### 1 

# Groundwater Chemistry Changes as a Result of CO 2 

# Injection at the ZERT Field Site in Bozeman, Montana

Project Report

Contributors in Alphabetical Order:

Lawrence Berkeley National Laboratory: John A. Apps, Jens T. Birkholzer, Nicolas Spycher, Liange Zheng

U.S. Geological Survey: Gil Ambats, William N. Herkelrath, Yousif, K. Kharaka, James J. Thordsen, Evangelos Kakouros, Sarah Beers

Montana State University: Kadie S. Gullickson, Lee H. Spangler

### 2 TABLE OF CONTENTS

1. INTRODUC TION

.......................................................................................................... 6 1.1. Understanding the Problem................................................................................................. .6 1.2. Scope and Objective.............................................................................................................7

2. DESCRIPTION OF THE ZERT FIEL D SITE ................................................................ 7 2.1. Regional and local setting ................................................................................................ ....8 2.2. Geology ................................................................................................................... .............9 2.3. Test Configuration and Groundwater Flow

......................................................................... 9 2.4. Water Sampling and Analysis ............................................................................................ 11

3. WATER SAMPLI NG RE SULTS

................................................................................ 12 3.1. Dissolved Inorganic Chemicals.......................................................................................... 13 3.2. Dissolved Organic Chemicals ............................................................................................ 20

4. MINERALOGICAL DATA

.......................................................................................... 21 4.1. Core Analysis and Methods ...............................................................................................21 4.2. Mineralogical Results..................................................................................................... ....22
5. HYDROGEOLOGICAL AND GEOCHE MICAL INTERP RETATION

......................... 25 5.1. Transport of CO 2 

### in Gas and Aqueous Phases................................................................... 26 5.2. Geochemical Response to CO 2 

### Release and Interpretation ............................................... 27 5.3 Geochemical Modeling ...................................................................................................... 55
6. SUMMARY AND CO NCL UDING RE MARKS............................................................ 69

7. ACKNOWLEDGMENTS and Discl aimer

.................................................................. 72

8. REFERE NCES

........................................................................................................... 72

APPENDI X A

.................................................................................................................... 79

APPENDI X B

.................................................................................................................... 82

APPENDI X C

.................................................................................................................... 83

APPENDI X D

.................................................................................................................... 85

### 3 List of Figures

Figure 2.1 Packer system installed in horizon tal well (from Spangler et al., 2009)...................... 8

Figure 2.2 Location of water monitoring wells in re lation to the surface tr ace of the slotted

horizontal CO 2 

### injection pipe in zone 6 of the ZERT site. The gridpoints in the figure show refer to a grid with 10 m spacing that was established to provide a common reference system for researchers. ............................................................................... 11

Figure 3.1 Cumulative precipitation and groundwater leve ls in selected ZERT wells during the sampling period in 2008............................................................................................. 13

Figure 3.2 Modified Stiff diagrams showing concen trations (equivalent units normalized to 100%) of major cations and anions, together with salinity and pH of groundwater from well 2B before and during CO 2 

### injection. .........................................................14

Figure 3.3 Field- measured groundwater pH values (a), alkalinities (b ), and electrical conductance (c), obtained from selected ZERT wells as a function of time of sampling. Note the systematic decrease in pH values from ~7.0 before CO 2 

### injection to values as low as 5.6 during injecti on, and subsequent pH increases after CO 2 

### injection ended. Alkalinities increas ed from about 400 mg/L before CO 2 

### injection to values close to 1200 mg/L as HCO 3 

### , and electrical conductanc e also increased from about 600 μS/cm before CO 2 

### injection to values hi gher than 1,600 μS/cm during injection......................................................................................................................15

Figure 3.4 Concentrations of major cations in gr oundwater from the ZERT wells plotted as a

function of water alkalinities. Note the relati vely constant concentrations of Na and K, but the general increases in the concentrations of divalent cations with water alkalinities, possibly in dicating dissolution of carbonate minerals............................16

Figure 3.5 Ca/Mg weight ratios plotted as a function of time of sampling. Note that the ratios are higher than the value expected from dissolution of dolomite (1.6), and are relatively constant for well 4B, but c ontinue to increase for well 2B until CO 2 

### injection was terminated, possibly indicating dissolution of calcite..........................17

Figure 3.6 Concentrations of Fe (a) and Mn (b) in groundwater from selected ZERT wells plotted as a function of time of sampling. Note the low Fe and Mn concentrations during July 20 to July 26, which we are attributing to the ox idizing conditions possibly caused by percolating oxygenated wa ter from rainfall events (see Figure 3.1)..............................................................................................................................18

Figure 3.7 Trace metal concentrations in groundwater from selected ZERT wells plotted as a

function of water alkalinities. Some of the scatter shown probably results because the reported values in some cases are close to the detection limits.................................. 19

Figure 3.8 Concentrations of BTEX compounds de tected in groundwater from selected ZERT wells plotted as a function of time of sampling. The B TEX concentrations are all significantly below the maximum contam inant levels (MCLs) (e.g., 5 ppb for benzene). .................................................................................................................... 20

Figure 4.1 Grain size fractions for di fferent zones in Well W6. ................................................. 22

Figure 4.2 Grain size fractions for 15 gravel sa mples analyzed from Wells 6, 7, and 8............. 23

Figure 5.1 Schematic representation of CO 2 

### transport in ZERT field site. The cross-section shown is along the axis of well pa irs 1A/B, 2A/B, 3A /B, and 4A/B. ........................27

### 4 Figure 5.2 Correlation plots showing relationships between vari ous species in solution ........... 33

Figure 5.3 Normalized saturation indices of magnetit e, goethite, siderite, green rust (fougeriteCO 3 

### ) and Fe(OH) 3 

### (ferrihydrite) plotted as a function of pe......................................49

Figure 5.4 Normalized saturation indi ces of pyrolusite, bixbyite, manganite, hausmannite and rhodochrosite, plotted as a function of pe. ................................................................50

Figure 5.5 Saturation index of calcite as a function of log P(CO 2 

### ) calculated for groundwater samples taken between July 7 and July 17, 2008 at the ZERT site. For an explanation of the trend lines , please refer to the text................................................ 54

Figure 5.6 Calculated pH with different co mbination of chemical reactions.............................. 62

Figure 5.7 Calculated calcium concentration vs pH with different calcite dissolution rate. ....... 63

Figure 5.8 Calculated sodium concentration vs pH..................................................................... 63

Figure 5.9 Calculated potassium concentration vs pH. ............................................................... 64

Figure 5.10 Calculated magnesium concentration vs pH. ............................................................. 64

Figure 5.11 Calculated lead concentration vs pH.......................................................................... 64

Figure 5.12 Calculated cadmium concentration vs pH.................................................................. 65

Figure 5.13 Calculated copper concentration vs pH......................................................................65

Figure 5.14 Calculated zinc concentration vs pH.......................................................................... 65

Figure 5.15 Calculated iron concentration vs pH in the base model, in which only cation exchange is considered a nd a sensitivity run in whic h both cation exchange and

siderite dissolution are considered ............................................................................. 66

Figure 5.16 Calculated bicarbonate concentration vs pH..............................................................67

Figure 5.17 Calculated arsenic concentration vs pH ..................................................................... 67

Figure 5.18 Calculated selenium concentration vs pH .................................................................. 68

Figure 5.19 Calculated phosphate concentration vs pH ................................................................ 68

Figure D.1 Normalized saturation indices of chrom ite, magnesiochromite, and the chromium spinels of Cd, Co, Cu and Zn, Eskolaite a nd Cr-ferrihydrite, plo tted as a function of pe. ...............................................................................................................................85

Figure D.2 Normalized saturation indices of copper, bornite, covellite, chal cocite, cuprite and tenorite, plotted as a function of pe. ..........................................................................86

Figure D.3 Normalized saturation indices of sele nium, molybdenum selenide, cadmoselite and clausthalite, plotted as a function of pe. ..................................................................... 87

### 5 List of Tables

Table 5.1 Correlation matrix of the log molar concentrations of all species analyzed in MSUZERT groundwaters for the period July 8 to July 17, 2008....................................... 31

Table 5.2 Principal components derived from an analysis of groundwater chemical compositions for the period July 7 to July 17, 2008 .................................................. 36

Table 5.3 Correlation of aqueous species with prin cipal components for the period July 7 to July 17, 2008 in relation to shor t-term chemical processes ....................................... 39

Table 5.4 Assumed initial composition of the groundw ater at the ZERT site prior to CO 2 

### injection......................................................................................................................43

Table 5.5 Saturation indices of redox-insensitive primary and seconda ry minerals with respect to groundwater at the ZERT site ................................................................................ 46

Table 5.6 Variation in principal aqueous species of selected redox-sens itive species as a function of pe at circum-neutral pH (pH = 6.9). ........................................................ 48

Table 5.7 Distribution of the predominant Cr, Fe and Se species in solution at pH = 6.82 and pe

= +9 ............................................................................................................................51

Table 5.8 Calculated calcite saturation index and CO 2 

### partial pressure for groundwaters taken at the ZERT site between July 7 and July 17, 2008 ................................................... 53

Table 5.9 Primary species and thei r initial concentration .......................................................... 57

Table 5.10 Minerals considered in the model. ............................................................................. 57

Table 5.11 Aqueous complexes and their dissociation constants ................................................58

Table 5.12 Cation exchange reactions a nd selectivity coefficients.............................................. 59

Table 5.13 Surface complexation reactions and surf ace complexation constants (logK) on HFO

............................................................................................................................... .....59

Table 5.14 Equilibrium constants for minerals 1 

### ...........................................................................61

Table 5.15 Initial volume fraction and kinetic parameters for minerals considered in the model 1



### ............................................................................................................................... .....62

Table 5.16 Initial (C i 

### ) and final (C f 

### ) adsorbed concentration and their difference C f 

### –C i 

### ...........67

Table A1 Chemical composition of water sa mples from ZERT monitoring wells.................... 79

Table B1 Well log for well 6 .....................................................................................................82

### 6 

## 1. INTRODUCTION

1.1. Understanding the Problem

Combustion of fossil fuels produces CO 2 

### , a common greenhouse gas linked to global climate change. Separation of CO 2 

### from emissions produced by large industrial point sources like power plants, cement kilns and refineries, and injection deep underground into geologic formations is one method of preventing CO 2 

### releases into the atmosphere. This process is referred to as Carbon Capture and Storage (CCS). CCS is one of several solutions being considered to mitigate global climate change. Other solutions include increased energy efficiency, renewables, nuclear power, advanced coal, and plug-in hybrid electric vehicles.

Injection of CO 2 

### into deep geologic reservoirs is one solution, but this method of sequestering carbon is not without its own unique set of environmental risks, summarized by the United States Environmental Protection Agency (US EPA) in its draft rules gove rning underground injection of

CO 2 

### (US EPA, 2008). The US EPA designates potabl e groundwater aquifers as those underground sources of drinking water (USDW) worthy of protection under th e Safe Drinking Water Act (US EPA, 2004). Many municipalities, businesses, and priv ate individuals across the United States rely on these USDW as their primary source of drin king water, or for process water used in manufacturing, power production, etc.

Compressed CO 2 

### (as a free phase liquid or supercritical fl uid) is less dense than water under most conditions, causing it to float (like oil on water) and move upward if unconstr ained by fine-grained, low permeability rocks or sediments (such as shale, mudstone, clay or salt). Therefore, if not properly capped (analogous to a soda bottle), CO 2 

### has the potential to leak from the storage formation where it was injected and move as a buoyant fluid into overlying USDW. Once CO 2 

### enters a USDW, it readily dissolves in the groundwater, forming a weak acid (i.e., carbonic acid), which lowers the pH of the groundwater and increas es its dissolved carbonate content, according to the following overall reaction:

CO 2(g) 

### + H 2 

### O ==> H + 

### + HCO3 – 

### (1-1)

The magnitude and extent of acidification and alka linity increase is exp ected to depend on the degree of buffering by aquifer materials (such as carbonate minerals found in many aquifers) and could result in the release of trace metals by di ssolution and/or desorption and ion exchange reactions. Both of these processes could mobilize heavy metals including lead (Pb), cadmium (Cd) and arsenic (As) into groundwater (e.g., Wang a nd Jaffe, 2004; Kharaka et al., 2006; Birkholzer et al., 2008; Kharaka et al., 2009b). The solven t properties exhibited by supercritical CO 2 

### may also leach organic compounds including benzene, toluene, ethylbenzene and xylenes (BTEX), phenols, polynuclear aromatic hydrocarbons (PAHs) and other compounds, from organic matter present in the aquifer. The release of these constituents at levels exceeding US EPA maximum contaminant levels (MCL) could become a potential health ha zard. Therefore, even though the risk of CO 2 

### migration from the storage reservoir is considered to be small with proper site selection and a regulatory framework designed to oversee, moni tor and control risk (IPCC, 2007), the potential impact the CO 2 

### can have on potable groundwater resources warrants further study.

### 7 1.2. Scope and Objective

In this report, we describe the geochemi cal and hydrological results of a shallow CO 2 

### -injection test performed during July and August 2008 at a field s ite created by the Zero Emission Research and Technology Center (ZERT), located on Montana St ate University (MSU) property near Bozeman, Montana. From July 9 to August 7, 2008, approximately 300 kg/day of food-grade CO 2 

### was injected through a perforated pi pe placed approximately 1.1–2.5 m below ground surface, and about 0.5–1 m below the water table. The objective of this study was to investigate, using a small-scale field experiment, whether sequestered CO 2 

### released from a hypothetical geologic storage reservoir would have an adverse impact on USDW. The objective was accomplished by collecting approximately 80 shallow groundwater samples before, during, and after CO 2 

### gas injection and analyzing these samples for their concentrations of major, minor, and trac e inorganic and organic compounds. The scope of the study was limited to the potential impact of gaseous CO 2 

### on potable groundwater and did not consider or evaluate the impact of displaced saline groundwater on shallow freshwater aquifers. The study was undertaken as a collaborative pr oject between the Earth Sciences Division at Lawrence Berkeley National Laboratory (LBNL) and the U.S. Geological Survey (USGS), in part funded by the Electroni c Power Research Institute (EPRI).



## 2. DESCRIPTION OF THE ZERT FIELD SITE

ZERT was initially created as a collaboration in volving two universities and six DOE national laboratories. ZERT started in 2006 to develop a fi eld test site where controlled, shallow subsurface

release of CO 2 

### would be possible over a significant spatial extent. The primary purpose of the ZERT field site, located on MSU property near Bozem an, Montana, was to allow controlled studies

of near-surface CO 2 

### transport and detection technologies. A slotted horiz ontal well of about 70 m length divided into six zones was installed into the shallow subsurface (Figure 2.1). The average depth of the slotted section of the horizontal well is on average about 1.8 m below ground surface, varying from up to 2.5 m at the southwestern end (Z one V and VI) to 1.1 m at the northwestern end (Zone I). On average, the horiz ontal well is about 0.5–1 m below th e water table (also see Section 5.1). The scale and CO 2 

### release rates were chosen to be re presentative of possible leakage rates from large-scale CO 2 

### storage projects. Controlled releases of CO 2 

### were performed in the summers of 2007 and 2008. A wide variety of detection techni ques were deployed by various collaborators to determine the efficacy of monito ring strategies, and to better understand transport from the saturated zone to the vadose zone and ultimately to the atmosphere (Spangler et al., 2009). This report focuses primarily on the geochemical and hydrological results of shallow CO 2 

### injection into the groundwater at the ZERT site, based on the groundwater samples and other relevant data obtained in the 2008 release experi ment (Kharaka et al., 2009a). Details on other aspects of the project can be found in Sp angler et al. (2009).

### 8 Figure 2.1 . Packer system installed in horizont al well (from Spangler et al., 2009).



### 2.1. Regional and local setting

The ZERT field site is located on a relatively flat 12 hectare agricultural plot at the western edge of the MSU-Bozeman campus in Bozeman, Montana, USA (see Spangler et al., 2009, for description of regional setting, and the various investigations completed or currently ongoing at this site). Plant coverage is about 70% gr asses with the remainder being alfa lfa, clover, dandelion, and thistle.

The ZERT site is located at an elevation of 1,495 m near the southeastern boundary of the Gallatin Valley, which is a north-south-trending intermontane basin of about 1,350 km 2 

### in area (Kendy and Tresch, 1996). Structurally, the Gallatin Valley is an eastward tilted grabe n, a broad alluvial plain filled with up to 1,800 m of poorly consolidated Cenozoic sandy gravel deposits. At the ZERT site, about a meter of silts and clays with a top soil la yer of variable thickness overlie the sandy gravel. Test pits and cores collected from several locations on the field consistently revealed two distinct soil zones. The topsoil of organic silt and clay with some sand ranges in thickness from 0.2 to 1.2 m. A caliche layer, high in calcite and clay, is observed at depths of ~50–80 cm, but only small

traces of carbonate are observed in fines above 2.5 m and in megascopic de trital carbonate. Beneath the topsoil layer is a cohesionl ess deposit of sandy gravel exte nding to 5 m, the maximum depth investigated (Spangler et al., 2009). More information on core an alysis and local mineralogy is given in Section 4.

### 9 2.2. Geology

The Gallatin Valley basement and surrounding mountai ns consist of complexly folded and faulted Tertiary through Archean bedrock.The eastern boundary of this valley is a series of steep normal faults along the fronts of the Br idger and Gallatin Ranges. The Ga llatin Range, which borders the south margin and crests above 3,000 m elevation, pr ovides source material to the basin-fill deposits in the portion of the Gallatin Valley which includes the ZERT field site.

