bouchaou et al., 2008). The evidence of cation
exchange indicated by ion ratios and 87Sr/86Sr implies that Mg, Ca, Sr, Ba, and Ra are inefficiently
removed during the salinization process. This chemical mechanism is evaluated in subsequent
discussion.
155
5.5.3 Boron’s sensitivity to seawater intrusion and ion exchange
Boron is a useful trace element for diagnosing seawater intrusion because of the high B
concentration in seawater and its distinctive adsorption behavior. The expected behavior of
boron in a seawater intrusion event is to undergo adsorption upon contact with solids in the
previously fresh part of the aquifer such that the residual water exhibits B/Cl‐ ratios lower than
seawater values (Jones et al., 1999; Vengosh and Spivack, 1999; Vengosh et al., 2005). Because
the B adsorption coincides with cation exchange reactions, the co‐occurrence of low B/Cl‐, low
Figure 40: Indicators of ion exchange in relation to each other and 87Sr/86Sr. Data indicated by x symbols are from Tagma (2005).
156
Na/Cl‐, and high Ca/Na both diagnose the salinity source and provide constraints on the removal
of radium in the salinization environment.
The boron concentrations in the Agadir coastal aquifer are not consistent with the
expected boron behavior of seawater intrusion settings for two reasons: (1) B/Cl‐ ratios are
higher than the seawater value, especially at higher boron and chloride concentrations (Table
18; Figure 33); and (2) The high values of B/Cl‐ are inconsistent with the observed cation
exchange behavior of the aquifer, in which net Ca release occurs to groundwater, whereas B
Figure 41: Map showing measured values and inferred contours of 87Sr/86Sr ratios. Base map content obtained from ESRI ArcGIS Online (World Topographic Map base map).
157
increases would be associated with waters receiving Na from reverse cation exchange during
freshwater flushing (Ravenscroft and McArthur, 2004; Faye et al., 2005; Goldberg et al., 2008;
Halim et al., 2010). Possible explanations for this trend include (1) mixing of an additional
salinity source with high Na/Cl‐ and high B/Cl‐; or (2) boron desorption. The highest values of
B/Cl‐ occur at Na/Cl‐ ratios near or above seawater values (Figure 40). These are waters in which
seawater intrusion cannot be the sole source of salinity due to the expected Na/Cl‐ decline in
base‐exchange reactions. Instead, these elevated Na/Cl‐ and B/Cl‐ ratios could result from
partial freshening subsequent to salinization, in which the elevated B concentrations are the
result of desorption and elevated Na/Cl‐ is the result of cation exchange. Although cyclic
salinization and freshening (Russak and Sivan, 2010) have not been documented in this coastal
aquifer, one potential source of fresh water to the aquifer is recharge of municipal water, which
is locally blended with high‐TDS groundwater to produce irrigation water of acceptable salinity.
5.5.4 Sulfate sources and barite solubility
While the behavior of alkaline earth metals (Mg, Ca, Sr) has been examined to elucidate
salinity sources, ion exchange, and potential mechanisms of divalent cation removal, barium
warrants special attention because it is analogous to radium and undergoes similar removal
mechanisms, including the potential for radium removal by coprecipitation into barite. Barium
control by barite was considered by calculating the activities of Ba2+ and SO42‐ ions using the
PHREEQC geochemical code and considering the solubility product of 10‐9.97 at 25°C (Drever,
1997). Also, Ba2+ species activity is distinct from total Ba because a significant amount (median
37%, range 12‐55%) of total Ba activity is complexed as the BaSO40 species (Table 19). The
inverse relationship between barium and sulfate activities is consistently associated with barite
158
saturation (Figure 42), and implies that barium concentrations are controlled by barite
precipitation. The median barite saturation index in the coastal aquifer is 0.10 (range ‐0.12‐0.35;
Table 17).
Figure 42: Sulfate and Ba2+ ion activities of Agadir groundwater. Lines represent barite saturation index values based on solubility product at 25°C.
Thus, the sulfate content of the system apparently suppresses barium activity through
barite coprecipitation, which is itself made possible by sufficient barium and sulfate
concentrations as well as apparently oxic conditions. The latter aspect is particularly important
as sulfate‐reducing conditions can induce reductive dissolution of barite (Bolze et al., 1974),
resulting in elevated levels of aqueous Ra (Phillips et al., 2001; Landa, 2003; Martin et al., 2003).
Seawater intrusion may commonly result in sulfate‐reducing conditions by contributing high‐
sulfate water to high‐organic carbon coastal sediments (Appelo and Postma, 1993; Andersen et
al., 2005; Sivan et al., 2005; Yamanaka and Kumagai, 2006; de Montety et al., 2008). The large
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range of δ34SSO4 could imply that these variations are the result of sulfate reduction, which is
faster for isotopically light sulfate and thus causes δ34SSO4 and δ18OSO4 of the residual
groundwater to increase with a predictable slope when δ34SSO4 and δ18OSO4 are plotted against
one another (Krouse and Mayer, 1999; Berner et al., 2002; Tuttle et al., 2009). However, there
seems to be no sulfate reduction induced by seawater intrusion because δ34SSO4 never exceeds
the modern seawater value of 21‰, nor does δ18OSO4 exceed the modern seawater value of
9.5‰ (Figure 35). Instead of the elevated values of δ34SSO4 expected from sulfate reduction, a
trend leading to ~ ‐10‰ was observed (Figure 35). Such a negative value cannot be obtained
from seawater or evaporite‐derived gypsum, nor do anthropogenic sources of sulfate exhibit
such negative δ34SSO4 values (Barker et al., 1998; Krouse and Mayer, 1999; Hosono et al., 2010).
Instead, these ratios are distinctive of sulfate derived from oxidation of biogenic sulfide (e.g.
pyrite; Strauss, 1997; Krouse and Mayer, 1999).
From these observations, it is inferred that the variations in δ34SSO4 are the result of
mixing of sulfate sources rather than sulfate reduction and thus, fractionation of δ34SSO4 within
the various sulfate sources should be negligible. Four possible sulfate sources may be
contributing to the aquifer (Figure 35): (1) Mesozoic marine gypsum from the Atlas Mountains
reworked in the coastal formations, represented by a hypothetical saline water with 3000 mg L‐1
sulfate with δ34SSO4 ≈ 15‰ and 1000 mg L‐1 Cl‐; (2) sulfate derived from oxidation of marine
sulfide, represented by a fresh water containing 500 mg L‐1 SO42‐ with δ34SSO4 = ‐9‰ and
250 mg L‐1 Cl‐ (similar to sample M9‐47); and/or (3) modern seawater with δ34SSO4 = 21‰
(Krouse and Mayer, 1999); and (4) anthropogenic sulfate from wastewater, not represented on
Figure 35. The mixing model indicates that the presence of sulfate sources 1 through 3 is
consistent with the observed isotope ratios, and most of the samples plot along one of the
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hypothetical two‐component mixing curves between seawater‐ and gypsum‐derived sulfate and
seawater‐ and sulfide‐derived sulfate (Figure 35). In general, it was possible to explain the
observed data without introducing anthropogenic sulfate to the mixing model.
A similar mixing model was not attempted using δ18OSO4 because the sulfate derived
from sulfide oxidation exhibits ambiguous δ18OSO4 if the sulfide oxidation pathway is not known.
Sulfide oxidation by molecular oxygen imparts a different δ18OSO4 than if water molecules are
the oxygen source (Lloyd, 1968;Tuttle et al., 2009). Despite these complications, the pattern in
δ18OSO4 is broadly consistent with the interpreted δ34SSO4 pattern, yet inconsistent with mixing of
only seawater and Mesozoic gypsum‐derived sulfate (Figure 35). In contrast to other studies of
seawater intrusion, some waters in this study clearly exhibit sulfide oxidation and no samples
indicate sulfate reduction. This apparent lack of widespread sulfate reduction also implies that
barite precipitation is plausible.
5.5.5 Radium activities
224Ra, 226Ra, and 228Ra activities are correlated with each other (Figure 37, first row), but
correlations of Ra isotopes with the components of salinity vary within the half‐lives of the
isotopes being considered (half‐lives of 224Ra < 228Ra < 226Ra). The chloride ion, whose primary
interaction with Ra is to form the soluble RaCl+ complex, would be expected to enhance radium
solubility. Chloride concentration is more strongly positively correlated with 224Ra than 228Ra or
226Ra (Table 20; Figure 37, first row), that is, the correlation weakens with increasing radium
isotope half‐life. The sulfate ion would be expected to form the uncharged complex RaSO40,
which could enhance Ra solubility in low‐Ba waters. However, saturation of the insoluble BaSO4
(section 5.5.4) suggests that elevated sulfate concentrations may be associated with Ra removal.
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In contrast to the correlations with chloride, sulfate is negatively correlated with long‐lived 226Ra
but is essentially uncorrelated with shorter‐lived 228Ra and 224Ra (Table 20; Figure 37, third row).
As an enhancer of aqueous radium, chloride is more strongly associated with short‐lived
224Ra because short‐lived isotopes are rapidly replenished in relationship to the rates of removal
processes. In contrast, removal terms of radium are more closely associated with longer‐lived
226Ra because its replenishment by recoil is relatively slow, preserving the signature of chemical
removal of radium. Thus, the shortest‐lived isotope is most responsive to the Ra‐mobilizing
effects of salinization. Considering 224Ra, 226Ra, and 228Ra together, the radium‐salinity
relationship depends on the balance of major anions present in the water rather than simply
total dissolved solids. Overall, the SO42‐/Cl‐ ratio is significantly negatively associated with 224Ra
(ρ=‐0.82), 226Ra (ρ=‐0.60), and 228Ra (ρ=‐0.55) in approximate proportion to the random
(counting) error associated with radium isotope analysis. Seawater‐like values of SO42‐/Cl‐
(~0.05), which are mostly associated with high chloride concentrations overall, exhibit the
highest levels of aqueous Ra (Figure 37, fourth row). If the three samples with distinctive
87Sr/86Sr > 0.709 (circled points in Figure 37), possibly representing different aquifer conditions
in the southeastern part of the study area (sample AC‐50 plus the Group 3 samples with high
SO42‐ and δ34SSO4), are excluded from the above analysis, the correlations between radium and
SO42‐/Cl‐ deteriorate somewhat (ρ=‐0.73 for 224Ra, ρ=‐0.45 for 226Ra, and ρ=‐0.41 for 228Ra).
Although it is possible that this observed pattern is the result of the geographic distribution of
uranium and thorium in aquifer rocks, that is, that the high‐Cl‐ waters occupy higher‐Th rocks
and the high‐SO42‐ waters occupy lower‐Th rocks, the chloride‐based and sulfate‐based
correlations with radium isotope activities are independent. Another independent line of
evidence is seen in the uranium data. The positive correlation between SO42‐ concentration and
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uranium concentration and negative correlation between SO42‐ concentration and 226Ra activity
(Figure 38) imply that high‐sulfate waters do not occupy low‐uranium rocks. This suggests that
either U or Ra mobility is affected by sulfate concentrations. Speciation calculations suggest that
sulfate complexes for oxidized uranium are negligible, whereas Ra‐sulfate interactions have a
clear ability to remove Ra through barite precipitation.
The precipitation of barite is probably a significant removal mechanism for Ra, indicated
by the correlation between Ra2+ and Ba2+ ion activities (Figure 43), even though the geochemical
sources of these elements differ (only radium is released by alpha recoil). Further evidence is
seen in the negative correlation between Ra2+ and SO42‐ ion activities, similar to the behavior of
Ba2+ (Figure 43). Overall, barite is a potentially significant removal mechanism for radium that
removes some of the radium released by recoil, whereas the RaSO40 complex is unavailable to
barite and diminishes the effectiveness of removal by barite (median 21% of total Ra, range
9‐40%; Table 17). Perhaps given the strong correlation between Ra and Ba activities (Figure 43),
salinization would increase Ba and sulfate concentrations, inducing barite precipitation, and
consequently removing 226Ra most efficiently because the short‐lived isotopes are replenished
more rapidly by recoil than is 226Ra.
5.5.6 Radium isotope ratios
Radium isotope ratios, here presented as the ratio of the short‐lived to the long‐lived
isotope (SRa/LRa), are another potential source of useful information on Ra sources and critical
Ra‐mobilizing processes. The 228Ra/226Ra isotope ratio is primarily sensitive to the Th/U ratio of
the aquifer solids that are within recoil distance of the water (Dickson, 1990; Sturchio et al.,
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2001; Vengosh et al., 2009) as long as groundwater is not adjusting to major changes in solid‐
phase radioactivity or water chemistry along its flowpath. The median 228Ra/226Ra of 1.9 is
within the expected range for sand or sandstone aquifers (e.g. Vengosh et al., 2009), but is too
high to represent radium derived from a carbonate aquifer, which may imply that marine
carbonate is not a major radionuclide source to the aquifer if it is assumed that the 228Ra/226Ra
ratio represents the solid‐phase 232Th/238U ratio. However, sulfate concentration is somewhat
correlated with 228Ra/226Ra (ρ= 0.65 or 0.46 when the high‐87Sr/86Sr points are excluded; Figure
44), but not chloride (ρ=0.19 or 0.32 when the high‐87Sr/86Sr points are excluded). This may
imply that 228Ra/226Ra is not completely controlled by the aquifer lithology, but may be higher
than the 232Th/238U of the aquifer rocks because the recoil‐removal balance is not identical
between 228Ra and the longer‐lived 226Ra.
Figure 43: Relationship between Ra2+ thermodynamic activity (derived from 226Ra) and Ba2+ and SO4
2‐ activities.
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During the salinization process, the recoil flux of radium remains constant but the first‐
order rate constant of radium removal, k1, slows, which permits the radium activity in water to
increase. k1 can vary by about 5 orders of magnitude between fresh groundwaters and brines
(Krishnaswami et al., 1982, 1991). Absent another Ra removal process, the difference in decay
constants between 228Ra and 226Ra (half‐lifes of 5.8 yr and 1600 yr respectively) implies that in
response to a change in the system, 228Ra would grow in from alpha recoil 280 times faster than
226Ra, resulting in extremely high 228Ra/226Ra ratios. In contrast, adsorption or coprecipitation
would remove all radium isotopes equally rapidly, which counteracts recoil's tendency to favor
short‐lived isotopes in solution. A related effect of effective Ra removal is that by quickly
removing most of the recoil flux per time unit, the steady‐state radium activity is much lower,
and is reached much more rapidly, than if there were little or no Ra removal. In other words, Ra
removal increases the effective decay constants of Ra isotopes (decreases the half‐lives) by
orders of magnitude. Thus, 228Ra/226Ra ratios respond strongly to large, rapid increases in salinity
on a time span shorter than both 228Ra and 226Ra's decay constants, perhaps months or shorter.
This effect is most dramatic in wells located in beach sands adjacent to the ocean, where high‐
salinity, low‐Ra seawater with 228Ra/226Ra ~1 is pumped into the coastal aquifer and recovered
on short time scales. When pumped, this water exhibits higher Ra levels than seawater and
much higher 228Ra/226Ra ratio (by a factor of up to 10) than both the original seawater and
regional groundwater (Lamontagne et al., 2008; Otero et al., 2011). In this study, the somewhat
elevated 228Ra/226Ra ratios in the highest‐salinity water could signal increased radium
mobilization during salinization. Under this interpretation, elevated 228Ra/226Ra would be caused
by all Ra atoms being removed at the same rate, but 228Ra being replenished more rapidly than
226Ra by alpha recoil. This process could be influenced or complicated by (1) active salinization
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causing salinity increase over time; or (2) pumping‐induced mixing of waters with different
compositions or Ra isotope activities. The potential significance of 228Ra/226Ra for radium
removal processes could be addressed with more certainty by analyzing aquifer rocks for their
radium isotope content.
Another short‐lived to long‐lived activity ratio is 224Ra/228Ra. Due to the effects of
adsorption in accelerating the systems’s response to change, the response of short‐lived 224Ra is
near‐instantaneous. In most groundwater systems observed values are near or slightly above 1
(Krishnaswami et al., 1982; Dickson, 1990; Luo et al., 2000; Focazio et al., 2001; Tricca et al.,
2001), and unusual values >2 are reported in dynamic environments such as intertidal
groundwater (Lamontagne et al., 2008). The observed 224Ra/228Ra values of the waters in this
study are slightly above 1 (median 1.3), which is expected for a system in dominated by alpha
recoil as the primary radium source balanced by rapid adsorption and other geochemical
removal mechanisms such as barite precipitation. Unlike other radium activities and ratios,
224Ra/228Ra is not well correlated with any geochemical parameters in this study.
Figure 44: Relationship between anion concentrations and 228Ra/226Ra activity ratio. Circled points represent three wells with distinct 87Sr/86Sr > 0.709.
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5.6 Conclusions
In summary, the radium isotopes and isotope ratios correlated with chloride
concentration seem to be those that respond most rapidly to the recoil mechanism modified by
diminished adsorption at higher salinity: 224Ra and 228Ra. Those correlated with sulfate
concentration seem to be those most responsive to chemical removal mechanisms – the long‐
lived 226Ra and the 228Ra/226Ra ratio. As with these anion‐traceable effects in which sulfate
concentration is linked to radium removal, major cation behavior is useful for illustrating the
relative importance of adsorption and barite coprecipitation. The cation exchange setting
strongly favoring divalent cation release (Mg, Ca, Sr) implies that Ra adsorption is relatively
inefficient, both by exchange sites rejecting divalent cations due to their surface charge.
Because several geochemical tracers are inconsistent with seawater intrusion being the only
dominant salinity source in the aquifer, the findings of this study may not be uniformly
applicable to all cases of seawater intrusion into coastal aquifers, but are broadly applicable to
the oxic salinization of coastal groundwater. Specifically, the effects of salinity and anion balance
on radium may occur in any aquifer where multiple salinity sources contribute varying amounts
of chloride and sulfate. Inferred barite precipitation, combined with the strong release of
divalent cations from exchange sites in response to high‐Na saline inputs, implies that barite
coprecipitation rather than ion exchange or adsorption is a significant radium removal
mechanism, and perhaps the dominant controlling mechanism of radium removal, in the coastal
aquifer at Agadir. Although it has not been documented that salinization is actively increasing
the abundance of radium in groundwater in the Agadir coastal aquifer, it seems evident that: (1)
higher‐chloride waters permit more radium in solution and future increases in chloride
concentration would also support increased radium levels; and (2) oxic, high‐SO42‐/Cl‐ waters are
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effective at radium removal by barite coprecipitation and the chemical conditions are at least
partially responsible for maintaining overall Ra activities in the fresh and saline waters. The
results of this study suggest that, while increasing salinity is associated with higher radium in
groundwater, sulfate may impart significant radium‐suppressing effects in oxic waters in which
barium is present. Therefore, the relationship between radium and salinity depends on the
major ion composition of the salinity, as well as other well‐documented factors including pH and
redox conditions.
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6. Radium isotope response to salinity gradients in confined Atlantic Coastal Plain aquifers
6.1 Introduction
The geochemistry of radium in aquifer systems is governed by the balance between
radium sources and removal mechanisms. However, unlike other alkaline earth metals and trace
elements of interest in aquifers, both geochemical mechanisms and radioactive decay can
strongly influence observed levels of radium in groundwater. The primary source of radium
isotopes in water is alpha recoil, in which solid‐phase parent thorium isotopes experience alpha
decay at or near the mineral surface and the energy of alpha decay propels Ra isotopes into the
groundwater. In contrast to these radioactive processes of Ra release, Ra removal from water is
much faster than radioactive decay of radium isotopes and is inferred to be due to adsorption in
many systems (Krishnaswami et al., 1982, 1991). The evolution of Ra in saline aquifer systems is
commonly attributed to salinity‐driven effects rather than redox effects on adsorption sites,
although deep saline groundwater systems tend to be anoxic (Kraemer and Reid, 1984; Moise et
al., 2000; Sturchio et al., 2001; Tomita et al., 2010). Less commonly, the concurrent salinity‐
sensitive and redox‐sensitive controls on Ra have been acknowledged in studies of Ra in saline
waters and brines (e.g. Bloch and Key, 1981; Herczeg et al., 1988).
While the salinity‐radium relationship is well established in prior studies, described by
decreasing first‐order rate constant of adsorption k1 (equation 1) as ionic strength increases
(Krishnaswami et al., 1982, 1991), in this research the relative importance of the components of
salinity is examined: anion effects driven by chloride and/or sulfate concentrations, and cation
effects driven by calcium‐ vs. sodium‐dominated waters. One application of this research is that
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studies of Ra and other trace elements along spatial gradients can facilitate prediction of their
response to temporal gradients in salinity, such as seawater intrusion or other salinization
events. In coastal plain aquifers, salinization events can include seawater intrusion or cross‐
formational inputs of natural saline water or brine, in which leakage is induced by pumping that
exceeds natural recharge (Vengosh, 2003). In practical terms, would the expected radium
increase with salinity become problematic for water quality before dissolved solids
concentrations became excessive for use as drinking water? This study focuses on the Atlantic
Coastal Plain, an aquifer system with distinctive cation exchange behavior (Chapelle and Knobel,
1983; Knobel et al., 1998) that provides a natural environment to test the radium source‐
removal balance under the conditions of highly efficient divalent cation removal across a salinity
gradient in which the solubility of Ra would otherwise be expected to increase. Also, the anoxic,
sulfate‐reducing conditions of the Atlantic Coastal Plain (Chapelle and McMahon, 1991) diminish
the importance of a significant complication in studies of Ra geochemistry: the dominance of
redox‐sensitive Mn and Fe oxide adsorption sites in oxic groundwater. In addition to the radium
isotope data, this paper reports 87Sr/86Sr, an isotopic tracer of divalent cation sources; boron
isotopes, which may record the cation exchange conditions favoring Na‐dominated waters with
low divalent cation concentrations; and S and O isotopes of sulfate to evaluate sulfate reduction
and its effects on barite precipitation, which can remove radium. Used in combination with ion
ratios, these isotopic tracers allow the relative importance of critical Ra‐mobilizing processes to
be identified.