The basin-fill deposits are subdivi ded by Kendy (2001) into three hydr ogeologic units: Quaternary alluvium (Qal), Quaternary and Tertiary undiffe rentiated deposits (QTd), and Tertiary Bozeman group (Tb). The Tertiary Bozeman Group is consider ed to underlie the entire Gallatin Valley; it outcrops as pediments on the Ma dison Plateau and benches east of Bozeman, but it is mostly covered by a veneer of Quaternary and/or Te rtiary alluvium. Quaternary and Tertiary undifferentiated deposits consists of unconsolidat ed Quaternary alluvi al-fan deposits overlying Tertiary sedimentary pediments. Alluvial fan deposits extend into the Gallatin Valley from the Bridger and Gallatin Range fronts and consist of a heterogeneous mixture of unconsolidated poorly sorted rock fragments in a sand, silt and clay matrix with some carbonate cement (Kendy and Tresch, 1996; Lonn and English, 2002) . The largest of these alluvial fans, extends northwards from the Gallatin Range front into the alluvial plain between Hyalite and Bozeman Creeks. The city of Bozeman and much of the surrounding land, includi ng the ZERT field site, is located on this alluvial fan and is mapped, in Kendy (2001) as Quaternary and Tertiary undifferentiated deposit (Qtd).

The Gallatin Range provides the source material fo r the alluvial fan deposits in and around the ZERT site. The bedrock geology for the Gallatin Range is grouped by Kendy (2001) into four hydrogeologic units: (1) Tertiary (E ocene) volcanic rock s (Tv), which include andesite and basalt flows, breccia, agglomerate, and tuff of the Ga llatin-Absaroka Volcanics; (2) Tertiary through Middle Proterozoic sedimentary bedrock (Tysaq ), which include a number of sandstone and limestone formations, and the Middle Proteroz oic Belt Supergroup. Sandst one, conglomerate, and limestone predominate the lithologies; (3) Cret aceous through Cambrian sedimentary bedrock (KCscu), in which fine-grained sandstone, shale, mudstone, siltstone, and chert predominate the lithologies; (4) Archean metamorphic rocks (Am), wh ich consists of crystalline rocks, mainly biotite gneiss, with some schist , quartzite, and marble. Tertiary volcanic rocks (Tv) and Archean metamorphic rocks (Am) predominate in the Gall atin Range and appear to be the major rock lithologies in poorly to moderately sorted sand and gravel at the ZERT field site.

2.3. Test Configuration and Groundwater Flow

Five pairs of polyvinyl chlori de (PVC) groundwater monitoring wells, 5 cm in diameter and extending to depths of about 1.5 m (B-wells) and about 3.0 m (A-we lls) below ground surface, were installed in the summer of 2007. These were emplaced at the southwestern end of the slotted section

of the horizontal CO 2 

### injection pipe (zone 6) to measure changes in the chemical composition of shallow groundwater and head-space gas, following CO 2 

### injection (Figure 2.2). The wells are screened at the bottom 0.76 cm, and are located at distances ranging from 1 to 6 m from the

horizontal CO 2 

### injection well. Comparison of the elevation of the horizontal inj ection well in zone 6 with the screening depth of the gr oundwater wells reveals that the CO 2 

### injection occurs

### 10 approximately 0.26 to 0.33 m above the top of the sc reening interval in the deep A-wells, and about 0.50 to 0.60 m below the bottom of the screening interv al in the shallow B-wells (see schematic in Figure 5.1).

The groundwater hydraulic gradient and flow direct ion were estimated to be 17 degrees west of north (Figure 2.2). Four of the monitoring well pair s were installed north of the horizontal injection well (2A/B, 3A/B, 4A/B, and 5A/B ) in the downgradient groundwater flow direction; one pair (1A/B) was installed south or hydraulically upgradient (Figure 2.2) from the horizontal well. In December 2008, three additional monitoring wells we re installed (W6, W7, and W8), primarily to

obtain core for mineral characterization and labor atory experiments. Well W6 is 2.6 m deep with

the bottom 0.76 m screened. Wells W7 and W8 are 3 m deep with the bottom 1.5 m screened. Four

additional water wells in the area have been us ed to estimate groundwat er flow gradients and direction (MW1 through MW4); however , these wells are too far away fr om the injection well to be affected by CO 2 

### injection.

The depth of the water table at the ZERT field site is close to the gr ound surface, but varies seasonally. Spangler et al . (2009) report a variation of about 1 m in the depth of the water table over

the last two and a half years. In the summers of 2007 a nd 2008, the depth of the water table measured in the water monitoring wells in zone 6 ranged from approximately 1.35 m to slightly more than 1.5 below ground surface, indicating th at the shallow wells were just deep enough to collect water samples. Percol ation tests made in 2006 yielde d hydraulic conductivity values approximately 0.76 and 1.4 m/day in the upper soil (sa ndy silt) layer and in th e sand-gravel aquifer, respectively. Results of tracer tests conducted in 2009 showed the lateral groundwater flow is high at about 2 m/day, and higher than previously estimat ed (Spangler et al., 2009). Calculated rates of

CO 2 

### diffusion and advection in the unsaturated zone are at least an order of magnitude higher than the groundwater flow rate (Lewicki et al., 2007; Lewicki et al., 2009). Also, anal yses of gas samples obtained from the head space of the wells yielded P(CO 2 

### ) values as high as 0.95 bar relatively rapidly (Strazisar, 2009; Strazisar et al., 2009). These data indicate that the rapid and systematic changes observed in the chemical composition of groundwater are likely driven by both fast spreading of gaseous CO 2 

### in the sand-gravel unsaturated zone under the upper soil layer, accompanied by re-dissolution into the groundwater, as well as by the transport of dissolved CO 2 

### in the groundwater (see also Section 5.1).

### 11 10m 5B 5A W6 W7

W8 1A 2A 3A 4A 4B

3B

2B

1B (-3,1)

groundwater gradient

(-3,0) 17 ° 

### N

(-4,0)



### Horizontal
well

(-4,-1) 

### Monitoring Wells

Depth = 3 m Depth = 1.5 m

Installed 2008

Depth ~ 3 m

Gridpoints (10m) 

### 10m 5B 5A W6 W7

W8 1A 2A 3A 4A 4B

3B

2B

1B (-3,1)

groundwater gradient

(-3,0) 17 ° 

### N

(-4,0)



### Horizontal
well

(-4,-1) 

### Monitoring Wells

Depth = 3 m Depth = 1.5 m

Installed 2008

Depth ~ 3 m

Gridpoints (10m) 

### Monitoring Wells

Depth = 3 m Depth = 1.5 m

Installed 2008

Depth ~ 3 m

Gridpoints (10m) 

### Monitoring Wells

Depth = 3 m Depth = 1.5 m

Installed 2008

Depth ~ 3 m

Gridpoints (10m)

Figure 2.2. Location of water monitoring wells in relation to the surface trace of the slotted horizontal CO 2 injection pipe in zone 6 of the ZERT site. The gridpoints in the figure refer to a grid with 10 m spacing that was established to provide a common reference system for researchers.



### 2.4. Water Sampling and Analysis

Most of the water samples for this study were obtained from the shallo w groundwater wells (1.5 m deep). The water from the deeper wells was not significantly impacted by the injected CO 2 

### , likely

because these are screened between 0.26 and 0.33 m be low the level of the injection pipe. Sample collection, preservation, and field and laboratory chemical analyses of water and some relevant isotopes were carried out by me thods described in Kharaka and Hanor (2007). Groundwater samples were collected using a peristaltic pump after purging the well by co llecting 0.5–1 L aliquot that was used for rinsing the sampling containe rs. On-site chemical analyses of groundwater samples, carried out in a USGS mobile laboratory, included meas urements of pH, Eh, conductivity, alkalinity, dissolved oxygen (DO) and temperature. Sample filtration, acidification, and preservation were also carried out in the USGS field laborat ory. Groundwater samples for metal analysis were filtered using a 0.1 micron filter, then acidified using ultrex HNO 3 

### and kept in a refrigerator and shipped on ice to the USGS laboratory in Menlo Park, California.

Laboratory analyses for cations and metals were evaluated by comprehensive inductively coupled plasma mass spectrometry (ICP-MS), with focus on concentrations of various trace elements, including As, B, Cd, Cr, Cu, Fe, Hg, Mn, Pb, Se, U, and Zn. Sophisticated ICP-MS equipment was necessary because these elements of interest have very low bac kground concentrations in the site groundwater, often close to the dete ction limits. Major and minor ani ons were determined in USGS laboratory by ion chromatography (IC). Benzene, toluene, ethyl-benzene, m-, p- and o-xylene (BTEX) were analyzed using an SRI-8610C gas chromatograph as described in EPA Method

### 12 5030C, a purge-and-trap procedure fo r the analysis of low levels of volatile organic compounds (VOCs) in aqueous samples and water miscib le liquid samples. The gas chromatographic determinative steps are found in EPA methods 8015 and 8021 (US EPA, 2003).



## 3. WATER SAMPLING RESULTS

Detailed chemical analyses of over 60 groundwater samples collected from the ZERT wells in July– August 2008 have been completed. Most of the sample s analyzed are from four of the five shallow 1.5 m deep “B” wells; results from Well 3B are incomplete because we were unable to obtain adequate volumes of water for our analytical requirements, as wa ter recovery in this well was generally slow. The chemical data obtained for th e samples are listed in Appendix A. The chemical composition of water obtained from the 3 m-deep “A” wells are similar to those listed for sample Z109 (Table A1), collected prior to CO 2 

### injection in Well 2B, and remained relatively unchanged following CO 2 

### injection primarily because their perforations are located below the CO 2 

### injection pipe.

The major cations and anions listed in Appendi x A and plotted in the included figures carry analytical uncertainties of a pproximately ±3%. The uncertaintie s are ±5% for trace metals, and these may be more than ±10% for trace metals that are close to their detection limits. For sample collection, preservation, analysis, and uncertainties applicable to these results, see Kharaka and Hanor (2007) and references therein.

Sample collection started in July 7, 2008, before the CO 2 

### injection, which began at 15:45 on July 9, 2008. Injection was continuous, except for a few shor t intervals when it was interrupted for various

reasons. CO 2 

### injection ended on August 7, 2008, but wa ter sample-collection continued through August 14, 2008. During the sampling period, there were several relatively large rainfall events that impacted the chemical composition of the groundwat er and raised the gro undwater levels in the wells (Figure 3.1). On the other hand, groundwater leve ls fell as a result of either drainage or evapotranspiration and water sampling (Figure 3.1).

### 13 1493.94 1493.98 1494.02 1494.06 1494.10 1494.14 1494.18 1494.22 1494.26 1494.30

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Water level (m above mean sea level)

0 1 2 3 4 5 6 7 8 9Cumulative precipitation (cm) 

### 1B

2B

4B

5B

0.3 cm 2.4 cm

0.26 cm 2.8 cm 0.26 cm 0.56 cm 1.6 cm

CO 2 start (7/9 15:45) CO 2 stop (8/7 10:00)

1493.94 1493.98 1494.02 1494.06 1494.10 1494.14 1494.18 1494.22 1494.26 1494.30

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Water level (m above mean sea level)

0 1 2 3 4 5 6 7 8 9Cumulative precipitation (cm) 

### 1B

2B

4B

5B

0.3 cm 2.4 cm

0.26 cm 2.8 cm 0.26 cm 0.56 cm 1.6 cm

1493.94 1493.98 1494.02 1494.06 1494.10 1494.14 1494.18 1494.22 1494.26 1494.30

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Water level (m above mean sea level)

0 1 2 3 4 5 6 7 8 9Cumulative precipitation (cm) 

### 1B

2B

4B

5B

0.3 cm 2.4 cm

0.26 cm 2.8 cm 0.26 cm 0.56 cm 1.6 cm

1493.94 1493.98 1494.02 1494.06 1494.10 1494.14 1494.18 1494.22 1494.26 1494.30

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Water level (m above mean sea level)

0 1 2 3 4 5 6 7 8 9Cumulative precipitation (cm) 

### 1B

2B

4B

5B

0.3 cm 2.4 cm

0.26 cm 2.8 cm 0.26 cm 0.56 cm 1.6 cm

CO 2 start (7/9 15:45) CO 2 stop (8/7 10:00)

Figure 3.1. Cumulative precipitatio n and groundwater levels in selected ZERT wells during the sampling period in

2008.



### 3.1. Dissolved Inorganic Chemicals

The chemical data obtained for samples from shallow and deep wells prior to CO 2 

### injection show that the groundwater in the area is a Ca-Mg-Na-HCO 3 

### type water, with a fresh water salinity of about 600 mg/L TDS (Figure 3.2; Sample Z- 109 in Table A1). The groun dwater has a pH of approximately 7, and HCO 3 

### is the dominant anion. The concentrations of Cl and SO 4 

### are relatively low. The initial concentrations of Fe, Mn, Zn, P b, and other trace metals are at ppb levels, which is typical of shallow groundwaters in regions of moderate precipitation.

Following CO 2 

### injection, the pH of groundw ater decreased systematically to values around 6.0 at wells 1B, 2B, and 5B, all in close proximity to the injection pipe, strongl y responding within one day. The strongest early response oc curred in Well 2B, only 1 m from the injection pipe and in the direction of groundwater flow, decr easing to a pH of 5.7 (Figure 3.3a). The pH of water for samples from Well 4B (6 m from pi pe) started decreasing only after three days following CO 2 

### injection, but pH values remained above pH 6.0 (Figure 3.3a). Th e pH in all wells started increasing towards 6.5

after CO 2 

### injection was terminated on August 7, 2008. The measured pH values of groundwater were controlled primarily by P(CO 2 

### ), which was measured by Stra zisar et al. (2009) in capped wells, and was computed with SOLMINEQ (Kharaka et al., 1988) using the measured temperature, pH, alkalinity, and other chemical parameters. The P(CO 2 

### ) values measured and computed for the ZERT samples were 0.035 bar before CO 2 

### injection; they increased to values close to 1.0 bar following CO 2 

### injection. It is evident from these results that pH is an excellent early indicator for

detection of CO 2 

### intrusion into groundwater.

### 14 100 75 50 25 0 25 50 75 100 

### Well 2B [08-109]

(pre-injection)

100 75 50 25 0 25 50 75 100

pH = 5.96; TDS = 1514 mg/L pH = 7.04; TDS = 614 mg/L



### Cl Cl

SO

4 

### HCO

3 

### HCO

3



### SO

4



### Mg Mg

Ca Na Ca Na



### Well 2B [08-146] (post-injection; T = 339hr)

Figure 3.2. Modified Stiff diagrams showing concentrations (equivalent units normalized to 100%) of major cations and anions, together with salinity and pH of groundwater from well 2B before and during CO 2 injection.



### The alkalinity of groundwater incr eased from about 400 mg/L as HCO 3 

### to values higher than 1,300 mg/L, following CO 2 

### injection (Figure 3.3b). Alka linity values for different wells, as with the pH values, show variable trends reflecting distance from the CO 2 

### injection pipe, impacts from precipitation events, and possibly lo cal variations in the mineral composition of soils and sediments. Given the reasonable assumption that the initial groundwater composition for Well 5B was similar to the average background composition at the site, the alkalinity values for this well and Well 2B increase first and show the highest alkalinities (up to 1330 mg/L). The alkalinity values in Well 4B, which is furthest away from the injection well, increase more slowly and only to about 700 mg/L.

The alkalinities in all wells decrease to values approaching 400 mg/L after termination of CO 2 

### injection (Figure 3.3b).

Values for the electrical conductance, also measured at the site, show simila r trends to alkalinity, increasing from ~600 μS/cm before CO 2 

### injection to approximately 1800 μS/cm following CO 2 

### injection (Figure 3.3c). Assuming that the electrical conductivity of the gr oundwater in Well 5B was initially similar to the aver age background value at the site, th e conductance values for wells 1B, 2B, and 5B increase first and show the highe st conductance. Conductance values for Well 4B increase more slowly and only to about 1,100 μS/ cm. All conductance measur ements decrease to values approaching 600 μS/cm after termination of CO 2 

### injection on August 7, 2008 (Figure 3.3c).

### 15 5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2 7.4

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5pH 1B

2B

4B

5B CO 2 start CO 2 stop

200 400 600 800 1000 1200 1400

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Alkalinity, as HCO3( mg/L ) 1B

2B

4B

5B CO 2 start CO 2 stop

400 600 800 1000 1200 1400 1600 1800

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Electrical conductance (S/cm ) 1B

2B

4B

5B

CO 2 stop CO 2 start



## (c) (b) (a)

5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2 7.4

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5pH 1B

2B

4B

5B CO 2 start CO 2 stop

200 400 600 800 1000 1200 1400

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Alkalinity, as HCO3( mg/L ) 1B

2B

4B

5B CO 2 start CO 2 stop

400 600 800 1000 1200 1400 1600 1800

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Electrical conductance (S/cm ) 1B

2B

4B

5B

CO 2 stop CO 2 start 5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2 7.4

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5pH 1B

2B

4B

5B CO 2 start CO 2 stop

5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2 7.4

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5pH 1B

2B

4B

5B CO 2 start CO 2 stop

200 400 600 800 1000 1200 1400

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Alkalinity, as HCO3( mg/L ) 1B

2B

4B

5B CO 2 start CO 2 stop

200 400 600 800 1000 1200 1400

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Alkalinity, as HCO3( mg/L ) 1B

2B

4B

5B CO 2 start CO 2 stop

400 600 800 1000 1200 1400 1600 1800

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Electrical conductance (S/cm ) 1B

2B

4B

5B

CO 2 stop CO 2 start

400 600 800 1000 1200 1400 1600 1800

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/1 5Electrical conductance (S/cm ) 1B

2B

4B

5B

CO 2 stop CO 2 start



## (c) (b) (a)

Figure 3.3. Field- measured groundwater pH values (a), alkalinities (b), and electrical conductance (c), obtained from selected ZERT wells as a function of time of sa mpling. Note the systematic decrease in pH values from ~7.0 before CO 2 injection to values as low as 5.6 during injection, and subsequent pH increases after CO 2 injection ended. Alkalinities increased from about 400 mg/L before CO 2 injection to values close to 1200 mg/L as HCO 3 , and electrical conductan ce also increased from about 600 μS/cm before

CO 2 injection to values higher than 1,600 μS/cm during injection.