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6.2 The Atlantic Coastal Plain aquifer system: Natural and increasing salinity
In the Cretaceous to Holocene sediments that constitute the Atlantic Coastal Plain
aquifer system fresh waters at the recharge area are modified down‐flow by cation exchange,
microbial redox processes, and mixing with saline waters. This gradient provides a natural
opportunity to evaluate the response of radium isotopes to salinity gradients under uniform
anoxic conditions. This study focuses on the deepest and shallowest fresh confined aquifers in
the stratigraphy of the North Carolina coastal plain (Figure 45): (1) the upper Cretaceous Cape
Fear formation, which overlies the crystalline basement in the study area. The Cape Fear
formation, interpretated as deltaic deposits including offshore components (Gohn, 1988;
Winner and Coble, 1996), contains a major regional confining unit in its middle such that
hydraulically distinct upper and lower aquifers are defined in North Carolina. While lithologically
similar to the Upper Cape Fear aquifer, the Lower Cape Fear aquifer exhibits higher salinity
thought most of its extent in North Carolina (Winner and Coble, 1996). For simplicity, the Upper
and Lower Cape Fear aquifers are referred to collectively as the Cape Fear aquifer system in this
paper; and (2) the Pliocene Yorktown aquifer, which is the youngest major confined aquifer in
the North Carolina coastal plain and is composed of marine sand and clay containing abundant
shell material (Winner and Coble, 1996), and is locally phosphatic in the lower part of the
formation (Riggs and Belknap, 1988). Both aquifers have large subsurface extent and become
more saline down‐gradient towards the Atlantic Ocean, reaching salinity levels that would
require treatment for use (Figure 46).
In North Carolina, the Coastal Plain aquifers are extensively pumped for public water
supply and industrial use. Pre‐development hydraulic heads were commonly above the land
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surface, but for several decades water levels in wells have been declining, especially near cities
(dePaul et al., 2008) and this is specifically the case in the wells sampled for this study, where
declines of up to 10 m have been observed since 1980 (Figure 47). In most areas, the natural
condition of the aquifer system includes higher heads with increasing well depth (Winner and
Coble, 1996), including the nested monitoring wells sampled in this study. This pressure
relationship allows deep, saline water to leak upwards through confining units in response to
high pumping rates of the fresh, shallow waters so that declining water levels in wells are
associated with increasing chloride concentrations (Vinson et al., in press).
The Atlantic Coastal Plain has been the site of numerous studies on the geochemistry of
confined aquifers, including detailed examination of cation exchange (Back, 1966; Chapelle and
Knobel, 1983) and organic carbon oxidation (Chapelle and Knobel, 1985; Chapelle et al., 1987;
Chapelle and McMahon, 1991; McMahon and Chapelle, 1991a,b; Haque et al., 2008). Briefly,
conditions proceed from acidic, low‐TDS groundwaters at the recharge area to anoxic, near‐
neutral Ca‐Mg‐HCO3‐ waters in the shallow aquifer, to Na‐HCO3
‐ fresh waters farther down‐
gradient subsequent to cation exchange, and eventually reaching a mixing zone with saline Na‐
Cl waters (Chapelle and Knobel, 1983; Knobel et al., 1998). Reduction of iron and sulfate
electron acceptors is widespread in the confined aquifer, although methanogenesis is mostly
inhibited by high sulfate concentrations (Chapelle and McMahon, 1991). A few previous studies
have also focused on the radium isotope geochemistry of the Atlantic Coastal Plain aquifers.
Much attention has been devoted to the far up‐dip extent of the Coastal Plain sediments where
dilute, acidic waters exhibit Ra in excess of drinking water standards due to the low‐pH (<4.5)
effects on radium adsorption (Michel and Moore, 1980; Bolton, 2000; Szabo et al., 2005). In the
confined, near‐neutral Coastal Plain aquifers, radium was evaluated in a phosphate mining area
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of North Carolina, given the affinity of marine phosphates to accumulate U. However, low levels
of 226Ra were documented (Mitsch et al., 1984).
Figure 45: Schematic cross section of the Atlantic Coastal Plain aquifer system in North Carolina. Dashed polygons represent the approximate extent of sampling in the Cape Fear and
Yorktown aquifers (modified from Winner and Coble (1996), plate 7).
6.3 Methods
6.3.1 Water Sampling
The Cape Fear aquifer water samples plus Yorktown aquifer samples J3O3, J7K5, and
J7K7 were collected from wells in the monitoring network of the North Carolina Division of
Water Resources (Figure 46). These monitoring wells are assigned to the Yorktown, Upper Cape
Fear, or Lower Cape Fear aquifers based on hydrogeologic criteria and have short screen lengths
(typically 3 m). Wells were sampled after purging three volumes, although low‐yielding (M25F6)
or deep (1‐510 and 1‐610) wells were sampled after approximately one well volume. One saline
well, I13X3, could not be purged but was sampled for Br‐, Cl‐, and 87Sr/86Sr, which should be
unaffected by the stagnant conditions. Four wells, M25F4, M25F6, J22P7, and O23L8, yielded
high concentrations of suspended particles resulting in rapid clogging during filtration. The
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Figure 46: Map showing the extent of the Yorktown, Upper Cape Fear, and Lower Cape Fear aquifers and well locations (data from North Carolina Division of Water Resources shapefiles).
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Figure 47: Historic water levels in Cape Fear aquifer wells sampled at this study and selected other Cape Fear wells at the same well nests. Data provided by North Carolina Division of
Water Resources (http://www.ncwater.org).
remaining Yorktown Aquifer samples were collected from monitoring wells and high‐capacity
water supply wells maintained by the Dare County Water Department in the community of Kill
Devil Hills (Figure 46). These include a vertical nest containing a shallow monitoring well (1‐310),
production well (PW1), and two deep monitoring wells (1‐510 and 1‐610).
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6.3.2 Major and trace elements
Na, Ca, and Mg concentrations were determined using direct current plasma (DCP)
spectrometry and potassium analyzed by flame atomic absorption spectrometry, both calibrated
using standards prepared from plasma‐grade single‐element standards. B, Ba, and Sr
concentrations were determined on appropriately diluted water samples by inductively‐coupled
plasma mass spectrometry, calibrated using NIST 1643e trace element solution. Major anions
(Cl‐, SO42‐, Br‐) were analyzed by ion chromatography, and bicarbonate concentrations were
determined by titration of unfiltered water samples to pH 4.5.
6.3.3 Environmental isotopes (δ11B, δ34SSO4, δ18OSO4,
87Sr/86Sr)
Filtered water samples were treated with H2O2 several days prior to loading to facilitate
oxidation of organic carbon in a laminar flow hood with B‐free filters, then loaded onto
degassed Re filaments in a B‐free synthetic seawater solution (Dwyer and Vengosh, 2008).
Boron isotope ratios were measured by negative ion thermal ionization mass spectrometry
(TIMS) at Duke University. After stabilizing BO2 signal (mass 43) in negative ion mode at ~1V, the
CN signal at mass 42 was monitored (<5000 cps) using the secondary electron multiplier to
evaluate potential interference (Hemming and Hanson, 1992). 11B/10B ratios were normalized to
replicate runs of NIST 951 and reported as δ11B. Long‐term replication of NIST 951 indicates
precision of 0.25‰ (1σ, n=238). Sulfur and oxygen isotopes of sulfate were determined by
isotope ratio mass spectrometry (IRMS) at the University of Calgary subsequent to BaSO4
precipitation of filtered water samples. For oxygen isotope analysis, the purity of BaSO4
precipitate was checked by weighing, and only those precipitates containing 27 ± 2% oxygen are
reported. Strontium isotope samples were prepared by drying ~ 3 µg of Sr, with subsequent
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HNO3 digestion and extraction with Eichrom SR‐B50‐S resin. The resulting separated materials
and a Ta load solution were loaded onto Re filaments, then analyzed by positive ion TIMS at
Duke University. The long‐term average value of NIST 987 is 0.710245 ± 0.000010 (1σ, n=109).
6.3.4 Radium isotopes
Unfiltered water samples were collected in polyethylene containers and processed later
the same day. Most samples were 52 L, but a few samples were larger or smaller (26‐78 L).
Samples were slowly pumped through sequential columns (typically two columns), each
containing 10 g (dry weight) of Mn oxide impregnated acrylic fibers (Moore and Reid, 1973). The
fibers were rinsed with deionized water to remove particles and salts, then squeezed of excess
water. Fiber samples from each separate column were analyzed 1‐3 days later for 224Ra using a
delayed coincidence counter (Moore and Arnold, 1996; Garcia‐Solsona et al., 2008) calibrated at
efficiency appropriate for hand‐squeezed fibers. Subsequent sample runs were used to correct
the 224Ra results for the count rate attributable to 223Ra (Moore and Arnold, 1996) and confirm
negligible 224Ra support by its parent 228Th. 226Ra was analyzed after a 21‐day incubation period
in an evacuated glass tube by analyzing the ingrown 222Rn on a Durridge RAD7 radon‐in‐air
monitor (Kim et al., 2001). 228Ra was analyzed on a compressed wafer (5 mm tall and 65 mm in
diameter) prepared from the first column’s fibers on a Canberra broad‐energy germanium
gamma detector using a weighted the 911 keV peak of 228Ac. This was then corrected for the
distribution of Ra between the first and subsequent columns using the 224Ra results. The stated
uncertainties of radium isotope results are based on counting statistics only, which include the
propagation of error (Eaton et al., 2005) associated with background subtraction and the
addition of activities from the individual columns of fibers. These calculated counting errors do
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not incorporate other potential uncertainties including moisture content variability (224Ra),
counting geometry variation (228Ra), or quantitative adsorption onto the Mn oxide. Sample‐
specific radium detection limits are estimated by the 2σ error. This detection limit is defined as
the radium activity with a 95% probability of detection at ±100% precision and is equivalent to
1.96σ (Eaton et al., 2005).
6.3.5 Speciation calculations
Speciation calculations were performed using the geochemical code PHREEQC version 2
(Parkhurst and Appelo, 1999). Additional stability constants were added to the default database
for Ra2+, RaSO40, RaCl+, and RaHCO3
‐ (Langmuir and Riese, 1985), which are similar to the
thermodynamic constants in widespread use for Ba.
6.4 Results
6.4.1 Major and trace element chemistry
Down‐gradient within the aquifer, total dissolved solids (TDS) increases gradually and
several changes in the major cation composition are observed: increasing dominance of the
cations by Na, increasing Cl‐ concentrations, and irregular but increasing sulfate concentrations
(Table 21). The same trends in TDS are seen vertically in nested wells: waters exhibit more Na‐
HCO3‐ and Na‐Cl‐ character with depth. Among the wells that were sampled for full chemistry,
sodium and chloride exhibit the largest ranges in concentrations ‐ two to three orders of
magnitude (28‐5820 mg L‐1 and 5‐9890 mg L‐1, respectively; Table 22). These large ranges in
cation and anion compositions result in distinctive ion ratios between the aquifers and relative
to seawater. In both aquifers, Mg is significantly lower in concentration than the expected
Mg/Cl‐ ratio from seawater. In the Yorktown aquifer, Ca/Cl‐ and Sr/Cl‐ are at or above seawater
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values; in the Cape Fear aquifer, Ca/Cl‐ and Sr/Cl‐ are at or below seawater values. Overall, the
Mg/Cl‐, Ca/Cl‐, Sr/Cl‐ ratios are higher in the Yorktown aquifer than in the Cape Fear aquifer
(Figure 48). Anion concentrations also exhibit significant differences between the two aquifers
sampled. The ratio of sulfate to chloride is consistently higher than seawater values in the Cape
Fear aquifer but lower than seawater values in the Yorktown aquifer (Figure 49). Throughout the
waters sampled in this study, Br‐/Cl‐ ratios are uniformly close to the expected seawater value of
1.5 x 10‐3 (Figure 50).
Sodium is the dominant cation in nearly every water sample (median 93% Na+K (as
equivalents), range 27‐98%). Consistent with the patterns observed in divalent cations (Figure
48), higher sodium dominance was observed in the in the Cape Fear aquifer (median 97% Na+K)
than in the Yorktown aquifer (median 85% Na+K). On average, sodium dominance is significantly
higher than seawater (78% Na+K). The two shallowest wells in each aquifer relative to modern
sea level (the four shallowest wells overall) are also the two least Na‐dominated in each aquifer.
In the Yorktown aquifer, these are wells J7K7 and J3O3 (depths 23‐64 m below sea level; 27‐52%
Na+K, compared to 81‐97% Na+K in the remaining wells). In the Cape Fear aquifer, the
shallowest two wells are M25F4 and J22P1 (depths 58‐65 m below sea level; 56‐89% Na+K,
compared to 91‐98% Na+K in the remaining wells). These four shallowest wells exhibit higher
Mg/Cl‐, Ca/Cl‐, Sr/Cl‐ (Figure 48), Ba/Cl‐, and Ra/Cl‐ ratios (Figure 51) than the deeper, more Na‐
dominated waters.
6.4.2 Strontium and boron
Significant differences in strontium concentrations were observed between the two
aquifers. As indicated in Figure 48, the groundwater Sr/Cl‐ ratio is either higher than (median
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4.9 x 10‐4 in the Yorktown aquifer) or slightly below (median 1.1 x 10‐4 in the Cape Fear aquifer)
the seawater Sr/Cl‐ ratio of 1.6 x 10‐4. In addition to the order‐of‐magnitude difference in Sr/Cl‐,
the Sr/Ca ratio varies between the two aquifers (median 0.025 and 0.0016 in the Yorktown and
Cape Fear aquifers, respectively). Strontium isotope ratios differ also significantly between the
two aquifers studied. In the Yorktown aquifer, 87Sr/86Sr falls within a narrow range (0.70898‐
0.70916, median 0.70904), but ratios occupy a larger range in the Cape Fear aquifer and are
generally lower (0.70813‐0.70895, median 0.70871; Figure 52; Table 22). There is no observed
relationship between 87Sr/86Sr and Sr concentration in either aquifer.
Both aquifers exhibit a large range of boron concentrations (0.11‐2.60 mg L‐1 and
0.36‐6.61 mg L‐1 in the Yorktown and Cape Fear aquifers respectively; Table 22), including values
that exceed water quality guidelines of the World Health Organization (0.5 mg L‐1; World Health
Organization, 2008) and the U.S. Environmental Protection Agency (5 mg L‐1 lifetime Health
Advisory; Environmental Protection Agency, 2008). In terms of δ11B, the Yorktown aquifer
exhibits about 15 ‰ of variability (10.9‐34.7 ‰), although the Cape Fear aquifer exhibits a wider
range extending both above and below the global seawater value (8.5‐51.8 ‰; Figure 52). Low
values of δ11B are associated with fresh waters containing low boron concentrations but high
B/Cl‐ ratios. High values of δ11B are associated with saline waters exhibiting high boron
concentrations and somewhat lower B/Cl‐ ratios (Figure 53). It was also observed that the two
aquifers exhibit parallel trends with respect to the relationship between between B/Cl‐ and δ11B,
but an essentially overlapping relationship between δ11B and well depth. Relationships were
also seen in the Na/Cl‐ ratio in which an increasing Na/Cl‐ ratio is associated with increasing B/Cl‐
ratio in both aquifers, in addition to trends showing the highest Na/Cl‐ ratios in each aquifer are
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in waters of lowest δ11B. Overall, at comparable values of B/Cl‐ and Na/Cl‐ ratios, the Cape Fear
aquifer groundwater exhibits higher values of δ11B than the Yorktown aquifer (Figure 53).
Figure 48: Ca, Mg, and Sr plotted relative to chloride, grouped by aquifer and water type. Dashed lines indicate seawater ion ratios. For clarity, the four shallowest samples are
excluded from plots in the left two columns.
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Figure 49: Sulfate concentration plotted relative to chloride concentration, grouped by aquifer. Dashed line represents the seawater ratio.
Figure 50: Plot of bromide vs. chloride concentrations in Cape Fear and Yorktown aquifer groundwater. Dashed line represents seawater ratio.
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Figure 51: Ba and 226Ra plotted relative to chloride, grouped by aquifer and water type. Dashed line indicates seawater ratio. For clarity, the four shallowest wells are excluded from
plots in the left two columns.
Figure 52: Environmental isotope ratios (strontium, boron, and sulfur) in Cape Fear vs. Yorktown aquifer groundwater. Dashed lines represent modern seawater values.
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Figure 53: Plots showing boron concentrations and isotope ratios.
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6.4.3 Sulfur and oxygen isotopes of sulfate
The aquifers contain similar ranges of sulfate concentrations (0.2‐1140 mg L‐1 overall),
but ion ratios and isotopic composition of sulfate differ between the samples collected in the
two aquifers. The Yorktown aquifer exhibits SO42‐/Cl‐ ratios lower than the seawater value,
whereas the Cape Fear aquifer generally exhibits higher SO42‐/Cl‐ than seawater (Figure 49). All
values of δ34SSO4 and δ18OSO4 exceed the seawater values of 21.0 ‰ and 9.5 ‰ respectively
(Table 22) and are generally higher in the Yorktown aquifer than in the Cape Fear aquifer (Figure
52). The six samples for which both δ34SSO4 and δ18OSO4 are available plot along a line with slope
~0.25 (Figure 54).
Figure 54: δ34SSO4 plotted against δ18OSO4 for Yorktown and Cape Fear aquifer groundwater.
6.4.4 Barium and Radium
While barium concentrations increase somewhat with salinity (plotted as chloride
concentration) in the Na‐dominated waters of the Yorktown aquifer, there is no apparent Ba‐
salinity trend in the Cape Fear aquifer (Figure 51). Activities of 224Ra, 226Ra, and 228Ra increase
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Table 21: Major element concentrations (in mg L‐1) of the Yorktown and Cape Fear aquifers. Data in italics appeared in Vinson et al. (in press).
Well Screen elev (m) Date T (°C) pH O2 Ca Mg Na K Cl‐ Br‐ SO42‐ HCO3
‐ Yorktown aquifer J3O3 ‐64 10 Dec 2008 17.5 7.2 0.3 49.5 4.2 67.0 4.9 49.8 0.2 0.2 265 J7K7 ‐23 11 Dec 2008 17.8 6.8 54.0 23.0 27.7 18.8 22.1 0.1 0.4 356 J7K5 ‐94 11 Dec 2008 18.5 7.9 0.3 17.9 29.7 1080 42.4 1340 4.9 64.2 586 Ocean Bay ~ ‐120 28 Oct 2008 8.6 0.3 4.6 6.0 511 4.5 336 1.3 27.2 717 Albatross ~ ‐120 28 Oct 2008 18.8 8.0 0.3 20.6 42.7 950 3.7 1150 4.1 21.5 515 PW 1‐310 ‐95 28 Oct 2008 8.0 0.7 4.8 6.1 469 1.5 357 1.3 24.2 574 Well 1 ~ ‐90‐120 28 Oct 2008 20.0 7.9 0.6 63.2 112 1570 82.9 2400 8.7 168 354 PW 1‐510 ‐156 28 Oct 2008 21.4 7.2 0.2 232 479 4990 168 8180 29.6 764 333 PW 1‐610 ‐187 28 Oct 2008 21.4 7.2 0.1 270 527 5820 207 9890 36.4 719 458 Well 2 ~ ‐90‐120 28 Oct 2008 20.0 7.8 1.6 72.2 128 1650 88.0 2520 9.1 191 371 Well 7 ~ ‐90‐120 28 Oct 2008 19.5 7.8 0.6 79.1 128 1640 74.1 2610 9.3 177 313 Well 11 ~ ‐90‐120 28 Oct 2008 7.8 0.7 66.5 122 1720 74.4 2740 10.1 187 410 Upper & Lower Cape Fear aquifers L24B5 ‐91 18 Oct 2007 18.5 7.6 4.7 7.0 658 29.0 379 1.3 286 334 M25F4 ‐58 19 Oct 2007 19.5 7.8 5.5 2.0 80.9 9.3 8.8 0.0 7.7 204 M25F5 ‐72 19 Oct 2007 19.1 8.1 0.4 1.3 0.9 133 10.6 62.3 0.2 26.1 213 M25F6 ‐145 19 Oct 2007 21.8 8.1 12.4 11.1 928 31.9 457 1.6 304 725 J22P6 ‐160 22 Oct 2007 20.4 8.0 0.3 17.1 20.2 1100 39.4 774 2.7 459 729 J22P7 ‐131 22 Oct 2007 20.0 7.7 0.2 5.6 6.3 530 24.8 402 1.5 232 567 J22P1 ‐65 22 Oct 2007 18.7 7.3 0.3 18.9 18.4 60.0 22.5 4.6 0.4 302 K21R5 ‐266 22 Oct 2007 22.4 7.4 0.2 95.0 132 3550 114 4680 16.6 1090 531 K21R6 ‐207 10 Dec 2008 20.8 7.5 0.4 53.9 81.7 2720 85.1 3040 10.8 1140 633 K21R7 ‐121 10 Dec 2008 19.1 8.5 0.3 3.9 1.4 258 12.7 79.7 0.3 58.0 467 P24O1 ‐209 11 Dec 2008 20.7 7.6 0.2 4.8 3.6 647 18.3 369 1.4 317 622 I13X3 ‐270 11 Dec 2008 16000 58.3 R23X10 ‐295 13 Jan 2009 22.5 7.6 0.8 25.6 20.9 1810 47.7 2450 10.9 305 390 S22J12 ‐305 13 Jan 2009 22.0 7.2 0.4 19.2 14.0 1600 45.5 2040 8.9 115 647 O23L8 ‐253 13 Jan 2009 20.5 7.7 0.8 8.9 8.0 926 26.2 825 3.2 252 762 Seawater (Millero and Sohn, 1991) 8.2 412 1280 10800 399 19400 67 2710 114
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Table 22: Trace metal concentrations and isotope ratios of Yorktown and Cape Fear aquifer groundwater. Concentrations are in mg L‐1. Data in italics appeared in Vinson et al. (in press).