### The alkalinity increases following CO 2 

### injection are balanced primarily by increases in the concentrations of Ca and Mg, wher eas the concentrations of Na (10 ±2 mg/L) are relatively constant (Figure 3.4). The molar and atomic concentrations of Ca increase the most, with atomic values for Ca increasing from about 80 mg/L to up to 240 mg/L, and those for Mg increasing from about 25 mg/L to up to 70 mg/L (Figure 3.4). In Section 4. 2, we report the presence of calcite throughout the sampled sediment column, as well as a small fraction of detrital limestone and dolomite. The presence of these carbonate minerals could e xplain increases of Ca and Mg upon dissolution. However, atomic Ca/Mg ratios obt ained initially and following CO 2 

### injection (Figure 3.5) cannot be

### 16 explained by dissolution of dolom ite alone; the values, increasi ng from 3.2 to 6.1 for Well 2B,

indicate that dissolution of calcite is dominant, but that disso lution of both calcite and dolomite (or Mg-rich calcite) are requir ed to explain the changes in alkalinity and concentrations of Ca and Mg (Figure 3.5, Table A1). Dissolution of calcite (reaction 3-1) and dolomite (reaction 3-2) can be represented by the following reactions:

CO 2(g) 

### + H 2 

### O + CaCO 3(s) 

### = Ca 2+ 

### + 2HCO 3 - 

### (3-1)

2H + 

### + CaMg(CO 3 

### ) 2(s) 

### = Ca 2+ 

### + Mg 2+ 

### + 2HCO 3 - 

### (3-2)

These conclusions are supported by th e initial characterization of minerals in core samples that

show that calcite is abundant in a caliche layer observed at depths of ~50-80 cm. Small amounts of carbonates (~2% detrital limestone and calcite coatings on grains) were also observed in gravels below the caliche layer (Secti on 4.2). Results of geochemi cal modeling with an updated SOLMINEQ (Kharaka et al., 1988) support dissolution of calcite and disordered dolomite as possible reactions at all pH valu es; also, groundwater is undersatura ted with respect to dolomite at the lowered pH values obtained following CO 2 

### injection, allowing for its dissolution. Desorptionion exchange reactions on clay minerals with H + 

### are also suggested as an alternative and/or additional mechanism leading to increases in the concentrations of Ca, Mg, and other dissolved species, as discussed further in Sections 5.2 and 5.3

0 50 100 150 200 250

300 400 500 600 700 800 900 1000 1100 1200
Alkalinity, as HCO 3 (mg/L)Ca, Mg, Na, K Concentration (mg/L)

0.0 0.2 0.4 0.6 0.8 1.0Sr, Ba Concentration (mg/L) Ca
Mg

Na

K

Sr

Ba

Figure 3.4 Concentrations of major cations in groundwater from the ZERT wells plotted as a function of water alkalinities. Note the relatively constant concentrations of Na and K, but the general increases in the concentrations of divalent cations with water alkalinities, possibly indicating dissolution of carbonate

minerals.

### 17 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Ca/Mg (weight ratio) 1B

2B

4B

5B

A wells

CO 2 start CO 2 stop

2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Ca/Mg (weight ratio) 1B

2B

4B

5B

A wells

CO 2 start CO 2 stop

Figure 3.5 Ca/Mg weight ratios plotted as a function of time of sampling. Note that the ratios are higher than the value expected from dissolution of dolomite (1.6), and are relatively constant for well 4B, but continue to increase for well 2B until CO 2 injection was terminated, possibly indicating dissolution of calcite.



### The concentrations of Fe and Mn, the two most a bundant trace metals in gr oundwater, also increase following CO 2 

### injection (Figure 3.6). The concentrations of Fe in gr oundwater could increase from dissolution of Fe(II) and Fe(III) minerals. These pot entially include Fe (hydr)oxides such as ferrihydite and green rust (see Sect ion 5.2), which are likely to be present from the oxidation of primary phases such as magnetite and biotite, obse rved in the sediments (Section 4.2). Dissolution

of siderite would also result in increased Fe concentrations. Ho wever, thermodynamic evaluations (Section 5.2) indicate the groundwater to be quit e undersaturated with respect to this mineral, suggesting it is not likely to be present, except poten tially as a detrital pha se. Example dissolution

reactions include:

Fe(OH) 3(s) 

### + 3H + 

### = Fe 3+ 

### + 3H 2 

### O (3-3)

4FeOOH (s) 

### + 8H + 

### = 4Fe 2+ 

### + 6H 2 

### O +O 2 

### (3-4)

4Fe(OH) 3 

### + CH 2 

### O + 8H+ ==> H 2 

### CO 3 

### + 4Fe +2 

### + 10H 2 

### O (3-5)

Fe(CO 3 

### ) (s) 

### + H + 

### = Fe 2+ 

### + HCO 3 - 

### , (3-6)

where CH 2 

### O is a generic representation of dissolved (and particulate) organic matter. The concentrations of Fe in groundwater are a strong function of Eh, which coul d only be measured in water from the deep wells that were not impacted by CO 2 

### injection, because sh allow wells did not yield enough water volume. The Eh values obta ined (150 to 200 mV; pe ~ 2.5–3.5) indicate suboxic

conditions that account for the low concentration of dissolved Fe; much higher Fe values are possible under more reducing conditions because of the higher solubility of Fe(II) minerals (Hem, 1985). Note that Eh measurements may be of lim ited use, as discussed in Section 5.2. The Fe concentrations (Figure 3.6a) incr ease from 5 to 1200 ppb, but show very low values during July 20 to July 26 following significant rain fall events (Figure 3.1) even when pH values were low. Dilution alone cannot explain the low Fe con centrations during July 20 to July 26, but the low values could be attributed to more oxic conditions, possibly caused by increased dissolved O 2 

### content in

### 18 groundwater transported with perc olating water from rainfall events. Ion exchange reactions on clays with H + 

### and major dissolved cations, such as Ca, Mg, and Na, are other possible controls on Fe concentrations, as discussed in Sections 5.2 and 5.3.

The concentrations of Mn show similar tre nds to those of Fe, increasing from 5 to 1400 ppb following CO 2 

### injection, but also show low values during July 20 to July 26 (Figure 3.6b).

Dissolution and redox reactions can be written for Mn as for Fe above. It should be noted that Mn concentrations are higher for Well 4B than for 2B and 5B, and this may be controlled by local mineral compositions.

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Concentration (mg/L) 1B

2B

4B

5B

A wells

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Concentration (mg/L) 

## (a) Fe

(b) Mn

CO 2 start CO 2 stop CO 2 start CO 2 stop

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Concentration (mg/L) 1B

2B

4B

5B

A wells

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15Concentration (mg/L) 

## (a) Fe

(b) Mn

CO 2 start CO 2 stop

CO 2 start CO 2 stop CO 2 start CO 2 stop

CO 2 start CO 2 stop

Figure 3.6 Concentrations of Fe (a) and Mn (b) in groundwater from selected ZERT wells plotted as a function of time of sampling. Note the low Fe and Mn concentrations during July 20 to July 26, which we are attributing to the oxidizing conditions possibly caused by percolating oxygenated water from rainfall events (see Figure 3.1).



### The concentrations of Pb, As, Zn, and other trace metals (Table A1 and Figure 3.7) generally show an increase with increasing alkalinit y (lower pH value) following CO 2 

### injection. The values reported, however, carry high uncerta inties (10–50% analytical error) as they are, in some cases, close to the analytical detection li mits, especially with the required dilutions (up to 5 fold) prior to analyses. The concentration increases are likely caused by desorption-ion exchange reactions with

H + 

### , Ca, and Mg resulting from lowered pH va lues (see Section 5). Because measured H 2 

### S values were below our detection limit (0.2 mg/L), we are unable to use geochemical modeling to check for solubility of sulfide minerals as controls on me tal concentrations. The concentrations, it should be noted, are all significantly below the maximum cont aminant levels (MCLs) for the respective trace metals (e.g., 15 ppb for Pb, 6 ppb for As). The initial values and the increases in concentrations of these trace metals, although sma ll, are readily measured by the sampling and analytical methods

### 19 used in this study. These results highlight the role of geochemical tools for early detection of CO 2 

### leakage into groundwater.

0.01 0.10 1.00 10.00

300 400 500 600 700 800 900 1000 1100 1200

Alkalinity, as HCO 3 (mg/L)Concentration (g/L )

As Cu Zn Cd Pb Al Se

Figure 3.7 Trace metal concentrations in groundwater from selected ZERT wells plotted as a function of water alkalinities. Some of the scatter shown probably results because the reported values in some cases are

close to the detection limits.



### The chemical changes observed in the ZERT groundw ater are similar in trends, though much lower in concentrations, to the change s observed in the Frio Brine Pilo t test, near Houston, where 1600 t

of CO 2 

### were injected at 1500 m depth into a 24 m th ick “C” sandstone unit of the Frio Formation (Freifeld et al., 2005; Hovor ka et al., 2006; Kharaka et al., 2006). Following CO 2 

### breakthrough, 51 hours after injection, Frio samples showed shar p drops in pH (6.5 to 5.5 measured at ground surface), pronounced increases in alkalinity (100 to 3000 mg/L as HCO 3 

### ) and in Fe (30 to 1100 mg/L), a slug of very high DOC values, and signif icant shifts in the isotopic compositions of H 2 

### O, DIC, and CH 4 

### . These data, coupled with geochemical modeling, indicated rapid dissolution of minerals, especially calcite and iron oxyhydroxides, both caused by lowered pH (initially ~3.0 at subsurface conditions) of the brine in contact with supercritical CO 2 

### (Kharaka et al., 2009b). The

differences between results from Frio and ZER T tests are related to several geochemical parameters, but the most important reason re lates to the difference in subsurface P(CO 2 

### ) value, which was approximately 150 bar at Frio (Kharaka et al., 2009b), while P(CO 2 

### ) values at ZERT measured (Strazisar, 2009; Strazi sar et al., 2009) and computed w ith SOLMINEQ (Kharaka et al., 1988) ranged from 0.035 to ~1 bar. The amount of CO 2 

### dissolved in water is a function of fluid pressure, temperature and salinity. The CO 2 

### solubility increases at hi gher pressures, but decreases at higher (moderate) temperatures an d salinities. In the present case, the large difference in pressure

between the ZERT and Frio sites ha s a stronger effect than differen ces in temperature and salinity between the two sites. Thus, in a shallow aquife r such as that at the ZERT site, much less CO 2 

### can dissolve in the groundwater compared to the amount that can dissolve in deep pressurized formation waters such as at the Frio site. Therefore, the CO 2 

### -induced decrease in groundwater pH will be less pronounced, and thus the corresponding change in dissolved groundwater constituents will also be less. Consequently, trace metal mobilization from solids is expected to be less pronounced at

shallow sites such as ZERT.

### 20 3.2. Dissolved Organic Chemicals

Dissolved organic carbon (DOC) values obtained in this study are generally about 4 mg/L, with a variation range from 2.6 to 6.9 mg/L, and th ese variations do not seem related to CO 2 

### injection. Concentrations of benzene, tolu ene, ethyl-benzene, m-, p- and o- xylene (BTEX) were determined in a relatively large number of samples collected when it became clear that the injection of CO 2 

### was likely responsible for very small but systematic in creases in the concentra tions of dissolved BTEX compounds (up to ~1.6 μg/L for m-, p-xylene), as de picted in Figure 3.8. It s hould be noted that the

concentrations are all below the maximum contam inant levels (MCLs) for the respective BTEX compounds (5 ppb for benzene). There is some scatter in the result s, but the concentrations are about 0.2 μg/L before CO 2 

### injection, increasing to values grea ter than 1.0 μg/L for m-, p-xylene and o-xylene, and to about 1.0 and 0.8 μg/L for toluen e and benzene, respectively. The measured values are higher in water from Wells 1B (Figure 3.8a) and 2B (Figure 3.8b) relative to Well 4B, which is farther away from the injection pipe (Figure 3.8c). All the measured BTEX values decrease to a value approaching pre-injection (~ 0.2 μg/L) after termination of CO 2 

### injection on August 7, 2008 (Figure 3.8d).

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8



### (d) all wells (c) well 4B (b) well 2B (a) well 1BConcentration (g/L )

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8

benzene

toluene ethyl-benzene m-p-xylene o-xyleneConcentration (g/L )

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15 CO 2 start CO 2 stop CO 2 start CO 2 stop

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8



### (d) all wells (c) well 4B (b) well 2B (a) well 1BConcentration (g/L )

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8

benzene

toluene ethyl-benzene m-p-xylene o-xyleneConcentration (g/L )

7/7 7/10 7/13 7/16 7/19 7/22 7/25 7/28 7/31 8/3 8/6 8/9 8/12 8/15 CO 2 start CO 2 stop CO 2 start CO 2 stop

Figure 3.8 Concentrations of BTEX compounds detected in groundwater from selected ZERT wells plotted as a function of time of sampling. The BTEX concentrations are all significantly below the maximum contaminant levels (MCLs) (e.g., 5 ppb for benzene).



### The origin of the low amounts of BTEX com pounds detected in groundwater following CO 2 

### injection was investigated during 2008, and more recently in 2009. Three main possible sources were examined: (1) A natural source from deep organic matter or a petroleum source; (2)

### 21 contamination from a surface source or when inst alling the injection pipe and groundwater wells; or (3) contamination with trace amounts of BTEX in the injected CO 2 

### , although the gas used was foodgrade quality. Our initial re sults obtained in 2008 we re inconclusive with regard to the injected CO 2 

### being the source of the observed BTEX. We analy zed core samples obtained from wells drilled in December 2008 for BTEX. Though progress was slow b ecause results obtained were close to the detection limits of available analytical methods, results showed that BTEX concentrations were below detection limits. Thus, natural sources were eliminated as a possibility for the observed concentration increases. Purg e-and-trap analysis of CO 2 

### samples collected from the injection source tank in 2009 showed conclusively that the injected CO 2 

### (though food-grade) contained a very small amount of organic chemicals and thus was identifi ed as the source of the trace concentrations of detected BTEX (Kharaka et al., 2009b). Our results indicate that BTEX compounds can be detected in groundwater at extremely low concentrations, below MCL levels.



## 4. MINERALOGICAL DATA

4.1. Core Analysis and Methods

Three wells were installed in December 2008 to obtain core for mineral characterization and laboratory experiments. Well W6 is 2.6 m deep with the bottom 0.76 m screened. Wells W7 and W8 are 3 m deep with the bottom 1.5 m of the wells screened (Figur e 2.2). Cores from these wells have been logged and sampled at different in tervals for detailed mineralogical analysis. Approximately 20 samples have been collected fro m the cores for rock, mineral, and grain-size analysis, and 17 samples of these have been dry-si eved to 3 size fractions with grain-size diameters of (1) >9.5 mm (medium gravel +); (2) 9.5–2.0 mm (f ine to very fine gravel), and (3) < 2 mm (coarse sand to clay/colloid). X-ray diffraction (XRD) work was co mpleted with the primary goal of identifying carbonates and clays. Short exploratory sessions were conducted with a scanning electronic microscope (SEM), including some semi-quantitative chemical analysis using the Energy-Dispersive X-Ray (EDX) system on the SEM. The short summary given below focuses on ZERT Well W6, which is located in the main well fi eld (Figure 2.2). A log for well W6 is given in Table B.1 in Appendix B. The grain size fraction fo r different zones in Well W6 is shown in Figure

4.1.

It should be noted that the co res were obtained by a “direct-p ush” geoprobe, using a 3-inch geoprobe system. The direct-push methodology creates a bias in the true grain-size distribution because the larger cobbles, which are abundant th roughout the cored interval (even the top “A” soil horizon), are simply pushed out of the way. Larg e well-rounded cobbles exceeding 10 cm diameter

exist in abundance throu ghout the cored interval.

### 22 ZERT - Well 6 (W6) grain-size distribution

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

123 4567% of size fraction (normalized to 100%

< 2mm

2-9.5 mm

> 9.5 mm

“A” soil

20-25 cm “caliche”

50-55 cm gravel

251-261 cm gravel

177-184 cm gravel

245-251 cm gravel

61-64 cm gravel

107-118 cm ZERT - Well 6 (W6) grain-size distribution

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

123 4567% of size fraction (normalized to 100%

< 2mm

2-9.5 mm

> 9.5 mm

“A” soil

20-25 cm “caliche”

50-55 cm gravel

251-261 cm gravel

177-184 cm gravel

245-251 cm gravel

61-64 cm gravel

107-118 cm

Figure 4.1 Grain size fractions for different zones in Well W6.