Well δ34SSO4 (‰) δ18OSO4 (‰) Sr (µg L‐1) 87Sr/86Sr B (mg L‐1) B/Cl‐ δ11B (‰) Ba (µg L‐1) 224Ra (mBq L‐1) 226Ra (mBq L‐1) 228Ra (mBq L‐1) Yorktown aquifer J3O3 275 0.709058 0.11 0.0073 8.6 4.8 ± 0.7 3.3 ± 0.7 6.0 ± 3.2 J7K7 368 0.709156 0.27 0.0396 20.9 15.9 6.4 ± 0.9 5.6 ± 0.8 4.6 ± 2.6 J7K5 701 0.708997 1.19 0.0029 33.4 6.9 7.0 ± 0.9 4.6 ± 1.7 2.8 ± 2.0 Ocean Bay 103 0.709148 1.85 0.0181 24.8 4.0 1.7 ± 0.4 2.5 ± 0.5 0.2 ± 0.9 Albatross 902 0.709052 1.42 0.0041 29.8 2.4 4.9 ± 0.8 4.9 ± 0.7 3.8 ± 1.3 PW 1‐310 88.9 0.709000 1.33 0.0123 22.6 1.3 1.9 ± 0.5 1.9 ± 0.6 1.1 ± 1.2 Well 1 34.4 21.1 2830 0.709052 1.28 0.0018 25.8 8.0 21.4 ± 0.9 19.1 ± 1.3 19.5 ± 3.4 PW 1‐510 34.7 21.8 13800 0.709028 2.19 0.0009 34.5 50.9 99.1 ± 7.3 186 ± 7 96.7 ± 21.7 PW 1‐610 36.5 22.4 18000 0.709028 2.60 0.0009 34.7 67.2 98.7 ± 6.1 214 ± 7 119 ± 26 Well 2 2970 0.709040 1.31 0.0017 27.9 8.8 19.1 ± 1.0 20.0 ± 1.3 24.0 ± 4.6 Well 7 3450 0.709038 1.26 0.0016 27.3 10.9 25.6 ± 1.7 26.9 ± 1.5 27.0 ± 6.3 Well 11 3440 0.709027 1.43 0.0017 28.8 9.3 29.2 ± 2.4 26.9 ± 1.7 26.3 ± 6.8 Upper & Lower Cape Fear aquifers L24B5 46.9 0.708709 2.60 0.0226 16.3 23.5 ± 1.6 9.4 ± 1.2 15.4 ± 4.4 M25F4 19.6 0.708683 0.67 0.2486 15.7 25.3 29.8 ± 1.8 21.5 ± 1.4 30.8 ± 9.1 M25F5 8.0 0.708877 1.25 0.0660 28.7 3.6 ± 0.7 4.4 ± 0.9 1.9 ± 2.9 M25F6 78.3 0.708881 3.03 0.0217 43.0 58.7 39.8 ± 2.9 28.0 ± 1.9 55.1 ± 15.9 J22P6 29.4 20.7 169 0.708506 3.83 0.0162 46.9 29.8 18.4 ± 1.2 10.6 ± 1.3 17.9 ± 4.9 J22P7 58.0 0.708641 3.61 0.0296 42.2 28.1 23.6 ± 1.9 39.5 ± 2.1 11.5 ± 6.0 J22P1 188 0.708945 0.36 0.2550 8.5 704 42.7 ± 2.0 18.6 ± 1.4 51.4 ± 14.1 K21R5 2060 0.709065 4.62 0.0032 51.8 27.7 257 ± 14 105 ± 5 175 ± 30 K21R6 29.0 1190 0.708922 4.26 0.0046 45.1 20.6 74.9 ± 3.4 32.0 ± 1.4 60.8 ± 5.6 K21R7 32.5 0.708673 2.19 0.0901 27.5 41.5 8.1 ± 1.0 4.3 ± 0.6 5.7 ± 1.3 P24O1 31.2 21.1 63.3 0.708943 4.81 0.0427 42.0 25.5 6.2 ± 0.8 14.1 ± 1.2 6.2 ± 1.5 I13X3 0.708173 R23X10 773 0.708127 5.92 0.0079 45.9 31.8 28.4 ± 2.2 26.5 ± 2.0 18.5 ± 2.3 S22J12 30.6 20.1 816 0.708084 6.61 0.0106 45.2 58.7 66.1 ± 3.2 30.0 ± 1.9 53.9 ± 4.7 O23L8 166 0.708994 5.46 0.0217 47.8 39.1 17.1 ± 1.1 14.8 ± 1.3 12.6 ± 2.0 Seawater 21.0 1 9.5 1 7900 2 0.70924 1 4.6 3 0.0008 39 3 14 2 1‐ Clark and Fritz, 1997 2‐ Millero and Sohn, 1991 3‐ Vengosh and Spivack, 1999
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with chloride in both aquifers studied (Figure 51, Figure 56), and the increase is somewhat linear
above about 1000 mg L‐1 chloride. Radium activities exceed US EPA drinking water standards
(185 mBq L‐1 combined 226Ra+228Ra activity) only in the highest‐salinity waters. The shallowest
two samples in each aquifer exhibit significantly higher Ba/Cl‐ and Ra/Cl‐ ratios than do the Na‐
dominated waters (Figure 51). The highest Ba concentration in the study was observed in one of
the shallow waters from the Cape Fear aquifer (704 µg L‐1; Table 22).
6.5 Discussion
6.5.1 Major element chemistry and the aquifers’ response to salinization
The major element chemistry observed in the Yorktown and Cape Fear aquifers is
consistent with the vertical and down‐gradient trends described in previous investigations, in
which a distinctive combination of ion exchange and mixing phenomena are observed: (1)
Recharging waters interact with marine carbonate material, yielding a fresh Ca‐Mg‐HCO3‐
groundwater; (2) Farther down‐gradient, divalent‐monovalent exchange in which divalent
cations are removed and Na released,, increases the proportion of Na in the groundwater and
yields a Na‐HCO3‐ water; (3) CO2 from microbial respiration drives aragonitic shell dissolution,
with subsequent Ca removal by ion exchange, yielding increased TDS within the Na‐HCO3‐
hydrochemical facies; and (4) Na‐HCO3‐ waters mix with saline Na‐Cl‐ waters at depth that
represent modified seawater from prior sea level fluctuations (Chapelle and Knobel, 1983, 1985;
Chapelle et al., 1987; Meisler, 1989; McMahon and Chapelle, 1991a; Knobel et al., 1998). The
cation ratios are consistent with this ion exchange sequence yielding Na‐HCO3‐ and Na‐Cl‐
waters; Sr/Cl‐ and Ca/Cl‐ values above the seawater value in the Yorktown aquifer are consistent
with aragonitic shell dissolution (step 3 above) that somewhat modifies the Na‐dominated
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nature of these aquifers. The freshening process yielding Na‐HCO3‐ waters (step 2 above)
represents the progressive reversal of very old seawater intrusion into the Coastal Plain
aquifers. However, the long residence time of fresh to brackish groundwater in confined Coastal
Plain aquifers (104‐105 years; Hathaway et al., 1979; Murphy et al., 1992; Purdy et al., 1992,
1996; Plummer, 1993; Kennedy and Genereux, 2007) indicates that progressive freshening along
the ~200 km‐long aquifer is extremely slow, at least in North Carolina where saline waters are
found far inland in all of the confined aquifers (Manheim and Horn, 1968; Meisler, 1989; Knobel
et al., 1998). In the wells sampled in the study, this saline Na‐Cl water exhibits marine Br‐/Cl‐
ratios indicating its origin as seawater (Figure 50). In the brackish to saline sections of the
aquifer system, exchange sites remain Na‐saturated from interaction of marine clays with
seawater (Back, 1966; Chapelle and Knobel, 1983). Conversely, seawater intrusion of previously
freshened aquifers can release divalent cations by ion exchange (Nadler et al., 1980; Mercado,
1985; Appelo and Postma, 1993; Vengosh, 2003; Sivan et al., 2005; Russak and Sivan, 2010).
Thus, the relationship between cation composition and salinity is the opposite of what would be
expected from intrusion of modern seawater. The relationship between sodium domination and
the Mg/Cl‐, Ca/Cl‐, Sr/Cl‐, Ba/Cl‐, and Ra/Cl‐ ratios is consistent throughout the samples analyzed
from both aquifers: the four shallowest wells (less sodium‐dominated waters) exhibit higher
ratios of all five alkaline earth metals relative to chloride (Figure 48; Figure 51) than the more
Na‐dominated samples. Because these alkaline earth metals come from different geologic
sources (e.g. Sr is concentrated in marine sources and Ba is much higher in continental sources),
the similarity in their behavior strongly suggests that they are affected by a shared removal
mechanism related to their divalent nature.
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The aquifer’s response to the salinization of heavily pumped wells in the Yorktown and
Cape Fear aquifers (Figure 47) is analogous to the mixing relationship between Na‐HCO3‐ and
Na‐Cl‐ waters that occurs at depth throughout the Atlantic Coastal Plain (Knobel et al., 1998).
Because of the negligible differences in the cation composition between these waters, the
increasing salinity (chloride concentration) of the heavily pumped areas of the Coastal Plain
aquifers does not appear to induce cation exchange. The cation exchange phenomena observed
in the study area are a major influence on water chemistry, but they correspond with the
natural, gradual, down‐gradient freshening processes that dominate the confined Coastal Plain
aquifers, rather than pumping‐induced processes.
6.5.2 Redox conditions
Although radium is not redox‐sensitive, solid‐phase removal mechanisms for radium,
Mn oxides and barite, are sensitive to reductive dissolution reactions linked to natural organic
carbon oxidation by bacteria in groundwater systems (Bolze et al., 1974; Stumm and Morgan,
1996). Evidence for anoxic, reducing conditions is seen the in the low dissolved oxygen
concentrations (Table 17) and sulfur isotope data. In both aquifers, S and O isotope ratios
(Figure 54) are enriched over expected sulfate sources to this system such as oxidized marine
sulfide (Chapelle and McMahon, 1991) with its expected negative δ34S values, and seawater
(δ34S = 21‰), which implies that sulfate in the groundwater is residual from extensive sulfate
reduction (Krouse and Mayer, 1999). As further evidence of sulfate reduction in the Yorktown
and Cape Fear aquifers, the slope of 0.25 between δ18OSO4 and δ34SSO4 (Figure 54) is the expected
slope for kinetic fractionation effects during sulfate reduction from one abundant sulfate source
(Claypool et al., 1980; Tuttle et al., 2009). Yet, paradoxically, the Cape Fear aquifer system
190
exhibits SO42‐/Cl‐ ratios higher than the seawater value (Figure 49) despite its strong evidence of
sulfate reduction. Thus, the aquifer is receiving a supply of sulfate at lower δ18OSO4 and δ34SSO4
than observed in the groundwater to prevent the exhaustion of sulfate by sulfate reduction. This
source was documented in Cretaceous Coastal Plain aquifers as diffusion of sulfate from
confining beds. This sulfate diffuses into the sandy aquifer sediments and maintains high sulfate
concentrations concurrent with active sulfate reduction (Chapelle and McMahon, 1991; Brown
and Schoonen, 2004). Along with organic acids also diffused from the confining units (McMahon
and Chapelle, 1991b), these sulfate inputs support sulfate reduction as the dominant organic
carbon oxidation reaction in the aquifers, rather than allowing methanogenesis to dominate
after the depletion of sulfate (McMahon and Chapelle, 1991a). Assuming that the SO42‐/Cl‐ ratio
somewhat indicates a higher degree of confining unit interactions, it implies that the Cape Fear
aquifer has more clay content than the Yorktown aquifer, consistent with other lines of evidence
such as higher Na/Cl‐ and lower Ca/Na ratios in the Cape Fear. Overall in both aquifers, these
sulfate‐reducing conditions are unfavorable for extensive Ra adsorption sites to be present as
highly reactive Mn and Fe oxides of high surface area, which, if present could be effective
adsorption sites for Ra. This is because Fe and Mn oxide reduction is more thermodynamically
favorable than sulfate reduction (Champ et al., 1979; Stumm and Morgan, 1996).
6.5.3 Strontium
If exchangeable Sr indicates different 87Sr/86Sr from modern seawater, then the release
of divalent cations from aquifer solids by saline influxes may be indicated by 87Sr/86Sr of salinized
groundwater (Vengosh et al., 2002; Bouchaou et al., 2008). In the Cape Fear aquifer, 87Sr/86Sr
ratios (0.70808‐0.70906) are significantly more radiogenic (Figure 55) than those of Cretaceous
191
seawater (0.7072‐0.7078; McArthur et al., 2001). However, 87Sr/86Sr is not correlated with any
indicators of salinity such as chloride concentration, providing no indication that salinity
increases in the aquifer correspond with Sr release, nor should any correlation be expected
because the salinization of the aquifer is defined by mixing of Na‐HCO3‐ with Na‐Cl‐ waters
without inducing cation exchange. This irregular distribution in upper and lower Cape Fear
waters was also seen in four wells by Sirtariotis (1998) in southeastern North Carolina (range
0.70799‐0.70892). The ratios observed in Cape Fear aquifer groundwater, being more radiogenic
than seawater, are derived from continental‐sourced Sr such as mica (Brown, 1959) and feldspar
(Gohn, 1988) present in the formation. In the Yorktown aquifer, 87Sr/86Sr falls within a narrow
range (0.70900‐0.70916) that corresponds to late Miocene to late Pliocene marine carbonate
(Figure 55; McArthur et al., 2001). These results are consistent with other studies from the
North Carolina coastal plain indicating agreement with the global seawater 87Sr/86Sr record in
Tertiary formations (Denison et al., 1993) and their groundwaters (Woods et al., 2000) in which
seawater 87Sr/86Sr has increased since the late Eocene epoch. Because these ratios are observed
in both fresh and saline Yorktown waters, including the recharge area, these values indicate that
87Sr/86Sr is governed by equilibration with the Pliocene carbonate of the aquifer rather than
modern seawater intrusion, probably by dissolution of aragonitic shell material given the high Sr
concentrations and Sr/Cl‐ ratio in the Yorktown aquifer and the high Sr content of aragonite
(Kinsman, 1969). Overall, the cation chemistry of the Yorktown aquifer is complicated by the
release of Ca, Mg, and Sr from carbonates and subsequent cation exchange removing divalent
cations and releasing Na. The highly efficient removal of alkaline earth metals throughout both
aquifers is not recorded by 87Sr/86Sr, but to some degree 87Sr/86Sr provides a record of marine vs.
continental divalent cation sources.
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Figure 55: Relationship between 87Sr/86Sr and 228Ra/226Ra. Labeled bands represent the 87Sr/86Sr ratios of Cretaceous and Pliocene marine carbonate (McArthur et al., 2001). Note that most Yorktown aquifer samples exhibit Pliocene marine 87Sr/86Sr and carbonate‐like median 228Ra/226Ra < 1; most Cape Fear samples exhibit 87Sr/86Sr modified by radiogenic continental
inputs and sandstone‐like median 228Ra/226Ra > 1.
6.5.4 Boron
The 87Sr/86Sr results of this study illustrate the one‐directional nature of strontium
isotopes as a tracer of exchange processes: Waters that gain Sr by cation exchange or
dissolution record the Sr source by their 87Sr/86Sr ratio, yet waters losing Sr to the sediments do
not record this process in the 87Sr/86Sr of residual groundwater Sr. Boron exhibits potential as a
tracer of this cation exchange scenario because of its high concentration in seawater and its
association with high‐Na fresh waters (e.g. Ravenscroft and McArthur, 2004; Faye et al., 2005;
Halim et al., 2010). Desorption of boron from sediments may be driven by two independent
mechanisms: (1) decreasing pH, during which the attraction of adsorption sites for B(OH)4‐
declines (Palmer et al., 1987; Spivack et al., 1987; Singh and Mattigod, 1992), and (2) dilution, in
which decreasing ionic strength stimulates desorption without any change in pH (Goldberg et
193
al., 2008). At near‐neutral pH, B is present as B(OH)30 and the minor, charged portion of boron is
subject to adsorption and desorption. The uncharged and poorly adsorbed B(OH)30 can
accumulate in near‐neutral pH waters, resulting in the high concentrations observed in
groundwater. The large observed variations in δ11B (Figure 53) could be caused by mixing of
distinct B sources and/or fractionation during speciation of boron between B(OH)30 and B(OH)4
‐.
The two North Carolina coastal plain aquifers exhibit consistent B isotope trends. In
both aquifers, decreasing boron concentrations are associated with decreasing δ11B and
increasing B/Cl‐ is associated with decreasing δ11B (Figure 53), which implies that
adsorption/desorption behavior of boron explains both B concentrations and isotope ratios.
However, the trends between the aquifers diverge in the δ11B of the apparent saline,
unmodified end point of the system. In B/Cl‐‐δ11B space, the unmodified saline end point of the
Yorktown aquifer is consistent with seawater, but in the Cape Fear aquifer system, waters with
B/Cl‐ ratios approaching the seawater ratio exhibit δ11B up to 12‰ higher than the seawater
value (Figure 53). These δ11B values and B/Cl‐ ratios both greater than the expected seawater
values in the Cape Fear aquifer system require special attention. This is because δ11B of
groundwater > 39‰ is typically associated with the residual boron subsequent to boron
adsorption, which results in low B concentrations and low B/Cl‐ ratios (Vengosh, 2003). A
possible explanation for this boron isotopic behavior lies in the evolution of Na‐HCO3‐ waters in
the Atlantic Coastal Plain as determined by previous studies. In the confined aquifers, the
freshening process that generates Na‐HCO3‐ waters is consistent with boron desorption due to
decreasing salinity, resulting in the high B/Cl‐ ratios seen in the up‐gradient waters (Figure 53).
As they flow down‐gradient, Na‐HCO3‐ waters become even more enriched in Na and
bicarbonate through a process of carbonate dissolution followed by Ca/Na rechange (Chapelle
194
et al., 1987; Knobel et al., 1998). At the site of carbonate dissolution, localized pH ≥ 9 induces
(OH)‐ removal from groundwater, releasing other anions such as fluoride from fluorapatite in
shark teeth and resulting in localized high‐fluoride groundwater in the overlying upper
Cretaceous Black Creek aquifer (Zack and Roberts, 1988; Konikow and Glynn, 2005). Similar high
F‐ concentrations (>2 mg L‐1) have been documented in Na‐HCO3‐ waters of the Cape Fear
equivalent aquifers in South Carolina (Lee, 1984), which implies that this mechanism is plausible
in the Cape Fear aquifer system where the unusual boron isotope ratios occur. The high‐pH
microenvironment is a favorable site of boron adsorption, both because of its high pH, which
dictates that B will occur as highly adsorbable B(OH)4‐ (Palmer et al., 1987), and because of the
the high Ca/Na ratio at the carbonate dissolution site, which also promotes B(OH)4‐ adsorption
due to the positive surface charge of adsorption sites (Mattigod et al., 1985; Singh and
Mattigod, 1992). This suggested microenvironment for high‐pH adsorption in the deep, down‐
gradient section of the aquifer is also consistent with the observed ion ratios showing that the
highest Na/Cl‐ ratios coincide with the highest B/Cl‐ ratios and lowest δ11B. In summary, δ11B
appears to record (1) B desorption that predominates in the shallow waters undergoing dilution,
generating waters of high B/Cl‐ and low δ11B; and (2) down‐gradient adsorption of B in the
deeper‐slower circulating waters that causes B/Cl‐ to decrease and δ11B to increase, although
absolute boron concentrations are elevated throughout the aquifer.
The high boron concentrations documented in this study, in excess of international
standards and recommendations, are among the highest reported B concentrations for fresh to
brackish groundwater, which can reach mg L‐1 concentrations in marine and fluvial/deltaic
aquifers in which freshening processes release B and Na (Ravenscroft and McArthur, 2004; Faye
et al., 2005; Stamatakis et al., 2009; Halim et al., 2010). Boron concentrations ≥ 2 mg L‐1 have
195
been documented in Na‐dominated Coastal Plain groundwater throughout the southeastern
USA (Lee, 1984), and thus are not isolated to the area of this study. Because uncharged boron
species are inefficiently removed by reverse osmosis (e.g. Vinson et al., in press), even treated
waters may exceed recommended maximum B concentrations.