### 4.2. Mineralogical Results

The cores have an organic-rich dark soil zone at the top (0 to ~40 cm ), which grades into a light to medium gray “caliche” zone. The caliche zone is high in carbonate and clay, with lesser fine sand to fine silt (predominantly quartz, feldspar and white mica, and very minor volcanic rock fragments). Clay extracts from the caliche zone of W6 appear to be smectite. The carbona te mineral is calcite; however, the XRD pattern shows it is poorly crystallized, with a peak shift towards Mg-rich calcite.

“A” soil horizon

The top 30-40 cm of each core consists of clay and organic-rich soil. The soil is generally brown, plastic, and matrix supported with some fine sand. The topsoil zone grades into a caliche layer that is ~20-40 cm thick. The mineralogy has not been st udied in detail. XRD analysis performed on the clays from earlier auger core samples indicated chlorite; however, our stu dy indicates the caliche clay is dominantly smectite. Some well-rounded pe bbles, possibly volcanic andesite, are found in the top soil.

Caliche or high carbonate zone

The caliche zone begins at ~40 cm , graded gently in color from da rk brown topsoil to medium and light gray. The caliche zone is very fine-grain ed, with most material passing easily through 2 mm sieve (Figure 4.1). Unlike the caliche from a hand-a ugered core taken previous ly from other nearby wells (located at gridpoint -5,- 3 on the reference grid), which wa s predominantly carbonate with lesser amounts of clay and fine sa nd, the caliche zone at Well 6 app ears to be more clay and quartzdominated fine sand and silt, and less carbonate . The XRD scan of a bulk powder sample shows quartz and clay (smectite), with significant calc ite and plagioclase, and minor but measureable kaolinite and mica. XRD results suggest that the main clay mineral is smectite (main 001 peak ~14

### 23 angstroms, with a shift to ~17 angstroms upon glyc olation, and a slightly higher peak shift using glycerol). More work would be needed to desc ribe the clays in more detail. An SEM stub was prepared with the caliche materi al; however, the carbonate and cl ay was so fine-grained and comingled that it was not possible to resolve the mor phology of the crystals or to isolate grains for elemental analysis. Mica is readily identified in si eved samples from the caliche; it is predominantly white (muscovite) and lesser green mica. Mafic rock fragments are a very minor constituent of the

caliche zone.

Gravel zone

Below the caliche zone, the core changes rapidly to coarse gravel, with large pebbles up to the full diameter of the core (~6 cm), and gravel (> 2 mm) representing ~75% or more of the core. The gravel continues to the bottom of each 10 ft deep well. Megascopically, there are no obvious distinct zones. A total of 15 samples, ranging fr om ~150 to 400 grams, were collected from the gravel zones of the three wells (5 samples from W6; 4 samples from W7; 6 samples from W8). Because the gravel was so coarse, the entire diam eter of the core was collected for each sample, from core intervals of ~5 to 10 cm. The grain size fractions for all 15 gravel samples analyzed from wells 6, 7, and 8 are plotted in Figure 4.2.

ZERT - W6, W7, W8 "gravel zones" - grain-size distribution

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15% of size fraction (normalized to 100%)

< 2mm

2-9.5 mm

> 9.5 mm

Figure 4.2 Grain size fractions for 15 gravel samples analyzed from wells 6, 7, and 8.



### Overall, the gravel material can best be described as “extremely heterogeneous,” and it would be a significant task to identify all the rock types and minerals in this gravel. An approximate distribution of rock types in the gravel is presented below, with results from W6:

 Approximately 60%–70% mafic/ba sic volcanic rocks, likely andesite (mostly aphantic, some small plagioclase laths). Usually subangul ar to subrounded habit, often with broken edges. Some volcanic fragments have a pale bl ue “oxidized” exterior , but when broken are

fresh. XRD shows dominant mineral is Ca-Na plagioclase (albite to anorthite). Semi-

### 24 quantitative SEM/EDX “spot scans” of several feldspars showed a wide range of Na to Ca. Include basalt and diabase in this group. O livine was identified in some rocks using

SEM/EDX.

 ~20-30% quartz –biotite -+/-amphibole gneiss. Hab it angular to subangular.

 ~10% misc crystalline rocks, such as red gran ite/gneiss (with dark red potassium feldspar). Sub-angular. Likely rhyolite (pink-orange, subrounded), quartz-mica schist (angular).

 ~2% limestone—described below.

In general, the darker, denser aphanitic volcanic ro cks dominate in the larger size fractions, often as well-rounded pebbles (up to 30 grams weight). A wi de range plethora of ro ck types exist in the smaller size fractions (large sand to fine gravel size)—diabase, schist s, granites, vesicular volcanics, rhyolite, and many exotic crystalline rocks, ranging from highly altered/ weathered to fresh—that will take considerable some effort to identify.

Most of the mafic volcanic rocks are highly magne tic, so that even rock fragments > 2 mm are easily separated with a 1-ounce horseshoe magnet. Th e XRD pattern of a separate of fine-grained magnetic material was determined to be magnetite. SEM/EDX spectra of several “Fe-rich” minerals in andesite showed iron and titanium (up to 7 atomic %).

Limestone (in gravel)

Detrital limestone (not found in the core-catcher samples) is present in the gravel zone of Well 6, as a very small fraction of the total gravel. Of 97 hand-checked pebbles (~1 cm+ diameter) from the five W6 gravel samples, only 3 pebbles are li mestone. Considered by weight, the 3 limestone pebbles total only 10.2 grams, from a total weight of 573 grams of is this size fraction. The

limestone has a resemblance in color and textur e to the Madison limestone; however, it appears even more fine-grained than the hand specimen we have of Madison Limestone. The conchoidal fracture and XRD pattern shows it to be well-crystallized calcite. Limestone is also present in the 2– 9.5 mm size-fraction, again in very small amounts (~2%).

Carbonate

Calcite, primarily as a coating, is a common mino r constituent throughout the gravel zone. Calcite also commonly occurs as crusts of secondary calcite on rounded rock surfaces, from small to medium pebbles. An XRD pattern of such a carbon ate crust was determined to be well-crystallized

calcite. All 94 noncarbonate pebbles from the > 9.5 mm size frac tion had a carbonate coating and/or crust and reacted slightly-to-vigorous ly to dilute HCl. Likewise, all of the matrix samples (sieved and unsaved) reacted to HCl. Calcite also occurs throughout the matr ix as very fine rounded grains,

which dissolve in a few seconds in dilute HCl. We find no dolomite in the coarse fractions, although the 2–9.5 mm fraction would need further study. The finer size fractions likely contain some dolomite, identified as small rounded grains that fizz very slowly in 10% HCl (as opposed to calcite, which fizzes rapidly. Because the dolomite fractions are too small to appear in bulk XRD powders (generally ~5% by weight is needed), SE M or hand-picked grains would be necessary for positive identification.

### 25 Matrix minerals

In the progressively finer size fractions, minerals predominate over rock fragments and lighter minerals predominate over darker minerals. For minerals, quartz is dominant, followed by feldspar. Micas are abundant in all the size fractions, with muscovite more abundant in the finer sizes compared to biotite. Green and brown mica are also present (possibly chlori te, phlogopite). For the darker minerals, amphibole appears most abundan t; pyroxene is not easily spotted in matrix, although common in some rock fragments. Magneti te is abundant. Carbona te, as a coating, and calcite as very small rounded grains are found throughout the grav el zone. Fe-oxides, as yet unidentified, are present in the matrix, both as fine coatings on other minerals (often quartz) and as small rounded and powdery grains. Some pistachio green epidote was observed on a pebble surface.

Sulfides

No sulfide minerals have been identified in the ma trix. Pyrite is rare in some rock fragments, typically on fresh surfaces. The SEM backscatter sy stem was used on many rock grains to recognize sulfide minerals, but none were found in our samp les. High Fe targets were usually found to be oxides, magnetite, or titaniferous magnetite.



## 5. HYDROGEOLOGICAL AND GEOC HEMICAL INTERPRETATION

The groundwater data collected to study the effect of CO 2 

### injection on groundwater quality at the ZERT site was presented in Section 3. Here, th e groundwater data collect ed during this study are evaluated in a more quantitative manner, in li ght of the site hydroge ology, its mineralogical characterization, and the various geochemi cal processes likely to result from CO 2 

### dissolution in

water.

It should be remembered that th is study was conducted in a sha llow water-table aquifer, only a couple of meters below ground surface, at esse ntially atmospheric pr essure, and thus under conditions that are not representative of deeper USDW. In addition, becau se of the thin vadose zone, recharge from rainfall events during the test (Figure 3.1) had a rapid and noticeable effect on groundwater composition trends, complicating data in terpretation. Nevertheless, geochemical data collected during this study are qui te useful to test some hypothe ses regarding the effect of CO 2 

### dissolution in groundwater, namely that the pH d ecrease and alkalinity increase accompanying CO 2 

### dissolution in water are likely to drive dissolutio n and desorption/exchange reactions that could impact groundwater quality. These reactions would be expected to take place in any aquifer, although to a lesser extent in a shallow site such as ZERT co mpared to a deep USDW. This is because, as mentioned earlier, less CO 2 

### can dissolve in water at low pr essure than at high pressure, thus limiting the pH decrease and alkalinity in crease (e.g., reaction 1- 1) accompanying CO 2 

### injection. Obviously, the geochemical response of an aquifer to CO 2 

### dissolution also highly depends on the aquifer mineralogy, salinity, redox conditions , and initial (background) pH. Because these data are site-specific, results of analyses presente d in this section should not be directly applied to other sites, except for illustrating the general type of reactions (and not their specific effect or magnitude) that could take place at other locations.

### 26 Section 5.1 summarizes wh at is known about CO 2 

### transport in the subsurface at the ZERT site, to provide context for evaluating th e water quality response to CO 2 

### injection. In Sect ion 5.2, statistical and thermodynamic analyses are presented to determ ine correlations between the concentrations of various species, and formulate hypotheses on the types of mechanisms likely responsible for observed concentration trends and groupings of certa in species. Based in pa rt on this information, geochemical modeling analyses are then presen ted in Section 5.3. The modeling incorporates processes (i.e., dissolution, sorption/desorption, and ion exchange) that are most likely to affect groundwater quality upon CO 2 

### injection. Besides providing possi ble explanations for observed water quality changes at the ZERT site, the modeling also illustrates the types of coupled geochemical mechanisms that could affect gr oundwater quality in deeper USDW.

5.1. Transport of CO 2 

### in Gas and Aqueous Phases

The transport of CO 2 

### in the subsurface at the ZERT site (e.g., as gas phase in the vadose zone and dissolved in groundwater) is assessed in other parallel studie s (Lewicki et al., 2007; Strazisar et al., 2009; Oldenburg et al., 2009). As discussed in Section 2.3, groundwater flow velocities at the site may be up to 2 m/d, which is very high, and rates of CO 2 

### transport in the unsaturated zone are estimated to be at least an order of magnitude hi gher. Therefore, the rapid and systematic changes observed in the chemical composition of groundwat er at ZERT are likely driven by both the fast spreading of CO 2 

### as gas phase in the sand-gravel unsat urated zone under the upper soil layer, followed by re-dissolution into the groundwater, as well as by transport of groundwater with

dissolved CO 2 

### (see schematic in Figure 5.1). Notice that in both transport scenarios, the deeper A-wells are not significantly affected by the injected CO 2 

### . Redissolution of gaseous CO 2 

### is limited to a small layer at the soil-groundwat er interface, while the flow of CO 2 

### -charged water from the injection well bypasses the A-wells, because the screening interval is too deep.

The fast transport in the gas pha se is directly supported by gas samples obtained from the head space of the groundwater wells. CO 2 

### concentrations in wells 1B, 2B , and 5B increased up to about 90% (by volume) only two days after injection st arted (Strazisar, 2009). Th ese wells are all less than 2 m away from the injection pipe, with 5B actually located upstream (Strazisar et al., 2009).

Well 3B about 3 m downstream arrives at 50% CO 2 

### after about 5 days; Well 4B about 6 m

downstream starts to res pond after the same time period. Model re sults from Oldenburg et al. (2009) also indicate that gaseous CO 2 

### spreads out fast through the vadose zone under the soil layer. It is thus quite certain th at dissolution of CO 2 

### from the vadose zone b ack into groundwater is a prominent process triggering the chemical changes measured in the groundwater samples. However, because the groundwater flow velocity is very high (up to 2 m/day), the tr ansport of dissolved CO 2 

### in groundwater cannot be ruled out as an additional main source of CO 2 

### responsible for the groundwater quality changes, at l east in the shallower wells located downstream from the injection pipe. Because uncertainties remain regarding gas a nd groundwater flow rates, recharge effects, and subsurface heterogeneity, no attempts were made , here, at modeling the reactive transport of CO 2 

### in the subsurface. We focus instead solely on geoche mical modeling of reactive processes, such that the transport specifics re sponsible for increased CO 2 

### levels at monitored locations are not needed.

### 27 Figure 5.1 Schematic representation of CO 2 transport in ZERT field site. The cr oss-section shown is along the axis of well pairs 1A/B, 2A/B, 3A/B, and 4A/B.



### 5.2. Geochemical Response to CO 2 

### Release and Interpretation

The characterization of the geochemical response of groundwaters in the sh allow phreatic zone at the ZERT site was undertaken in response to the opportunity afforded by an independent field program to evaluate phenomena associat ed with the atmospheric leakage of CO 2 

### from underground storage sites. Both time and funding constrai nts limited the scope of the study, which was recognized at the outset as be ing exploratory in nature, and aimed at defining the relevant geochemical processes relating to CO 2 

### contamination of subsurface potable water sources, rather than their quantification. This se ction discusses severa l steps undertaken to interpret chemical analyses of groundwaters from the ZERT site, in an effort to provide a framework for the geochemical modeling conducted in Section 5.3.

The large number of species analyzed in ground waters from the ZERT site, and their temporal variation, provides a basis for understanding the tr ansient chemical processes operating in the phreatic zone near the water table and the ove rlying vadose zone. These processes include the following:

 Dissolution of CO 2 

### in the groundwater

 Evaporative concentration

 Infiltration of rainfall

 Bacterial mineralization of orga nic matter with generation of CO 2



###  Oxidation/reduction

 Ion exchange

### 28  Adsorption

 Mineral precipitation/dissolution

The depression of pH due to CO 2 

### uptake in the groundwater can i nduce dissolution of a number of primary and secondary minerals, and drive sorpti on/desorption and ion exchange reactions. The bicarbonate ion generated during CO 2 

### dissolution can also compete with other species, such as the arsenate anion for adsorption. Fluctu ations in the concentrations of oxidants and reductants, and the potential inhibition of bacterial activ ity by increasing concentrations of CO 2 

### , can likewise affect precipitation, ion exchange, and ad sorption of those species that ar e subject to changes in oxidation state. The evaporation induced by the bubbling of CO 2 

### through the groundwater, with concomitant

removal of dissolved redox-se nsitive gaseous species (e.g., O 2 

### , N 2 

### , CH 4 

### , and H 2 

### S), may disturb the

oxidation state and also lead to evaporative concentra tion of nonvolatile species . Random periods of rainfall are likely to flush re sidual capillary waters from the overlying soil horizons of the vadose zone, which have been subject to evapotranspiration and modification by pedogenic processes.

These vadose zone soil waters al so contain the concentrated resi dues of dissolved species from earlier rainfall events. Bacterial oxidation of orga nic matter transported from the overlying soil zone in the dissolved or particulate state modifies the concentration of disso lved concentration of carbonate species in the groundw ater and leads to elevated partial pressures of CO 2 

### in the gaseous phase in the overlying vadose zone. Such bacterial activity also modifies the concentrations of nutrients in the groundwater and generates reductants that can likewise induce the reductive dissolution or precipitation of minerals in the saturated zone. I on exchange can take place in more than one mineral. Adsorption characteristics al so vary, depending on the nature of the mineral substrates, of which there may be several.

The information available to help identify transi ent chemical processes in the saturated zone consists of some 60 comprehensive chemical anal yses of groundwaters take n at intervals throughout the test, the temperature and el ectrical conductivity of the gr oundwaters samples, information on rainfall events, the height of the water table, and mineralogical analyses of core samples taken from a number of shallow boreholes (see Sections 2 to 4). In addition, data on the chemical composition

of rainwater collected at two sites in Montana were downloaded from US EPA’s National Pollutant Discharge Elimination System (NPDES) web site to estimate contributions of dissolved species to the groundwater from that source. These data are evaluated below usi ng Principal Component Analyses (PCA) to determine whether certain sp ecies correlate with e ach other, potentially indicating similar or coupled geochemical behavi or, and to shed light on possible geochemical mechanisms at play upon CO 2 

### injection. Because the geochemical behavior of many metal species in groundwater is tied to oxida tion state, and because uncertainties remain concerning the controlling redox processes within the saturated zone at the ZERT site, thermodynamic analyses are also presented to help constrain the identity of potential phases exercising thermodynamic control over the redox potenti al at this site.

5.2.1 Principal Component Analysis

Principal Component Analysis (PCA) is a statistical tool for the inte rpretation of datasets consisting of multiple observations in each of many samples, and where it is desired to determine linear

relations between the observations in the dataset. The aim is to reduce the observed relations among the variables to simpler relations between fewe r variables, called principal components. These

### 29 principal components are linear comb inations of the variables, whic h are not correlated, but which explain the total vari ance of the data.

In evaluating the groundwater analyses from the ZERT site, PCA does not identify what each principal component relates to in terms of the chemical processes taki ng place in the field. However, it does identify which chemical species are strongly correlated with a given principal component, and this in itself provides a basis for hypothesizing the relationship of each principal component to a chemical process that is likely to be operative. Unequivocal assignment of any specified chemical process to each of the deri ved principal components is plausible only in conjunction with an independent evaluation of th e mineralogy, interpreta tion of measured soil chemical and physical propertie s, and independent thermodyna mic and kinetic analyses. The interpretation can be further strengthened th rough an independent comparison of modeled predictions of groundwater composit ion with field observations, as is illustrated in a subsequent section of this report.

Shlens (2009) points out that ce rtain assumptions are implicit in PCA for it to be effective:

1. Relations between the variables should be linear, although nonlinear regimes have been explored.

2. It is assumed that large variances have important structure. In other words, the data have a high signal-to-noise ratio, and the principal components with larger associated variances represent interesting structure.

3. The principal components are orthogonal. This assumption makes PCA solvable using linear algebra decomposition techniques.

Treatment of the data

The underlying relations between the various analy zed chemical species in the groundwaters are likely to obey thermodynamic or kinetic mass action re lationships, which are de fined in terms of the logarithms of the activities of the independent components. For a truly rigorous analysis, these

activities would have to be calcu lated using a distribution of speci es code such as EQ3/6 (Wolery, 1993) or PHREEQC (Parkhurst a nd Appelo, 1999), taking into account species complexation and species activity coefficients. However, the suppo rting thermodynamic databases of such codes are typically incomplete, especially regarding comple xation by trace elements, which are of particular interest to this study. Furtherm ore, since it is unlikely that the sampled groundwaters were in homogeneous equilibrium with resp ect to redox state, as is discu ssed further in Section 5.2.2, it is not possible to ascertain with a ny degree of confidence the likely oxidation stat e of several of the trace elements or their complexing agents. For these reasons, calculation of the thermodynamic activities of the aqueous basic sp ecies was not warranted. For the purpose of PCA, the measured species concentrations were ther efore transformed to the logari thmic molar concentrations as proxies for the corresponding activities. The loss of accuracy introduced by this approximation is mitigated somewhat by the sampled groundwaters bei ng very dilute (total dissolved solids (TDS) 

 1,500 mg/L). Furthermore, since the activity coefficients of all charged species vary systematically within the limited range of ionic st rength at the ZERT site, the ratio of the activity coefficients of any two species is approximately constant.

A preliminary study of the behaviors of individual chemical constituents indicated that the sampling period should be subdivided into discrete time in tervals for evaluation and interpretation, because

### 30 different chemical processes became pre-eminent over time, as atmospheric influences changed. Here, we restrict our analysis to a data set consis ting of the period utilized to establish the baseline ambient chemistry of the groundwater (July 7–8, 2008), and the initial period following the start of

CO 2 

### injection up to the time when a significant ra infall event perturbed the water table elevation (July 8–17, 2008). This time period allows evaluati on of the short-term response of groundwater chemistry to increasing partial pressure of CO 2 

### , P(CO 2 

### ). Because of the broad similarities in the chemical responses of the groundwaters to CO 2 

### injection during this peri od, regardless of location of the sampled wells, analyses from all well s were combined for the purpose of PCA.

The individual datasets were ev aluated using the StatisticXLTM p ackage, which is designed for use with the Microsoft ExcelTM spreadsheet softwa re under the Windows operating system. The number of principal components was not co nstrained, as prior knowledge of th e geochemistry of the system suggested that many geochemical processes could be operative, and that independent processes affecting different mineral substr ates might condition the initial response of the system to CO 2 

### uptake.

Results

Table 5.1 displays the correlation matrix of the log molar concentrations of all analyzed species.

The correlation coefficients are color-code d in ranges 0.9, 0.8, 0.7, 0.6, 0.5 and 0.4. Very strong correlations exist between the alkali earth metals, Mg, Ca, Sr, Ba, HCO 3 

### , Al, and to a slightly lesser

extent with SiO 2 

### . Li, Cd, Pb, Al, Cr, SiO 2 

### and B are also strongly correlated. Fe and Co are strongly, but inversely correlated with PO 4 

### . Se is strongly correlated with Cr , whereas Mn is fairly strongly correlated with Co and moderately with Fe. Several species show notably weak correlations with other constituents, notably Zn, U, Cl, Br, and NO 3 

### . H (-pH) is most str ongly correlated with K (0.856), but is otherwise only modera tely correlated with most othe r constituents. Although some of

the observed correlations could be fortuitous , many probably reflect underlying thermodynamic, kinetic, stoichiometric, or mixing relations . These may involve complexing, exchange, or adsorption or, with negative correlations, precipitation. Figure 5.2 illustrates some examples of

relevant correlations between va rious species in solution, which are intended to show that these results might ultimately be used to initiate broader lines of inquir y regarding the behavior of trace metals in response to system perturbations caused by CO 2 

### injection.31Table 5.1Correlation matrix of the log molar concentrations of all speciesanalyzed in MSU-ZERT groundwaters for the period July 8 to July 17, 2008H Li Na K Mg Ca Sr Ba Mn Fe Co Cu Zn Cd PbH 1.000Li0.677 1.000Na0.6450.822 1.000K0.8560.5800.524 1.000Mg0.7930.8500.9020.736 1.000Ca0.8500.8130.8510.8460.974 1.000Sr0.8230.8280.8930.7900.9920.992 1.000Ba0.8300.7620.7990.8350.9590.9740.972 1.000Mn 0.160 0.438 0.333 0.188 0.329 0.257 0.290 0.327 1.000Fe0.5540.7290.7230.5290.8140.7670.7940.8290.645 1.000Co0.4110.6650.595 0.3480.6200.5390.5840.6190.8890.877 1.000Cu0.5430.8720.8330.5420.7860.7690.7810.690 0.2540.6370.465 1.000Zn 0.067 0.347 0.289 0.099 0.262 0.216 0.243 0.219 0.118 0.143 0.225 0.284 1.000Cd0.7280.9510.7050.6430.7700.7610.7600.7220.5000.6890.6810.803 0.265 1.000Pb0.6120.9270.6760.5710.7040.6850.6900.6270.4760.6200.6240.828 0.2550.967 1.000Al0.7750.9530.8430.7050.9310.9160.9200.883 0.3460.7790.6150.878 0.2830.9130.867Cr0.6650.9440.6570.5760.6980.6870.6840.6470.4090.6060.5960.802 0.3030.9650.950As 0.0990.4740.461 -0.015 0.320 0.241 0.282 0.180 -0.116 0.097 0.0600.449 0.379 0.267 0.318Se0.5250.9180.7200.4130.6950.6340.6590.596 0.3360.5580.5440.789 0.3950.8360.843Mo 0.2250.6330.594 0.2140.478 0.3960.444 0.356 0.3880.4430.4700.693 -0.0260.6080.729U 0.3980.4300.652 0.2880.5920.5630.6090.542 0.0020.449 0.2930.501 0.003 0.325 0.301F-0.491 -0.039 -0.070-0.559 -0.332 -0.432 -0.375-0.490 0.043 -0.381 -0.122 0.011 0.376 -0.084 0.063Cl 0.217 0.253 0.099 0.367 0.306 0.246 0.254 0.3290.458 0.364 0.375 0.262 -0.041 0.305 0.333Br -0.058 0.310 0.260 -0.018 0.214 0.149 0.169 0.082 0.103 0.098 0.0670.518 0.181 0.255 0.259NO3 -0.325-0.415 -0.406 -0.096 -0.345 -0.232 -0.293 -0.316-0.636-0.473-0.670 -0.229 -0.181-0.440-0.409PO4-0.479-0.697-0.668-0.431-0.722-0.665-0.696-0.740-0.758-0.950-0.912-0.569 -0.152-0.678-0.596SO40.4480.4920.4090.4510.6490.5520.5860.621 0.3510.5870.490 0.389 0.0400.4610.468HCO30.8060.8280.8870.7630.9900.9830.9930.965 0.2910.8040.5940.772 0.2550.7520.680SiO20.7530.9630.8600.6610.9090.8790.8930.847 0.3460.7530.6250.868 0.3540.8890.855B0.7270.9110.6930.6160.7910.7470.7580.7260.4710.6620.6420.751 0.2640.9390.926LEGEND> 0.9> 0.8> 0.7> 0.6> 0.5> 0.432Al 27 Cr 52 As 75 Se 82 Mo 98 U 238 F Cl Br NO3 PO4 SO4 HCO3 SiO2 B 11Al 1.000Cr0.885 1.000As 0.347 0.414 1.000Se0.8210.9100.727 1.000Mo0.5730.634 0.3210.629 1.000U 0.462 0.242 0.234 0.302 0.411 1.000F -0.204 0.009 0.293 0.127 0.322 -0.152 1.000Cl 0.299 0.282 -0.120 0.228 0.387 -0.077 -0.114 1.000Br 0.265 0.239 0.261 0.307 0.272 0.150 0.200 0.313 1.000NO3 -0.338 -0.427 -0.126 -0.463 -0.378 -0.059 -0.149 -0.419 -0.004 1.000PO4-0.715-0.603 -0.086-0.567 -0.417 -0.318 0.264 -0.367 -0.1410.615 1.000SO40.573 0.412 0.100 0.424 0.463 0.238 -0.2170.795 0.168 -0.426-0.506 1.000HCO30.9190.681 0.2750.653 0.4460.603 -0.362 0.260 0.179 -0.296-0.6980.608 1.000SiO20.9670.9020.5150.8990.595 0.449 -0.134 0.285 0.239 -0.404-0.6920.5560.894 1.000B0.9060.916 0.3070.8370.617 0.233 -0.032 0.459 0.219-0.536-0.6540.6560.7580.888 1.000LEGEND> 0.9> 0.8> 0.7> 0.6> 0.5> 0.4

### 33 y = 1.0102x - 0.7914

R 2 = 0.9795

-3.2 -3.1 -3.0 -2.9 -2.8 -2.7 -2.6 -2.5 -2.4

-2.4 -2.2 -2.0 -1.8 -1.6

Log (HCO3-) (mol/L)Log (Mg) (mol/L) y = 0.8888x - 0.7204

R 2 = 0.9656

-2.8 -2.7 -2.6 -2.5 -2.4 -2.3 -2.2

-2.4 -2.2 -2.0 -1.8 -1.6

Log (HCO3-) (mol/L)Log (Ca) (mol/L)



### (a) (b)

y = 0.9334x - 3.4564

R

2 = 0.9855

-5.6 -5.5 -5.4 -5.3 -5.2 -5.1 -5.0

-2.4 -2.2 -2.0 -1.8 -1.6

Log (HCO3-) (mol/L)Log (Sr) (mol/L) y = 0.9381x - 4.0444

R 2 = 0.9311

-6.2 -6.1 -6.0 -5.9 -5.8 -5.7 -5.6

-2.4 -2.2 -2.0 -1.8 -1.6

Log (HCO3-) (mol/L)Log (Ba) (mol/L)



### (c) (d)

y = -1.5063x - 15.425

R 2 = 0.903

-7.0 -6.5 -6.0 -5.5 -5.0 -4.5

-7.0 -6.8 -6.6 -6.4 -6.2 -6.0 -5.8 -5.6

Log (PO4) (mol/L)Log (Fe) (mol/L)

y = -0.7856x - 12.793

R 2 = 0.8318

-8.6 -8.4 -8.2 -8.0 -7.8 -7.6 -7.4 -7.2

-7.0 -6.8 -6.6 -6.4 -6.2 -6.0 -5.8 -5.6

Log (PO4) (mol/L)Log (Co) (mol/L)



### (e) (f)

Figure 5.2. Correlation plots showing relationships between various species in solution

### 34 Figures 5.2a–d represent plots of th e concentrations of the alkali earth metals with respect to alkalinity (reported as HCO 3 - 

### ), which shows a much stronger correlation than with H + 

### . Minimal scatter characterizes all four plots, even for Ba, which is present in micro-molar concentrations, and all regressions show linear correlations with slop es near unity. The trendlines of all alkali earth metal concentrations with respect to each other and with respect to alkalinity are also linear and extrapolate through the origin, suggesting a common source for all four metals. The correlations exhibited in Figures 5.2a–d cannot be attributed to complexing of the respective alkali earths by

HCO 3 - 

### , because an independent thermodynamic analysis demonstrates that the MHCO 3 - 

### complex (M = Mg, Ca, Sr, Ba) is present only as a minor species, the predominant species being M +2



### . Equilibration with respect to a carbonate mineral according to the reaction

MCO 3 

### + CO 2 

### (aq) + H 2 

### O  M +2 

### + 2HCO 3 - 

### (5-1)

also cannot explain the unit slope positiv e correlation of log M with log HCO 3 - 

### . Furthermore, the linear correlation between Mg and Ca yields a molar ratio, Mg/Ca of approximately 0.45, which, even for the most magnesian rich calcite, seems to be excessive (maximum observed values are approximately 0.25), and would imply that the observ ed secondary calcite pr ecipitates at the MSUZERT site are proto-dolomites (Busenberg and Plummer, 1989). The perturbation of carbonate solubility according to the reaction

MCO 3 

### + H + 

###  M +2 

### + HCO 3 - 

### (5-2)

is also not supported by the data, because correl ations between all log M and pH are poor and nonstoichiometric. This is in part explaine d by the buffering effect that dissolved CO 2 

### has on pH according to the reaction:

CO 2 

### (aq) + H 2 

### O  HCO 3 - 

### + H + 

### ; pK ≈ 6.5 @ 10  C (5-3)

Finally, a comparison with Figure 5.5 (see Section 5.2.3 belo w) shows that sign ificant resaturation of calcite started taking place only afte r the partial pressure of dissolved CO 2 

### exceeded 0.2 bar, which implies that calcite may not have been active ly dissolving until that partial pressure had been exceeded. The linear trend observed between log (Ca +2 

### ) and log (HCO 3 - 

### ) in Figure 5.2b suggests

that Ca +2 

### was initially released from some source other than calcite, and that only later in the time period under investigation did calci te dissolution contribute significan tly to the total dissolved Ca +2


### in solution. This hypothesis is consistent with the argument that the increase d acidity resulting from

the dissociation of CO 2 

### (aq) would depress the chemical poten tial of the CaO component in solution (defined in terms of log [Ca +2 

### ]/[H + 

### ] 2 

### ), and that this would have induced the release of Ca +2 

### , most probably from cation exchange sites, and conc urrently through the dissolution of calcite.

The short-term and immediate resp onse of the system to changes in alkalinity as demonstrated by these correlation plots can be best explained by ion exchange on the basis of the following arguments: (a) the strong correlation observed in Figures 5.2a–d and the inferred rapid response of the system to increasing alkalini ty; (b) the linear relations between the log concentrations of the

alkali earths and other divalent cations, suggesting a common sour ce for all divalent cations, and

even monovalent and trivalent cat ions; (c) the inferen ces drawn from the principal component analysis described below; (d) the close correl ation of the geochemical modeling described in

### 35 Section 5.3 in which an ion exchange model, usin g independently derived ion exchange parameters, closely matches field measurements of pH against an alytical data for the alka li earth metal and other

divalent cation concentrations in solution; and (e) the mineralogical evidence for the presence of smectite in both vadose and phreatic zone sediments. We are led to the tentative conclusion that a shifting state of ion exchange equilibrium with re spect to interlayer site occupancies on smectite may be responsible for the observed res ponse of the groundwater composition to CO 2 

### uptake, although not necessarily driven exclusively by Ca +2 

### released by the dissolution of calcite. This process can be represented by the following equati on, where X represents a negatively charged (-2) ion exchange site:

X-M(II) + 2H 2 

### CO 3 

### (aq) = X-H 2 

### + (M +2 

### •2HCO 3 - 

### ) (5-4)

Alkalinity (reported as HCO 3 - 

### ) is represented approximately by the concentration of HCO 3 - 

### in this particular system (where minor complexation by HCO 3 - 

### and the contribution due to CO 3 -2 

### can be ignored). Therefore,

log (HCO 3 - 

### )  1⁄2-log(M +2 

### ) (5-5)

The trendlines in Figures 5.2a–d gi ve slopes that approximate unity rather than 0.5, as would be expected from the above relationship. The obser ved correlation between alkalinity and the alkali earth metals is therefore not readily interpreted. Synthetic geochemical modeling of the response of a groundwater composition in e quilibrium with a cation exch anger to changes in P(CO 2 

### ) should be conducted to further elucidate the many complex system interactions that might explain the observed relationships.