6.5.5 Radium and barium
6.5.5.1 Salinity effects
The observed levels of Ra and other divalent cations in both aquifers can be separated
into (1) the up‐gradient, comparatively shallow waters in which Ca and Mg are more significant
than in deeper parts of the aquifer; and (2) down‐gradient waters along the transition from Na‐
HCO3‐ to Na‐Cl‐ hydrochemical facies (section 6.5.1). The four shallowest samples exhibit
generally higher Ra relative to chloride concentration (Figure 51), apparently representing those
wells in which cation exchange reactions are less efficient at removing alkaline earth metals
including Ba and Ra. This corresponds with the freshening process (section 6.4.1). In the down‐
gradient portion of each aquifer, radium isotope activities increase with salinity (Figure 56). This
is generally consistent with radium‐salinity trends observed in other saline aquifers and brines
(Kraemer and Reid, 1984; Miller and Sutcliffe, 1985; Krishnaswami et al., 1991; Walters, 1995;
Wood et al., 2004; Tomita et al., 2010). Under conditions in which Ra2+ and the other alkaline
earth metals are efficiently removed to the aquifer solids, the increase of Ra observed with
salinity must be due to variations in radium removal processes. The increased ion‐ion
interactions at higher salinity (activity coefficient γ < 1) diminish the thermodynamically‐active
portion (chemical activity) of the Ra2+ concentration.
In the context of overall salinity increases, variations in individual anion compositions
may also exert important controls on Ra across the salinity gradient. Speciation calculations
196
indicate that Ra is primarily composed of Ra2+ and RaSO40; monovalent Ra species such as RaCl+
and RaHCO3+ are negligible. Calculated γRa2+ decreases at ionic strength approaching seawater
concentration, whereas γRaSO40 remains near 1 because uncharged species are little affected by
ion‐ion interactions that occur at increased ionic strength (Drever, 1997). Thus, the chemical
activity of RaSO40 closely approaches that of Ra2+ (Figure 57) above 1 mM SO4
2‐ and in some
saline waters with higher sulfate concentration and SO42‐/Cl‐ ratio, RaSO4
0 is the
thermodynamically dominant species of Ra, with a maximum RaSO40 /Ra2+ ratio of 1.8 (Figure
58). When expressed as molar concentration rather than activity, reflecting the actual
distribution of radium in the system, the same general trend was observed with sulfate
concentration but the overall role of uncharged species is less pronounced. In this case, the
maximum RaSO40/Ra2+ ratio is 0.59, indicating that the dominant Ra species under these
conditions is always Ra2+ on a molar basis. The proportion of total radium as RaSO40 is in an
uncharged, poorly adsorbed form, implying that the balance between sulfate and other major
anions is a significant component in radium’s response to salinity.
Although Sr isotopes provide no direct tracer of this process because Sr is continually
removed rather than released (section 6.5.3), the speciation of alkaline earth metals (Mg, Ca, Sr,
and Ba) provides evidence for the importance of uncharged species for Ra behavior. The ratio of
the activity of the uncharged alkaline earth metal sulfate complex (MeSO40) to the divalent ion
(Me2+) increases with sulfate concentration for the waters in this study. Among alkaline earth
metals, affininty for cation exchange (to the diffuse ion swarm or outer‐sphere complexes) and
adsorption (to inner‐sphere complexes) varies with ionic radius such that the larger radii (Ra >
Ba > Sr > Ca > Mg) are preferentially removed (Sposito, 2008). Although the specific mechanism
of alkaline earth metal removal is beyond the scope of this field‐based study, cation exchange is
197
Figure 56: Relationships between radium isotope (224Ra, 226Ra, 228Ra) activities and chloride concentrations.
198
inferred to be a significant process affecting Ra and other divalent cations because of the
documented release of Na. The greater significance of the uncharged species for Ra and Ba than
for Sr, Ca, and Mg is consistent with this selectivity sequence because Ra2+ and Ba2+ are more
effectively removed than Sr2+, Ca2+, and Mg2+, and the equilibrium values of MeSO40/Me2+ are
higher for Ra and Ba as a result (Figure 58).
Due to the rapid equilibration of speciation reactions, the presence of RaSO40
effectively reduces the proportion of Ra2+ that is easily removed. Through this mechanism,
equilibrium Ra molar concentration is up to 59% higher than if all Ra in these waters were the
efficiently removed Ra2+, as noted above, given that Ra removal is species‐sensitive but that the
Ra source (alpha recoil) is not. This also implies that the dimensionless radium distribution
coefficient K decreases accordingly. However, K can vary by at least two orders of magnitude
along salinity gradients from fresh water to seawater salinity in anoxic aquifers (Sturchio et al.,
2001), and similarly, about one order of magnitude of variation in Ra/Cl‐ was observed in the
gradient from fresh water to a chloride concentration of ~10,000 mg L‐1 in this study (Figure 51).
Therefore, uncharged sulfate complexes are a secondary, enhancing factor to the ion‐ion
interactions and diminished adsorption rate that result in lower values of and K at higher
salinity.
6.5.5.2 Redox effects and the barite removal mechanism for radium
As noted previously, anoxic conditions provide additional support for recoil‐generated
radium to remain in water beyond those provided by ionic strength and aqueous complexation.
The oxidation of organic carbon in the confined aquifers of the Atlantic Coastal Plain is linked to
reduction of Mn oxides, Fe oxides, and sulfate (Chapelle and Knobel, 1985; Chapelle et al., 1987;
199
Figure 57: Calculated chemical activities of Ra2+, RaSO40, and RaCl+ along gradients of sulfate
activity.
Figure 58: Ratio of sulfate‐complexed (uncharged) to divalent alkaline earth metal activities with increasing sulfate activity.
Chapelle and McMahon, 1991; Chapelle and Lovley, 1992) and is supported by anoxic conditions
(Table 21) and the sulfur isotope evidence (Figure 54) in this study. Thus it is inferred that Fe and
200
Mn oxides are less abundant adsorption sites for Ra than in oxic systems to the extent that they
are microbially‐reducible. However, it cannot be said that these metal oxides are absent, as Fe
and Mn oxides have been observed in anoxic Atlantic Coastal Plain aquifer sediments, primarily
in non‐labile forms that have relatively low reactive surface area (Haque et al., 2008). In contrast
to redox‐sensitive adsorption onto Fe and Mn oxides, Ra removal onto the abundant clays in the
aquifer by cation exchange or adsorption is not expected to be affected by reducing conditions.
Another potential redox‐sensitive radium removal mechanism is coprecipitation into
barite (BaSO4). While it has been documented in this study that sulfate complexes increase the
amount of radium that can remain in water, whether high sulfate concentrations can facilitate
radium removal depends on whether barite saturation is reached and precipitation actually
occurs. In aquifers with active barite precipitation, Ba and SO42‐ concentrations may be
negatively correlated (Underwood et al., 2009). However, precipitation of barite is difficult to
observe directly in groundwater systems due to its trace nature (e.g. Sturchio et al., 1993).
Pervasive sulfate‐reducing conditions may (1) maintain sulfate concentrations below barite
saturation, ultimately allowing Ba2+ to accumulate (Gilkeson et al., 1983) and/or (2) subject
barite crystals to reductive dissolution (Bolze et al., 1974; Phillips et al., 2001; Martin et al.,
2003). Furthermore, cation exchange may compete with Ba precipitation by rapidly removing
Ba2+. In the two aquifers in this study, there is not a consistent negative correlation beween Ba
and sulfate concentrations (Figure 59a). A majority of the Cape Fear aquifer waters are at or
near barite saturation, but most Yorktown aquifer waters are undersaturated with respect to
barite (Figure 59a). In each aquifer, the shallow, less Na‐dominated waters exhibit the most
undersaturated conditions (Figure 59a), the highest Ba/Cl‐ ratios (Figure 59b), and the highest
Ra/Cl‐ ratios (Figure 59c). The overall undersaturation of the Yorktown aquifer with respect to
201
Figure 59: Radium and barium in relationship to barite saturation. Dashed line represents barite saturation based on Ksp = 10
‐9.97 (Drever, 1997). Circled points are the shallowest two wells in each aquifer.
barite rules out barite precipitation as a major removal mechanism for Ra in that aquifer. In
contrast, a majority of Cape Fear aquifer waters are near barite saturation, indicating that Ra
removal by barite is possible. However, the similar behavior of radium in the two aquifers, with
no obvious radium removal at high sulfate concentrations (Figure 57), may imply that despite
the near‐saturation conditions, barite precipitation is not the main radium removal mechanism
in either aquifer, perhaps because barite precipitation is precluded by sulfate reduction and/or
that Ba and Ra atoms are removed rapidly onto clays. For example, in systems with higher clay
content than aquifer conditions, Ra removal onto clays is faster than and becomes dominant
over barite precipitation (Shao et al., 2009). However, in the Coastal Plain aquifers there is not
enough information to assess any potential competition between these two mechanisms of
radium removal. Overall, given that the main mechanism(s) of radium removal in this system are
charge‐dependent, the effect of high sulfate concentrations is to make radium removal less
efficient and support higher levels of radium in water than in comparable waters of low sulfate
202
concentration. At a minimum, this seems to be the case in the Yorktown aquifer with its
relatively low barium concentrations and persistent barite undersaturation.
6.5.5.3 Radium isotope ratios
Radium isotope ratios may provide additional information about radium‐mobilizing
processes because radium is removed at a rate dependent on the atomic concentration,
whereas individual Ra isotopes are replenished by alpha recoil at different rates depending on
their respective half‐lives. In most systems, the 228Ra/226Ra ratio is primarily influenced by the
relative abundances of 232Th and 238U. Within the aquifer formations, lithologic differences in
Th/U could influence 228Ra/226Ra, indicated by the large difference in 228Ra/226Ra between the
Yorktown and Cape Fear aquifers. Values below 1 in the Yorktown aquifer (median 0.8) are
consistent with the mixed carbonate and sand lithology of the Yorktown aquifer, given that
carbonate systems have a higher affinity for U than Th (Moise et al., 2000; Sturchio et al., 2001)
while sandstones containing a significant amount of Th from continental sediments are
expected to yield 228Ra/226Ra ≥ 1 (Lively et al., 1992; Vengosh et al., 2009) . The moderate
228Ra/226Ra values of these samples also imply that marine phosphates, if present, are not a
major radium source to these waters because phosphates would exhibit low Th/U and thus
would impart low 228Ra/226Ra below ~0.5 (Moise et al., 2000; dePaul and Szabo, 2007). In the
Cape Fear aquifer, higher 228Ra/226Ra ratios (median 1.5; Figure 60) are consistent with
sandstone aquifers containing Ra from continental sources. Given that 87Sr/86Sr in this aquifer
system seems to record marine carbonate vs. continental detrital sources of Sr (section 6.5.3),
228Ra/226Ra ratios averaging above 1 (continental sandstone) and below 1 (marine carbonate)
support consistent geologic sources and interactions to explain the observed radium distribution
203
(Figure 55). Within analytical uncertainty, 224Ra/228Ra ratios are broadly near the expected value
of at or slightly above 1 for a system in which Ra is contributed by recoil and chemical removal is
rapid, validating this assumption about the primary Ra source mechanism (Krishnaswami et al.,
1982). Also, median 224Ra/228Ra is broadly similar between the aquifers, with median of 1.1 in
the Yorktown aquifer and 1.3 in the Cape Fear aquifer (Figure 60). Overall, evidence such as
228Ra/226Ra suggests that solid‐phase radionuclide sources are different between the two
aquifers, and that the higher Ra/Cl‐ ratios in the Cape Fear aquifer (Figure 51) are probably the
result of higher solid‐phase radionuclide levels in the Cape Fear aquifer rather than Ra removal
being less effective in that aquifer.