Figures 5.2e and f show the negative correlations between phosphate (log PO 4 

### ), and log Fe and log Co, with trendline slopes of ~1.5 and 0.75, respec tively. The correlation between log Fe and log

PO 4 

### is strongly suggestive of an equilibrium state between these spec ies and some ferrous phosphate mineral such as vivianite:

Fe 3 

### (PO 4 

### ) 2 

### .8H 2 

### O = 3Fe +2 

### + 2PO 4 -3 

### ; log [PO 4 -3 

### ] = K ́ + 3/2 log[Fe +2 

### ] (5-6)

Calculations show, however, that prior to CO 2 

### injection, the MSU-ZERT groundwaters were unsaturated with respect to bot h vivianite and strengite (FePO 4 

### .2H 2 

### O) prior to CO 2 

### injection. The slope coefficient of the trendline for cobalt could be explained by achievement of solution equilibrium with respect to the cobalt analogue of, for example, gormanite (Fe II 3 

### Al 4 

### (PO 4 

### ) 4 

### (OH) 6 

### •2(H 2 

### O)), mclausanite (HFe II 3 

### Al 2 

### (PO 4 

### ) 4 

### F•18(H 2 

### O)) or keckite Ca(Mn,Zn) 2 

### Fe III 3 

### (PO 4 

### ) 4 

### (OH) 3 

### •2(H 2 

### O), where:

log [PO 4 -3 

### ] = K ́ + 3/4 log[Co +2 

### ] + ..., or log [PO 4 -3 

### ] = K ́ + 3/4 log[Co +3 

### ] + ... (5-7)

These examples are illustrative of the potential complexity associated with the mineralogy of shallow sedimentary systems. In reality, Co(II) is likely to substitute for Fe(II) in solid solution in Fe(II) hydroxides, such as green rust, but the pr esence of trace nanophases exercising solubility

control over trace elements ca nnot be discounted, and insolu ble phosphates might under some circumstance be potential candidates.

### 36 We now return to the correlation of the chemical species with a restri cted set of principal components. Information on the first ten principal components is presented in Table 5.2. The first four principal components explain over 80% of the variance, whereas principal components No. 5 (PC5) and No. 6 (PC6) explain another 10%. As w ith Table 5.1, the correlation coefficients are similarly color coded in range s 0.9, 0.8, 0.7, 0.6, 0.5 and also 0.4.

Table 5.2. Principal components derived from an analysis of groundwater chemical compositions for the period July 7 to July 17, 2008

Explained Variance (Eigenvalues)

Value PC 1 PC 2 PC 3 PC 4 PC 5 PC 6 PC 7 PC 8 PC 9 PC 10 Eigenvalue 17.353 3.073 2.668 1. 480 1.345 1.071 0.892 0.644 0.466 0.289

% of Var. 57.844 10.242 8.894 4. 934 4.484 3.569 2.975 2.145 1.553 0.962

Cum. % 57.844 68.086 76.980 81.914 86.398 89.967 92.942 95.087 96.641 97.603

Component Loadings (correlations between in itial variables and principal components)

Variable PC 1 PC 2 PC 3 PC 4 PC 5 PC 6 PC 7 PC 8 PC 9 PC 10

H 0.774 -0.419 0.114 0.090 -0.289 -0 .079 -0.111 0.051 0.279 0.018

Li 0.951 0.202 0.120 -0.010 -0.093 -0 .063 0.071 -0.056 -0.026 0.086

Na 0.874 0.023 0.216 -0.246 0.223 -0 .015 -0.064 0.062 0.113 -0.169

K 0.711 -0.500 0.071 0.245 -0.252 0. 031 0.076 0.136 -0.005 -0.226 Mg 0.946 -0.215 0.111 -0.044 0.117 0. 115 -0.059 0.029 -0.002 -0.002

Ca 0.914 -0.339 0.175 -0.014 0.018 0.062 0.023 0.051 -0.006 -0.070

Sr 0.931 -0.282 0.154 -0.060 0.084 0.077 -0.028 0.061 0.004 -0.044

Ba 0.902 -0.385 0.030 -0.053 0.025 0. 127 -0.035 0.028 -0.032 -0.034

Mn 0.484 0.229 -0.735 -0.236 -0.001 -0. 029 0.199 0.025 -0.033 -0.127

Fe 0.841 -0.169 -0.316 -0.257 0.156 0.019 0.088 -0.118 -0.164 0.044

Co 0.734 0.100 -0.537 -0.351 0.032 -0.015 0.077 -0.022 -0.070 0.045

Cu 0.857 0.201 0.275 0.113 0.176 -0 .060 0.211 0.018 0.000 -0.116

Zn 0.282 0.372 0.221 -0.287 -0.268 0.663 0.076 0.317 -0.098 0.051

Cd 0.914 0.156 -0.013 0.080 -0.234 -0.173 0.163 0.030 0.044 0.119

Pb 0.872 0.287 0.007 0.155 -0.184 -0 .233 0.109 0.089 -0.060 0.060

Al 0.972 -0.021 0.131 0.055 -0.054 -0 .019 0.066 -0.004 -0.043 0.065

Cr 0.872 0.276 0.070 0.116 -0.298 -0 .164 0.079 -0.061 -0.007 0.068

As 0.352 0.482 0.529 -0.105 0.044 0.174 -0.330 -0.410 -0.073 -0.109

Se 0.832 0.417 0.189 0.013 -0.142 0.001 -0.115 -0.228 -0.020 0.005

Mo 0.620 0.441 -0.058 0.196 0.313 -0.411 -0.150 0.166 -0.145 -0.143

U 0.497 -0.203 0.325 -0.258 0.562 -0.203 -0.126 0.194 0.086 0.208

F -0.201 0.870 0.095 -0.056 0.047 0.006 -0.108 0.368 0.004 -0.045

Cl 0.385 0.035 -0.573 0.615 0.168 0.257 -0.102 0.029 -0.021 -0.081

Br 0.252 0.387 0.148 0.340 0.464 0.286 0.497 -0.140 0.275 0.046

NO3 -0.477 -0.303 0.547 0.195 0.083 -0.056 0.352 0.005 -0.431 0.026

PO4 -0.791 0.034 0.434 0.310 -0.094 -0.030 -0.125 0.165 0.027 0.012

SO4 0.631 -0.104 -0.347 0.406 0.217 0.245 -0.359 0.014 -0.153 0.179

HCO3 0.928 -0.271 0.141 -0.061 0.099 0. 097 -0.032 0.065 -0.014 -0.007

SiO2 0.967 0.079 0.165 0.002 -0.079 0.028 -0.056 -0.076 -0.043 -0.009

B 0.910 0.172 -0.084 0.202 -0.223 -0.045 -0.061 0.031 0.038 0.107

LEGEND

> 0.9

> 0.8

> 0.7

> 0.6

> 0.5

> 0.4



### With respect to principal component No. 1 (PC1), nine of the chemical species show correlations of 0.9 and above, and another six show correlations between 0.8 and 0.9. With the exception of HCO 3 

### ,

### 37 Li and Na, all of the remainder ar e probably present as divalent cat ions or divalent hydroxy-cations of trivalent species, i.e., B, Al, Cr, and Mn. Principal component No. 2 (PC2) correlates strongly only with F, and relatively weakly with As, Se and Mo, all potential oxy-anionic species, and inversely, but weakly with H and K. Principal co mponent No. 3 (PC3) invers ely correlates with Mn, Co and Cl, and directly with As, NO 3 

### and PO 4 

### . Principal component No. 4 (PC4) correlates with Cl

and SO 4 

### . Correlations of chemical species with the re maining six components are for the most part weak and are of questionable significance.

Interpretation

Each of the primary principal components most pr obably relates to a chemical process that takes place rapidly, given the short interval over which th e analyses were taken, and considering the fact that a close logarithmic relationshi p is observed with respect to HCO 3 - 

### in the first principal component. The most likely short term processes ar e, in order of reaction rate: adsorption/desorption, ion exchange (both cationic and anionic), dissolution of carbonates and metastable Fe hydroxides and green rusts, some homogeneous re dox reactions in the aqueous phase, and pore

water infiltration from the vadose zone. Other unid entified processes may also be involved, which future investigations might reveal.

Table 5.3 lists the correlated variable for each of the first six principal co mponents, and identifies the likely processes in which the dominant aqueous species might participate. Note that over the possible range of redox states in the shallow phreatic zone at the ZERT site, several aqueous species might transform to other species with different ox idation states, notably Fe(III)/Fe(II), Cu(II)/Cu(I), Cr(VI)/Cr(III), N(VII)/N(0), U(VI)/U(IV), Se(VI)/ Se(IV)/Se(-II). It is also possible, although less likely, that extremely reducing conditions might arise through the progressive decomposition of organic matter, leading to transformations w ith respect to As(III)/A s(-I), S(VI)/S(-II), Mo(VI/Mo(II), C(IV)/C(0) /C(-IV), and N(0)/N(-III ). It should also be noted that many trace elements could be present in solid solution in minerals that are known to dissolve rapidly in response to changing pH, such as calcite and/or hydroxide minerals including green rust. These have not been checked in Table 5.3 under “Precipitation/dissolution,” because the extent of such

solid solutions is unknown.

Table 5.3 suggests that PC1 may be related to io n exchange or desorption processes because the

alkali earth metals and other divalent cations co rrelate very strongly with this component. Several monovalent cations (H + 

### , Li + 

### , Na + 

### , K + 

### ) are also moderately to strongl y correlated. These correlations suggest cation exchange on smectite in terlayer sites. Trivalent species such as B, Al, Fe and Cr, and tetravalent Si (as SiO 2 

### (aq)) are also moderately to strongly correlated, and are known to sorb onto clays. However, the strong positive correlation with HCO 3 - 

### is not readily expl ained in relation to sorption or ion exchange on smectite. Other oxy-anionic species, such as SO 4 -2 

### , Mo (as MoO 4 -2 

### ), Se (as SeO 4 -2 

### or SeO 3 -2 

### ) are also positively correlated. It is te mpting to assume that these species likely occupy the interlayer sites of an ionic clays such as fougerite, a Mg-rich form of green rust, especially when it is recognized that such occupancy in green ru st has been confirmed in the laboratory for all but MO 4 -2 

### . However, the oxy-anions, PO 4 

### an NO 3 

### , are negatively correlated, implying a different association with PC-1, whic h is not readily explained without further study.

PC2 appears to be dominated by anioni c species (F, As (probably as HAsO 4 -2 

### ), Se (as SeO 4 -2 

### or SeO 3 -2 

### ) and Mo (as MoO 4 -2 

### ), which are positively correlated, and by H and K, which are negatively correlated, and suggest adsorption on hydrous ferric oxide (HFO), which could be ferrihydrite,

### 38 lepidocrocite, or goethite. PC3 shows a negative correlation with respect to Mn, Co and Cl, and a positive correlation with respect to As (probably as HAsO 4 -2 

### ), NO 3 

### and PO 4 

### . This correlation is suggestive of saturation control by some transition metal phosphate with substitution of AsO 4 -3 

### for

PO 4 -3 

### . PC4 may be related to infiltration of pore water from the overlying vadose zone. U is correlated with PC5 and Zn appears to show the strongest correlation by far with PC-6. No explanation for these correlations is apparent. The behavior of Zn appears to be unique, and is unexpected.

Because smectite has been identified as a mineral th at is present in the sediments in the saturated zone, it is reasonable to assume as a working hypot hesis that ion exchange on smectite is operative, and is responsible for the species correlations obse rved with respect to PC1. It is not entirely clear, though, what specific ion exchange mech anism is taking place. Increasing H + 

### caused by the

dissolution of CO 2 

### might be responsible. Howe ver, the correla tion between H + 

### and the first principal component is only moderate, and as no ted in preceding discussions, does not explain the strong positive correlation with HCO 3 - 

### . Another possibility is that increased P(CO 2 

### ) leads to calcium carbonate dissolution, which in turn induces ion exchange with other divalent cations

observed to increase in concentration in the groundwater. This mechan ism would therefore be controlled by the rate of calcite dissolution. As a working hypothesis, we will assume that the primary response to increasing c oncentrations of dissolved CO 2 

### is cation exchange and desorption induced by Ca +2 

### due to calcite dissolution, and anion exchange induced by increasing HCO 3 - 

### in

solution. We will also assume that an adsorbent is present whose surface behaves analogously to HFO such as ferrihydrite, or other ferric oxy -hydroxides and other metastable phases (e.g., schwertmannite and green rust) commonly resulti ng from the transformation of HFO (see further discussion in Section 5.2.2). The presence of these minerals can also provide anionic exchange sites as suggested by the correlati on between PC1 and PC2 with respec t to aqueous anionic species. The modeling of this system is presented s ubsequently in Section 5.3 of this report.

### 39 Table 5.3. Correlation of aqueous species with principal components for the period July 7 to July 17, 2008 in relation to short-term chemical processes

Component Variable Correlation Dominant Aqueous Species Cation Exchange Anion Exchange Adsorption Precipitation/

dissolution Redox

Reactions Infiltration

1 H 0.774 H + √ √

Li 0.951 Li + √ √

Na 0.874 Na + √ √

K 0.711 K + √ √

Mg 0.946 Mg +2 √ √ √

Ca 0.914 Ca +2 √ √ √

Sr 0.931 Sr +2 √ √

Ba 0.902 Ba +2 √ √

Mn 0.484 Mn +2 √ √ √

Fe 0.841 Fe +2 √ √ √ √

Co 0.734 Co +2 √ √ √

Cu 0.857 Cu +2 /Cu + √ √ √

Cd 0.914 Cd +2 /CdSe √ √

Pb 0.872 PbCO 3 /PbSe √ √

Al 0.972 Al(OH) 3 - H 3 SiO 4 - ? ? √ √

Cr 0.872 CrO 4 -2 /CrO + ? √ √ √

Se 0.832 SeO 4 -2 √ √ √

Mo 0.620 MoO 4 -2 √

U 0.497 UO 2 (CO 3 ) 2 -

2 /UO 2 √ √

NO3 -0.477 NO 3 - √ √ √ √

PO4 -0.791 H 2 PO 4 - √ √

SO4 0.631 SO 4 - 2 √ √ √ √

HCO3 0.928 HCO 3 - √ √ √ ? SiO2 0.967 Si(OH) 4 √ √

B 0.910 B(OH) 3 √

2 H -0.419 H+ √ √

K -0.500 K + √ √

As 0.482 AsO 3 F --2 √ √ ?

Se 0.417 SeO 4 -2 /HseO 3 - √ √ √

Mo 0.441 MoO 4 -2 √ ?

U -0.203 UO 2 (CO 3 ) 2 -

2 /UO 2 √ √

F 0.870 F

- √

3 Mn -0.735 Mn +2 √ √ √

Co -0.537 CO 2 (OH) 3 + √ √

As 0.529 AsO 3 F -2 √ ?

Cl -0.573 Cl

- √ √

NO3 0.547 NO 3 - √ √

PO4 0.434 H 2 PO 4 - √ √

4 Cl 0.615 Cl- √

SO4 0.406 SO 4 -2 √

5 U 0.562 UO 2 (CO 3 ) 2 -

2 /UO 2 √ √

Br 0.464 Br

- √

6 Zn 0.663 Zn +2 √ √

Mo -0.411 MoO 4 -2 √ √ ?

### 40 5.2.2 Thermodynamic Evaluation of Base line Groundwater Composition and Redox
State

One of the more difficult problems associated with interpreting the geochemistry of the groundwaters in the saturated zone at the ZERT Site is the determ ination of the redox state. In addition, as discussed further below, it has been clearly established that complete redox equilibrium in shallow groundwater systems is rarely, if ever, achieved (e.g., Lindberg and Runnels, 1984; Washington et al., 2004; Stefansson et al., 2005). Furthermore, many of the trace elements and some of the major constituents can undergo reduction or oxidation in response to periodic infiltration of rainfall, which transports not only dissolved constituents already present in the rain (Na + 

### , Mg +2 

### , Ca +2 

### , NH 4 + 

### , HCO 3 - 

### , NO 3 - 

### , SO 4 -2 

### , Cl -, 

### Br - 

### ), which have been subject to evapotranspiration, but also dissolve d inorganic constituents and orga nic matter from the overlying soil zone. The transported organic matter is present bo th in the form of dissolved organic acids and associated complexes with inorga nic cations, but also is probably present in colloidal form. Both during transport and after reaching the water table, the organic ma tter in the pore waters undergoes oxidative mineralization through the agency of bacteria and fungal matter, depleting all dissolved and gaseous phase oxygen in the vadose and phreatic zones, and then reducing both dissolved and solid-phase electron donors. These include N(V) as NO 3 - 

### , N(III) as NO 2 - 

### , N(0) as N 2 

### , S(VI) as

SO 4 -2 

### , and dissolved minor elements and trace me tals involving species and complexes and associated solid phases in the following oxidation states; C(IV)/C(0) /C(-IV), Fe(III)/Fe(II), Mn(IV)/Mn(III)/Mn(II), As(V )/As(III)/As(-I)/As(-III), Se(VI)/Se( IV)/Se(0)/Se(-II), Cr(VI)/Cr(III), Co(III)/Co(II), Cu(II)/Cu(I), Mo(V I)/Mo(II) and U(VI)/U(IV).

The traditional approach taken in trying to char acterize the redox state of groundwaters is through the use of a platinum (Pt) elec trode and associated reference el ectrode, e.g., calomel or Ag/AgCl, which measures the redox potential of the gr oundwater when in thermodynamic equilibrium. However, as noted above, shallow groundwaters are rarely if ever at thermodynamic equilibrium, and therefore the measured mixed potential cannot be utilized as a meaningf ul representation of the redox potential. Furthermore, a Pt electrode can locally catalyze redox reactions on its surface, leading to precipitation of PtO, which, in combina tion with the Pt substrate, results in it responding to pH rather than Eh (see Baas Becking, 1960). The Pt electrode can also be poisoned. Thus, unless great precaution is taken, the measured potential may have more to do with the electrode sensor

than the solution in which it is immersed.

The problems associated with direct in situ measurement of redox poten tial using a Pt electrode have been known for many years (Baas Becki ng, 1960). Following Lindberg and Runnels (1984), therefore, attempts have been made to chemi cally analyze dissolved redox pairs in groundwaters (e.g., Washington et al., 2004; Stef ansson et al., 2005), and then calculate the apparent redox potential for each pair. As expected, the analyz ed redox pairs yield disparate redox “potentials.” indicating a nonattainment of homogeneous equ ilibrium. Washington et al. (2004) made a significant contribution, when they showed that the potentials asso ciated with the analyzed redox pairs are not random, but behave in a systematic manner, and that “...redox couples composed of concentrated reactants achieve nearly thermodynamic equilibrium concomitantly, and those with dilute reactants deviate from equilibrium by an amount roughly inverse to thei r concentrations.” In typical groundwaters, an upper group of redox couples is essentially controlled by the presence of

NO 3 

### , which is involved in the oxidation of speci es such as, for example, Fe(II), thus:

NO 3 - 

### + 2Fe +2 

### + 5H 2 

### O  NO 2 - 

### + 2Fe(OH) 3 

### (ppt) + 4H + 

### (5-8)

### 41 and a lower group controlled by the presence of disso lved organic matter, which can be represented by equations of the type:

3CH 2 

### O + 2H + 

### + SO 4 -2 

###  H 2 

### S(aq) + 2CO 2 

### (aq) + 2H 2 

### O , (5-9)

where CH 2 

### O is a generic representation of dissolved (a nd particulate) organic matter. According to Washington et al. (2004), the upper cluste r of redox pairs is centered on pe ≈ +10, whereas the lower cluster is centered on pe ≈ -2. Over time, or with increasi ng depth below the water table, all

NO 3 

### in solution is progressively c onsumed, and the redox pairs, initia lly in local partial equilibrium with respect to NO 3 

### , drift to lower redox states, where cont rol of the redox stat e is assumed by the reservoir of organic matter and microbial activity, which is usually the dominant reducing agent in shallow unconfined aquifers underlying soil zones w ith active vegetative cover. This scenario is particularly pertinent to the redox state of the ZERT site, as is described further below.

While the direct characterization of redox pa irs represents the mo st rigorous means of characterizing the redox state in a groundwater, such an approach is rarely implemented in practice, because it requires elaborate sampling procedur es to avoid inadvertent contamination by atmospheric oxygen, and time-consuming and expens ive analytical procedures. A more typical scenario is represented by the current characte rization of the ZERT site and elsewhere—e.g., see Bourrie et al. (1999) and Fede r et al. (2005)—where collecte d groundwaters are subjected to standard field and laboratory analytical procedur es, and where only the total concentrations of

elements in solution are determined rather than th eir speciation, including th e distribution of species

between oxidation states.

An alternative means of indirect ly characterizing the oxidation state of a gr oundwater is through a thermodynamic analysis of the groundwater compositi on using a distribution of species code such as PHREEQC (Parkhurst and Appelo, 1999) or EQ3/6 Wolery, 1992, 1993) and calculating the

saturation states of minerals th at might control the redox state of the system by varying the redox potential over a specified range and examining at what pe (or Eh) potential redox controlling minerals equilibrate with the redox affected aque ous species in solution. Ideally, the redox state can be established where one or more redox sensitive phases containing the same redox pair coexist in a state of saturation at a particular value of pe. Such an analysis is in some respects more fundamental, in that solid phases participating in redox reactions are likely to exert a substantial buffering capacity on the system as a whole, and are more likely to define a temporally averaged redox state, rather than a mo mentary, but not necessarily representative oxidation state characterized by the analysis of an isolated groun dwater sample. Such utili zation of redox sensitive minerals to help define the oxidation state of an aqueous system is subject to several limitations and

caveats:

1. The redox sensitive phases must be reactive, i.e. , they must be able to respond by dissolving or precipitating sufficiently rapidly, so that they reas onably approach thermodynamic equilibrium under ambient conditions. Such phase s are most likely to be metastable and either finely crystalline or amorphous.

2. The redox sensitive phases must be clearly id entified, and their thermodynamic properties reasonably well characterized.

3. Ideally, several redox sensitive components should be considered in the analysis in order to provide independent corroboration.

### 42 4. The participating aqueous species must be pres ent at sufficiently high concentrations to be detectable and quan tifiable in solution.

5. Groundwater sampling procedures exclude cont amination, and custodial integrity can be assumed (i.e., samples have not re-equilibrated under post-sampling redox conditions).

6. The thermodynamic properties of the aqueous species must be known and their complexation with other species in so lution must be known and quantified.

7. Suitable corrections to the thermodynamic data used in the analysis should be made to account for ambient field temperatures of 10oC rather than the standard reference temperature of 25oC used in most therm odynamic calculations concerning groundwaters.

8. Validation of the thermodynamic analysis should be made through an independent mineralogical study of the abundance and dist ribution of redox sens itive phases in the soil/sediment profile through the vadose and phreatic zones where the groundwaters were sampled.

In the following analysis, a modified EQ3/6 ther modynamic database was used as described in Birkholzer et al. (2008). The databa se is a substantial revision of an earlier database, Data0. ymp.R4, which was developed and qualified for use by the Yucca Mountain Project on behalf of the U.S. Department of Energy. A very importa nt requirement for conducting a successful thermodynamic evaluation of the ZERT site is the incorporation of the solubility products of metastable redox sensitive minerals that might exerci se partial or complete secular contro l over the redox state of the system. Among these is a class of minerals referred to as “green rusts” of which fougerite is the most important na turally occurring representative. Because this mineral is unstable under oxic conditions, its recovery and preservation during coring for soil and sediment samples is questionable unless appropriate pr ecautions are taken to avoid c ontamination with atmospheric oxygen. Despite the fact that such precautions were not anticipated or implemented, circumstantial evidence suggests that it may be present at the ZERT site, and some geochemical similarities between the ZERT site and the type locality for fouge rite in Brittany mandate that it be taken into consideration in the following thermodynamic eval uation. (By “type locality”, we refer to the original site from which the f ougerite was initially obtained, ch aracterized mineralogically, and reported in the scientific liter ature.) Appendix C presents a su mmary description regarding the known thermodynamic properties of fougerite and its compositional variation, as cited in the recent literature, and describes the procedure adopted for calculating its solubi lity product in a form appropriate for inclusion in Data0.ymp.R4.

Baseline Groundwater Composition

The groundwater composition selected for evaluation of the redox state is based on the average of seven comprehensive chemical analyses of gr oundwaters taken on July 7 and 8, 2008, immediately prior to CO 2 

### injection. With the exception of sample 08ZERT-104, all samples were taken from the deeper A-wells and from two mon itoring wells drilled further away from the injection well (MW1 and MW2). It is possible that a redox gradient exists between th e water table and deeper horizons within the phreatic zone, and that sample 08ZERT-104 may reflect more oxidizing conditions. However, the chemical analysis does not suppor t this interpretation, and even suggests a lower redox potential induced by exogenic contamination, as is noted further below. These analyses and the average values are given in Table 5.4.

### 43 Table 5.4 Initial composition of the groundwater at the ZERT site prior to CO 2 injection

SAMPLE T pH Li Na K Mg Ca Sr Ba Mn Fe

°C mg/L mg/L mg/L mg/L mg/L mg/L mg/L mg/L mg/L

08ZERT

- 102 10 6.8 0.003 7.5 6.2 24 85 0.27 0.13 0.002 <0.01

08ZERT

- 103 10 6.8 0.003 9.2 5.5 24 95 0.31 0.12 0.001 <0.01

08ZERT

- 104 15 6.9 0.006 9.7 5.6 27 90 0.30 0.10 0.11 0.05

08ZERT

- 105 8 6.9 0.003 7.8 5.6 23 84 0.27 0.11 0.001 <0.01

08ZERT

- 106 8 6.9 0.004 8.1 5.6 25 86 0.28 0.10 0.31 <0.01

08ZERT

- 107 8 6.7 0.005 7.5 5.3 24 84 0.26 0.11 0.001 <0.01

08ZERT

- 109 12 7.0 0.007 9.1 5.4 28 92 0.30 0.10 0.028 <0.01

Average 10.286 6.871 0.004 8.414 5.583 24.992 87.988 0.285 0.110 0.065 0.015

SAMPLE Co Cu Zn Cd Pb Al Cr As Se Mo U

ug/L ug/L ug/L ug/L ug/L ug/L ug/L ug/L ug/L ug/L ug/L

08ZERT

- 102 <0.2 1.4 3.4 0.07 <0.02 2 <0.4 0.8 <2 0.4 4.3

08ZERT

- 103 0.3 1.5 4.0 0.09 <0.02 2 0.4 1.0 <2 0.3 4.9

08ZERT

- 104 0.8 2.4 2.4 0.26 0.08 3 10 1.0 <4 1.3 4.5

08ZERT

- 105 <0.2 1.2 2.8 0.08 0.02 2 <0.4 0.9 <2 0.4 4.0

08ZERT

- 106 0.6 1.1 3.2 0.10 0.02 2 <0.4 0.8 <2 0.3 4.0

08ZERT

- 107 0.2 1.6 3.1 0.23 0.05 3 12 1.0 <4 0.4 4.1

08ZERT

- 109 0.4 2.5 3.8 0.29 0.06 3 12 1.3 5 0.7 4.3

Average 0.376 1.668 3.245 0.162 0.038 2.383 5.010 0.976 3.003 0.545 4.291

SAMPLE F Cl Br NO3 PO4 SO4 HCO3 SiO2 B 11

mg/L mg/L mg/L mg/L mg/L mg/L mg/L mg/L mg/L

08ZERT

- 102 0.09 6.5 0.051 0.12 0.20 8.3 396 27 0.012

08ZERT

- 103 0.10 3.7 0.042 0.79 0.17 5.8 432 27 0.009

08ZERT

- 104 0.17 5.3 0.046 0.10 0.09 7.3 431 33 0.017

08ZERT

- 105 0.10 4.7 0.040 0.55 0.17 7.5 386 26 0.012

08ZERT

- 106 0.11 5.1 0.044 0.19 0.13 7.4 398 26 0.013

08ZERT

- 107 0.12 4.8 0.042 0.60 0.18 7.4 387 32 0.018

08ZERT

- 109 0.17 5.4 0.099 0.26 0.10 7.2 434 32 0.017

Average 0.123 5.071 0.052 0.373 0.149 7.270 409.129 28.9

31 0.014



### In general, most of the major chemical constituen ts and trace elements show relatively consistent values, giving one confidence that the analyses truly reflect the actual concentrations present in the groundwater. However, several specie s of critical interest in defi ning the oxidation state of the groundwater and the saturation state of redox sensitiv e phases are at or close to their detection limits. These include Mn, Fe, Cr and Se. Furt hermore, sample 08ZERT-104 shows anomalously high concentrations of Fe, Cu, Cd, Cr and Mo, su ggesting contamination by steel. The averaged

concentrations should therefore be considered as maximum values in interpreting the subsequent results. However, it should also be noted that th ese same species show variations in concentrations subsequent to CO 2 

### injection that place them well within the range of detectability for the analytical

methods used.

### 44 Thermodynamic Analysis of Baseline Groundwater

The average concentrations of the chemical species given in Table 5.4 were entered into the input file of EQ3/6, and an initial run was made to calculate the extent of complexation by HCO 3 - 

### at 25oC. The analytical value for HCO 3 - 

### was adjusted to account for cation complexation, and the revised

value inserted as a “free” concentration. In this manner, the total equilibri um concentration of all carbonate species, including CO 2 

### (aq), can be accounted for in solution, and P(CO 2 

### ) determined. EQ3/6 was then re-run, setting the temperature to the average ambient value of 10.28oC. The resulting charge imbalance was cal culated to be –0.50% of the mean charge of all charged species, which is excellent, and therefore no attempt was made to ensure electrical neutrality by balancing on any selected species.

The calculated saturation indices (SI) for several primary and secondary minerals of interest, but

which are not sensitive to the redox state of the system, are li sted in Table 5.5. The saturation

indices of aluminosilicate minerals have been nor malized to both Al and Si to afford a better respective comparison with the gibbsite (Al(OH) 3 

### ) and silica (opal-CT or quartz) chemical components. These results are typica l of what would be expected in a soil zone of this type, where primary rock-forming minerals asso ciated with detrital igneous, metamorphic, and sedimentary rock fragments observed in the shallow subsurface envi ronment at the ZERT site are weathering in the presence of neutral to weakly acid groundwater s containing dissolved organic matter from the overlying soil zone. The aluminum concentration is elevated and actually supersaturated with respect to gibbsite, which is consistent with field observations elsewhere (Dobrzynski, 2007), and is caused by the formation of metastable amorphous or poorly crystalline hydroxy-aluminum silicates (HAS) such as imogolite or allophane as weatheri ng products of detrital minerals of igneous or metamorphic provenance. These HAS minerals are al so closely associated with secondary opal-CT, which forms concurrently in systems with excess s ilica. Recent work has s hown that HAS minerals form as metastable precursors to secondary minera ls actually observed in the fine fractions of sediments recovered from the site, i.e., smectite, illite, and kaolinite. The transformation from HAS to these secondary clays is exceedingly slow, and therefore the former, together with opal-CT, controls the activity of Al(OH) 3 

### and SiO 2 

### components in the aqueous phase. As a result, the weathering minerals can approach thermodynamic equilibrium, and indeed, secondary epitaxial overgrowths of K-feldspar on detrital feld spar grains have sometimes been observed.

The groundwater at the ZERT site does not appear to be saturated with respect to either HAS or to opal-CT. Calcite also appears to be undersaturate d, with SI(calcite) = -0. 283. Given the fact that secondary calcite precipitates are ubiquitous both below and above th e water table at the ZERT site (Section 4), it is perhaps surprising that it is not closer to saturation. Furthermore, mineralogical studies of secondary calcite sugges t that it may be magnesium rich. Such calcites are less stable than the pure form (Busenberg and Plummer, 1989), which would imply even greater degrees of undersaturation. However, undersaturation with respect to HAS, opal-CT and calcite can be explained by transient dilution of the groundwater due to rainfall prior to the start of the test. Table 5.5 shows that all of the listed potential second ary phyllosilicates except celadonite, including those actually observed to be present in sediments at the ZERT site, are supersaturated, as expected. The supersaturation of K-feldspar is consistent with some observati ons elsewhere (Zhu et al., 2006), but may be related to potassium retention by vegetati on in the presence of illite, rather than any thermodynamic or kinetic control by this mineral, especially when it is recognized that the pH is ≤ 7. The observed undersaturation with resp ect to low albite is expected.

### 45 Table 5.5 also lists the sa turation indices of some other minerals of passing interest. Dolomite is strongly undersaturated, as is char acteristic of shallow groundwaters (Edmunds et al., 1982; Back et al., 1983; Hopkins and Putnam, 2000; James et al ., 1993; Pacheco and Szocs, 2006), and therefore dolomite, if present, would tend to dissolve. As men tioned in Section 4, individu al grains of detrital dolomite and occasional fragments of dolomitic limest one were identified in the coarse fractions of

sediments from the ZERT site. This material shows evidence of corrosion, wh ich is consistent with the findings of the thermodynamic analysis. Earl ier work (Birkholzer et al., 2008) involving the thermodynamic evaluation of 38,000 potable groundwater s showed that the concentration of Ba was controlled primarily with respect to barite (BaSO 4 

### ). Table 5.5 shows that barite is significantly undersaturated with respect to both barite and to alstonite, the barium an alogue of dolomite. This finding is also suggestive of dilution by rainfall , preventing the accumulation of the products of weathering to saturate, except possibly after prolonge d dry spells. Finally, the saturation indices of hydroxy-apatite, fluor-apatite and crandallite ar e compared. Hydroxy-apatite is strongly undersaturated, whereas fluor-apatite is strongly supersaturated. Many of th e detrital lithic fragments at the ZERT site consist of amphibolites or extrusive mafic rocks including andesitic volcanics. These rocks typically contain apatite, commonly with significant F - 

### substitution for OH - 

### . Therefore, it is possible that the shallow groundwater s may be saturated with respect to hydroxy-fluor apatite. It is more likely, however, that ap atite is unstable with re spect to secondary crandallite, which is close to

saturation. Other alkali-earth metals such as Sr a nd Ba, and also Pb and Cd are known to substitute for Ca, and AsO 4 -3 

### , VO 4 -3 

### , CO 3 -2 

### , MoO 4 -2 

### , SO 4 -2 

### , SeO 4 -2 

### and even CrO 4 -2 

### can substitute for PO 4 -3 

### in the crandallite crystal structure. Fe +3 

### and Cr +3 

### can also substitute for Al +3 

### in crandallite—see Schwartz, et al. (2000). Therefore, this minera l or an isostructural mixed phosphate-sulfate, woodhouseite (CaAl 3 

### (PO 4 

### )(SO 4 

### )(OH) 6 

### ), could exercise partial thermodynamic control over the

activities of a number of trace elements in the groundwater.