6.6 Conclusions
Overall, the primary removal mechanism for Ra and other alkaline earth metals is
charge‐dependent removal of divalent cations, probably onto exchange sites on clays, although
adsorption sites such as non‐microbially‐reducible metal oxides may also be present. Ra levels
are not especially high compared to some other sand and sandstone aquifers; only in the most
saline samples do radium activities approach drinking water standards. The observed low to
moderate activities of Ra can be explained primarily by (1) Ra sources consistent with alpha
recoil from the sandstone aquifer matrix and (2) charge‐sensitive mechanisms that remove Ra
and the other alkaline earth metals. Radium removal mechanisms become progressively less
effective as salinity increases because of ion‐ion interactions at increasing ionic strength and the
secondary enhancing role of uncharged RaSO40. When considering all alkaline earth metals (Mg,
Ca, Sr, Ba, and Ra), the uncharged sulfate species is increasingly important with salinity for the
heavier elements Ba and Ra (indicated by higher MeSO40/Me2+). In terms of molarity, the
204
Figure 60: Differences in radium isotope ratios (228Ra/226Ra and 224Ra/228Ra) between the Yorktown and Cape Fear aquifers.
presence of the RaSO40 species supports up to 59% higher total radium than if all radium were
present as Ra2+. Radium isotope results are consistent with the behavior of the complementary
Sr and B isotope systems. 87Sr/86Sr, like 228Ra/226Ra, seems to fingerprint marine vs. nonmarine
lithologic sources of Sr and radionuclides. Boron isotopes provide insights on the freshening
process that removes alkaline earth metals, including Ra, and yields Na‐HCO3‐ waters with high
boron concentrations. Overall, the geochemistry of this aquifer system influenced by cation
exchange imparts strong effects on the isotopic systems of Ra and B. The formation of
uncharged species of alkaline earth metals and boron, such as RaSO40 and B(OH)3
0, is a
significant means by which these elements are not removed to the solids in a strongly retentive
system such as the Atlantic Coastal Plain.
205
7. Synthesis
7.1 The components of salinity: Cation and anion effects
The five groundwater systems in this study encompass a large range of fresh to saline
and oxic to strongly reducing conditions commonly observed in aquifers (Figure 2). Direct
comparisons between these field‐based studies are approximate because the solid‐phase
radionuclide distribution varies between these aquifers. However, as components of the overall
salinity in a system, certain anions and cations impart consistent influences on radium because
they directly or indirectly affect radium removal mechanisms (Table 23).
Table 23: Components of salinity and redox conditions that are correlated with increasing (↑) or decreasing (↓) radium activities in the five studied groundwater systems.
Aquifer system Cl‐ SO42‐
Ca or Ca/Na
Ba Fe/Mn oxide reduction
Sulfate reduction vs. BaSO4
Crystalline rocks ↑ ↑
Willcox Basin ↑
Jordan sandstone ↓ ↓ ↑
Agadir coastal aquifer ↑ ↓ ↑ ↓
Atlantic Coastal Plain ↑ ↑ ↑ ↑ ↑ ↑
The chloride ion uniformly supports higher levels of radium in water because it is
associated with higher ionic strength, which makes adsorption of Ra2+ less efficient, because
high‐Cl‐ waters also contain cations that may compete with Ra for adsorption sites, and to a
much lesser extent through formation of the RaCl+ complex. In both of the studies with chloride
concentrations exceeding 1000 mg L‐1, linear relationships were seen between chloride
concentration and 224Ra (Agadir coastal aquifer) and with all radium isotopes (Atlantic Coastal
Plain). Moreover, these relationships exhibit similar estimated slopes of ~12‐16 mBq L‐1 of
206
radium activity increase per 1000 mg L‐1 of TDS concentration increase in Na‐Cl type waters
(Figure 38, Figure 56) despite the different inferred Ra removal mechanisms in each aquifer. A
recent analysis of radium in Na‐Cl‐ waters and brines in several studies worldwide with TDS of
103‐105 mg L‐1 (Tomita et al., 2010) found a slightly higher slope, but similar magnitude in light of
the uncertainties involved: ~100 mBq L‐1 per 1000 mg L‐1 TDS.
High sulfate concentrations favor the uncharged and thus weakly adsorbed RaSO40
species, which should inhibit the overall removal of Ra in high‐sulfate waters. However, once
barite saturation is reached, high sulfate concentrations could contribute to Ra removal by
coprecipitation into barite. Because of the redox‐sensitive nature of barite and the possibly
slower rate of radium removal by barite precipitation relative to other removal mechanisms, the
effect of barite on groundwater radium is mixed, depending on redox conditions and cation
ratios in addition to the saturation state of the mineral.
Calcium and other alkaline earth metals exhibit some association with radium in all of
the studies with significant variations of Ca concentrations and/or Ca/Na ratios. Only in the
Jordan aquifer, where the entire aquifer exhibits similar concentrations and proportions of Ca,
was no such correlation observed. Consistent with the apparent effects of Ca, however, the Ca‐
dominated Jordan aquifer exhibits the highest Ra relative to salinity of all five study areas. These
consistent indications between the different study areas suggest that in fresh to slightly saline
waters, the competition among alkaline earth metals for charge‐sensitive removal sites strongly
influences the removal of radium. Thus, although radium released by cation exchange is not
expected to be a significant mechanism of direct radium mobilization to waters, cation exchange
can nonetheless promote higher levels of radium by interfering with Ra removal upon
207
introducing competing divalent cations. Conversely, systems in which divalent cations are
removed from groundwater in exchange for Na exhibit more rapid and effective Ra removal.
As an alkaline earth metal and chemical analogue to radium, high Ba concentrations also
support higher levels of Ra in water due to increased competition for adsorption sites as long as
barite is undersaturated. At barite saturation, the precipitation of BaSO4 is a significant removal
mechanism for radium, and thus some high‐Ba aquifers exhibit enhanced radium removal that
partially mitigates the effects of divalent cation concentrations, e.g. high Ca/Na ratio. As with
sulfate concentrations, evaluating the effects of Ba concentrations requires understanding of
the barite saturation state, the possible effects of sulfate‐reducing conditions, and other Ba
removal mechanisms such as cation exchange, which may be more rapid.
7.2 Salinity vs. redox effects
A similar analysis to that of Tomita et al. (2010), combining the results of the sand and
sandstone aquifers in this study (fractured‐rock waters excluded), indicates that fresh waters
with TDS < 1000 mg L‐1 consistently exhibit radium‐226 levels higher than the expected
downward extrapolation of the salinity trend, but especially in the anoxic fresh waters of the
Jordan aquifer (Figure 61). In contrast, the oxic Willcox Basin waters exhibit very low radium,
both in absolute activities and relative to the other fresh water systems. It must be emphasized
that radionuclide levels of the aquifer solids differ between aquifers and also influence these
results, but even the low‐U quartz sandstones in the Jordan sandstone seem to support high
groundwater Ra relative to the low salinity. This must be due to the anoxic, Fe‐ and Mn‐oxide
reducing conditions of the Jordan aquifer.
208
Figure 61: Relationship between total dissolved solids (TDS) and 226Ra activity in sand aquifers in this study. Dashed line is a linear regression of the Atlantic Coastal Plain waters with TDS >
103.5 mg L‐1 (slope ~15 mBq L‐1 per 1000 mg L‐1 TDS).
Therefore, within the salinity range considered in this study (TDS < ~20,000 mg L‐1),
redox conditions impart a substantial influence on radium mobility. Unlike some of the
components of salinity that exhibit mixed effects on radium mobility, like sulfate concentration,
reduction along the thermodynamic sequence from oxic to sulfate‐reducing conditions would
apparently always enhance radium mobility in a consistent way (Table 23) by removing
adsorption sites (Mn and Fe oxides) and by inhibiting barite precipitation through sulfate
reduction. Overall, the observation that radium mobilization increases by approximately one
order of magnitude between oxic and Mn/Fe‐oxide reducing conditions suggests that this
commonly‐encountered redox gradient can impart a similar effect on radium removal as, for
209
example, a TDS increase from 103.5 to 104 mg L‐1 (~3000‐10,000 mg L‐1; Figure 61). Thus, from the
perspective of evaluating water quality impacts of changing groundwater conditions on the
naturally‐occurring radionuclide distribution in an aquifer, commonly observed redox conditions
in fresh waters used for drinking correspond to the salinity effects in significantly more saline
waters than typically used for drinking without desalination. The highest levels of radium
activities in groundwater are therefore expected in both saline and reducing aquifer conditions.
210
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Biography
David Vinson was born in Birmingham, Alabama in 1978. He received a B.S. degree from
the University of Alabama (2000) with double major in Geology and American Studies, and a
M.S. degree in Earth and Planetary Sciences from the University of New Mexico (2002). First‐
authored journal articles in print include: (1) Biogeochemistry at the zone of intermittent
saturation: Field‐based study of the shallow alluvial aquifer, Rio Grande, New Mexico
(Geosphere, v. 3, p. 366‐380, 2007); (2) Radon transfer from groundwater used in showers to
indoor air (Applied Geochemistry, v. 23, p.2676‐2685, 2008); and (3) Relationships between
radium and radon occurrence and hydrochemistry in fresh groundwater from fractured
crystalline rocks, North Carolina (USA) (Chemical Geology, v. 260, p. 159‐171, 2009). Vinson
received the Kelley‐Silver departmental fellowship of the University of New Mexico Department
of Earth & Planetary Sciences (2000‐2002), Outstanding Student Paper Award from the
American Geophysical Union, Biogeosciences Section (2002 spring meeting), Outstanding MS
student in 2002‐03 award from the University of New Mexico Department of Earth & Planetary
Sciences, a Geological Society of America student research grant (2006), Best Student
Presentation award from Geological Society of America Geology & Society Division (for 2006
meeting – awarded 2007), an International Association of Geochemistry student research grant
(2009), and a Clay Minerals Society student research grant (2009).