### 46 Table 5.5. Saturation indices of redox-insensitive primary and secondary minerals with respect to groundwater at

the ZERT site

Name Formula SI SI(n. Al) SI(n. Si)

Quartz SiO2 +0.633 - +0.633

Opal-CT SiO 2 -0.267

- -0.267

Gibbsite Al(OH) 3 +0.402 +0.402

-
Hydroxy Aluminum and Magnesium Silicates

Imogolite Al 2 SiO 5 .xH 2 O -0.897 -0.448 -0.897

Hydroxy

aluminium silicate Al 2 Si 2 O 5 (OH) 4 -2.310 -1.155 -1.155

Kaolinite Al 2 Si 2 O 5 (OH) 4 +2.440 +1.2 20 +1.220

Talc Mg 3 Si 4 O 10 (OH) 2 -5.540 - -1.385

Clinochlore Mg 5 Al 2 Si 3 O 10 (OH) 8 +2.332 +1.1 66 +0.777

Pyrophyllite Al 2 Si 4 O 10 (OH) 2 +1.919 +0.9 59 +0.480

Smectites

Beidellite-Ca Ca 0.165 Al 2.33 Si 3.67 O 10 (OH) 2 +2.206 +0.9 47 +0.601

Smectite-Ca Ca 0.145 Mg 0.26 Al 1.77 Si 3.97 O 10 (OH) 2 +4.639 +2.6 21 +1.169

Montmorillonite-Ca Ca 0.165 Mg 0.33 Al 1.67 Si 4 O 10 (OH) 2 +2.175 +1.3 02 +0.544

Phengites

Celadonite KMgAlSi 4 O 10 (OH) 2 -0.355 -0.355 -0.089

Illite K 0.6 Mg 0.25 Al 1.8 Al 0.5 Si 3.5 O 10 (OH) 2 +1.647 +0.7 16 +0.544

Feldspars

K-feldspar KAlSi 3 O 8 +1.009 +1.0 09 +0.336

Albite-low NaAlSi 3 O 8 -1.887 -1.887 -0.629

Carbonates

Calcite CaCO 3 -0.283

- -
Dolomite CaMg(CO 3 ) 2 -3.196*

- -
Alstonite BaCa(CO 3 ) 2 -2.435

- -
Phosphates

Hydroxy-apatite Ca 5 (OH)(PO 4 ) 3 -4.992

- -
Fluor-apatite Ca 5 (F)(PO 4 ) 3 +5.721

- -

Crandallite CaAl 3 (PO 4 )(OH) 5 .H 2 O -0.230 -0.077 -

Sulfates

Barite BaSO 4 -0.440

- -
* Solubility product is based on data reported by Rock et al. (2001)



### To evaluate redox sensitivities, EQ3/6 was re -run specifying the redox state at successive increments ranging from –2 to +14 in units of pe (-log[e-]) (mol/kg H 2 

### O) at 2-unit pe intervals. The

### 47 principal aqueous species of Fe, Mn and selected trace elements ar e displayed for each pe increment in Table 5.6. Note that Mn, Co, Cd, Pb and Zn re main in the (II) state throughout the range of pe’s investigated. However, solid Mn and Fe phase s coexisting with aqueous Mn and Fe do show changes in oxidation state. The dominant specia tion for Pb and Zn also changes under strongly reducing conditions. Changes in aq ueous speciation associated with changes in oxidation state are observed for Fe, Cu, Cr, Se, U, and NO 3 

### , but changes are not observed with respect to As, which remains in the As(V) state, and S (as SO 4 

### ), which remains in the S(VI) state, although reducing conditions are likely to stabilize As(O), As (-I), As(-II), S(-I), and S(-II), respectively.

The candidate elements that show the most pot ential for identifying redox controls over the expected range of values at the ZERT site appear to be Mn, Fe, Cr, Cu and Se. In the following evaluation, only Fe and Mn are addressed. The be havior of Cr, Cu, and Se are addressed in Appendix D. In this evaluation, the concentrations of all species in solution are held constant, in conformity with the average composition given in Table 5.4, regardless of the oxidation state specified by the pe. The calculations assume complete homogeneous e quilibrium with respect to the groundwater, and the saturation index (SI) for each redox-sensitive solid phase is calculated. If the trend of calculated SI’s of a give n redox-sensitive phase as a functi on of pe intersects the zero axis, i.e., SI = 0, then that phase is in equilibrium with the groundwater at the pe where SI = 0, regardless of the number of essential redox-sensitive components constituting that phase. For example, consider the mineral bornite (Cu 5 

### FeS 4 

### ) (Appendix D, Figure D.2). In th is mineral, all three elements can take up different equilibrium oxidation states in solution, depending on pe, as is illustrated in Table 5.6, i.e., Cu(II)/Cu(I), Fe(III) /Fe(II) an S(VI)/S(-II). The translation equation defining the solubility product of bornite in Da ta0.ymp.R4 of EQ3/6 is given by:

Cu 5 

### FeS 4 

### + 4H + 

### = Cu + 

### + Cu +2 

### + Fe +2 

### +4HS - 

### (5-10)

Each redox-sensitive species is also related to the primary basis redox-sensitive species by other translation equations defining th e appropriate dissociation consta nts relating redox pairs, thus:

Cu + 

### + H + 

### +1/4O 2 

### (g) = Cu +2 

### +1/2H 2 

### O (5-11)

Fe +3 

### + 1/2H 2 

### O =1/4O 2 

### (g) + Fe +2 

### + H + 

### (5-12)

HS - 

### + 2O 2 

### (g) = SO 4 -2 

### + H + 

### (5-13)

where

O 2 

### (g) +4H + 

### + 4e - 

### = 2H 2 

### O; (5-14)

pe  log  K  log O 2 g 

###   

##   1 2 H  

##  

H

2 O 

##   1 2 

### (5-15)

When EQ3/6 calculates the distribut ion of species in the aqueous pha se, it ensures that all aqueous redox pairs are in a state of homogeneous equilibrium at the sp ecified pe, defined in relation to

O 2 

### (g) in the above-cited equation. Ther efore, when SI(bornite) is calc ulated, it does so in relation to all equilibrated redox-sensitive species at the specif ied pe. In the following figures, and those given

### 48 in Appendix D, the SIs of redox-sensitive minerals for a given redox-sensitive component can be compared, regardless of the presence of any other redox-sensitive essential components in any given phase.

In the following evaluation, the purpose is to iden tify coexisting pairs of redox-sensitive minerals that buffer the pe (Eh) of the system at equilibrium , i.e., where the stability fields of each intersect each other at equilibrium with the coexisti ng groundwater, which can be recognized by the intersection of their respective SIs, where both SIs = 0. In real ity, the intersection may occur incrementally above or below SI = 0, because of un certainties in the thermochemical data used in the analyses, analytical errors (which are accentuated near the detection limits for elements in solution), and other factors, such as temperatur e, ionic strength, and accoun tability of all aqueous complexes of the participating species.

The saturation indices of relevant iron and manganes e phases as a function of pe were tabulated and the data graphed as shown in Figures 5.3 and 5.4, respectively. To make meaningful comparisons

between the saturation indices of different minerals of a given element, each saturation index was normalized to represent the stoi chiometry of a mineral containi ng one mole of the referenced element, e.g., in the iron system, SI(magnetite (Fe 3 

### O 4 

### )) is divided by 3 to yield SI(FeO 1.333 

### ).

Table 5.6. Variation in principal aqueous species of selected redox-sensitive species as a function of pe at circumneutral pH (pH = 6.9).

Dominant Aqueous Species at Specified pe Aqueous Species 14 12 10 8 6 4 2 0 -2

Mn Mn +2 Mn +2 Mn +2 Mn +2 Mn +2 Mn +2 Mn +2 Mn +2 Mn +2

Fe FeO + FeO + FeO + FeO + Fe +2 Fe +2 Fe +2 Fe +2 Fe +2

Co Co +2 Co +2 Co +2 Co +2 Co +2 Co +2 Co +2 Co +2 Co +2

Cu Cu +2 Cu +2 Cu +2 Cu +2 Cu +2 Cu + Cu + Cu + Cu +

Cd Cd +2 Cd +2 Cd +2 Cd +2 Cd +2 Cd +2 Cd +2 Cd +2 Cd +2

Pb PbCO 3 PbCO 3 PbCO 3 PbCO 3 PbCO 3 PbCO 3 PbSe PbSe PbSe

Cr CrO 4 -2 CrO 4 -2 CrO 4 -2 CrO + CrO + CrO + CrO + CrO + CrO +

As FeAsO 4 (aq) FeAsO 4 (aq) FeAsO 4 (aq) FeAsO 4 (aq) FeAsO 4 (aq) FeAsO 4 (aq) FeAsO 4 (aq) AsSe(OH)(SeH)

- AsSe(OH)(SeH)

-

Se SeO 4 -2 SeO 4 -2 SeO 4 -2 HSeO 3 - HSeO 3 - HSeO 3 - ZnSe FeSe FeSe

Mo MoO 4 -2 MoO 4 -2 MoO 4 -2 MoO 4 -2 MoO 4 -2 MoO 4 -2 MoO 4 -2 MoO 4 -2 MoO 4 -2

U UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 (CO 3 ) 2 -2 UO 2 Zn Zn +2 Zn +2 Zn +2 Zn +2 Zn +2 Zn +2 Zn +2 Zn(SCN) 4 -2 Zn(SCN) 4 -2

NO3 NO 3 - N 2 N 2 N 2 N 2 N 2 N 2 N 2 N 2 PO4 H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - H 2 PO 4 - SO4 SO 4 -2 SO 4 -2 SO 4 -2 SO 4 -2 SO 4 -2 SO 4 -2 SO 4 -2 SO 4 -2 SO 4 -2

HCO3 HCO 3 - HCO 3 - HCO 3 - HCO 3 - HCO 3 - HCO 3 - HCO 3 - HCO 3 - HCO 3 -

### 49 Figure 5.3. Normalized saturation indices of magnetite, goethite, siderite, green rust (fougerite-CO 3 ) and Fe(OH) 3 (ferrihydrite) plotted as a function of pe.



### Inspection of Figure 5.3 indicates that the phase boundary between fougerite-CO 3 

### and Fe(OH) 3 

### (ferrihydrite is located where SI(fougerite-CO 2 

### ) –SI(Fe(OH) 3 

### ) ≈ 0 at pH =6.82 and pe ≈ +9 (Eh ≈ +0.51 V). This oxidation potential is somewhat higher than that calculated for the boundary between green rust sulfate where pe ≈ +6 at the same pH, as calcu lated by Majzlan et al. (2004). However, these investigators used the thermodynamic properties for gr een rust sulfate estimated by Bourrie et al. (1999), which ignored the pres ence of Mg(II) substitution for Fe(II) in the trioctahedral sheets. This substitution considerably e xpands the stability field of naturally occurring fougerite in pe-pH space (Bourrie et al., 2004).

Given the reactive nature of both fougerite-CO 3 

### and ferrihydrite, and the potentially rapid transformation between one mineral and the other, we assert that the redox potential is most likely to be buffered through the coexistence of these two minerals. Note that detrital magnetite is also observed in the ZERT sediments (Section 4), an d both fougerite and ferrihydrite could be reductively transformed to magnetite within th e pe range between +2.5 and +13. However, the reaction is likely to be inhibi ted kinetically. The groundwater woul d also be supersaturated with respect to goethite at any pe ≥ 1. However, it is well recognized that the presence of SiO 2 

### (aq) inhibits the transformation of green rust and fe rrihydrite to more stable oxy-hydroxides (Kwon et al., 2007). Finally, it should be noted that side rite remains unsaturate d, i.e., SI(siderite) ≈ -2, even under reducing conditions where all dissolved Fe(T) is transformed to Fe(II), i.e., pe ≤ +4. In this context, it should be mentioned th at at greater unsampled depths below the water table, reducing

conditions could be such that a ll Fe(III) in metastable iron oxides, i.e., ferrihydrite and green rust, and presumed to be prevalent near the interface w ith the vadose zone, could be reduced to Fe(II) and transformed to siderite under ambient P(CO 2 

### ).

### 50 Figure 5.4. Normalized saturation indices of pyrolusite, bixbyite, manganite, hausmannite and rhodochrosite, plotted as a function of pe.



### Examination of Figure 5.4, which displays the calcu lated normalized saturation indices of various manganese oxides, shows that pyrolusite (MnO 2 

### ) would saturate at pe ≈ +11.5, bixbyite (Mn 2 

### O 3 

### ) and manganite (MnOOH) at pe ≈ +12, and hausmannite (Mn 3 

### O 4 

### ) at pe ≈ +14.5. Rhodochrosite remains undersaturated, that is, SI(rhodochrosite) ≈ -1.8, throughout the examined range of pe values (-2 < pe < +14). Thus, if any of the mi nerals pyrolusite, bixbyite, manganite, or hausmannite were present as pure substances , and in equilibrium with the groundwater, the pe would, at a

minimum be ≈ +11 (for pyrolusite). However, given th at the average crustal abundance of Mn is about 1.7% that of Fe, it is likely th at most, if not all Mn is presen t in solid solution in reactive Fe oxy-hydroxides such as ferrihydrite or goethite, e. g., see Sileo et al. (2001) . Thus, assuming an ideal solid solution in ferrihydrite, Mn +2 

### ions would be in equilibrium with a solid solution at pe ≈ 12 – 1.8 = 10.2, which is about 1 pe unit higher th at the pe estimated from the pe buffering equilibrium between fougerite-CO 3 

### and Fe(OH) 3 

### (ferrihydrite). Note that in the presence of significant concentrations of Mn in sediments, Fe(II) can reduce pyrolusite with concomitant precipitation of the manganese spinel, jacobsite (MnFe 2 

### O 4 

### ), e.g., see Villinski et al. (2000).

In summary, the foregoing analysis of the satu ration indices of redox-sensitive minerals indicates that the redox potential of the system is to a large extent controlled by the metastable redox buffering equilibrium between fougerite-CO 3 

### and ferrihydrite at a pe circa +9 (Eh ≈ +0.51 V). The

observed initial concen tration of Mn in solution is consistent with saturation with manganese oxide minerals at an oxidation state pe ≤ +11.5, but because the average cr ustal abundance of Mn is less than that of iron, it is likely that the system pe with respect to Mn is determined by a solid solution of Mn(II) in ferrihydrite, and theref ore the pe would be << +11.5. Equilibrium between potential Cr oxide minerals suggests a pe-buffering phase equ ilibrium between Cr-ferri hydrite and chromite at

pe ≈ +8.7, although both the thermodynamic data us ed in the interpretation and the potential involvement of solid solution substitution in Fe(III) bearing phases such as ferrihydrite discounts

### 51 the significance of this observati on. Nevertheless, the behavior of th e Cr system is consistent with

those of Fe and Mn.

The preponderance of the evidence from the foregoi ng analysis is that the pe of the system is probably buffered at approximately +9 ±1 (see also Appendix D rega rding Cr minerals). This value is close to the pe value of +10 of the local partial equilibrium cl uster of redox pairs controlled by

nitrate observed in shallow Georgia groundwaters by Washingt on et al. (2004). However,

examination of the behavior of Cu and Se minera ls as a function of pe (see Appendix D) suggests that organic matter might also be playing a role with respect to the establishment of a lower pe cluster of redox pairs at pe ≈ +1. This compares with an average value of pe ≈ -2 determined by Washington et al. (2004)

Note the distribution of Cr, Fe and Se species in solution at pe = +9, as shown in the following Table 5.7. With respect to Cr, Cr(VI) and Cr(III) sp ecies are present in s ub-equal concentrations, i.e., 82 mol.% versus 18 mol.%. A similar situa tion obtains with respect to Se(VI) versus Se(IV) species with respective c oncentrations of 85.1 and 15 mol.%. W ith respect to Fe, however, Fe(III) dominates Fe(II) by about 10 3 

### . It is possible that all three species are impacted by NO 3 - 

### /NO 2 - 

### in solution, although unfortunatel y, the concentration of NO 2 - 

### was not measured in groundwater samples from the ZERT site.

Table 5.7. Distribution of the predominant Cr, Fe and Se species in solution at pH = 6.82 and pe = +9

Species Molality Percentage

Chromium Species

CrO 4 -2 1.99E-08 43.4

CaCrO 4 (aq) 1.11E-08 24.2

HCrO 4 - 6.55E-09 14.31

CrO + 6.43E-09 14.04

CrOH +2 1.80E-09 3.92

HCrO 2 (aq) 4.58E-11 1.00E-01

Cr +3 8.91E-12 1.95E-02

Iron Species

FeO + 1.28E-07 41.83

HFeO 2 (aq) 1.05E-07 34.49 Fe(OH) 2 F(aq) 5.44E-08 17.83 Fe(OH) 3 H 3 SiO 4 - 1.73E-08 5.68

FeOH +2 1.49E-10 4.89E-02

Fe +2 1.33E-10 4.36E-02

Selenium Species

SeO 4 -2 2.59E-08 85.1

HSeO 3 - 3.16E-09 10.38

SeO 3 -2 1.37E-09 4.51

### 52 5.2.3 Evolution of Calcite Saturation I ndices During Injection

Examination of core material from the ZERT site shows that coatings of secondary calcite are ubiquitous throughout the section, including that section below the water table from which groundwater samples were drawn for analysis. It would therefore not be surprising if a thermodynamic analysis showed that initial groundw ater samples taken prior to injection of carbon dioxide were saturated with respec t to calcite. However, it must be recognized that the equilibrium

state of the shallow subsurface environment at that ZERT site is subject to perturbations due to seasonal and long-term variations in the following parameters:

 Fluctuations in temperature

 Groundwater recharge, leading to fluctuati ons in the height of the water table.

 Infiltration of rainwater

 Evapotranspiration from the vadose zone

 Capillary transport

 Bio-metabolism

That some or all of these proce sses are operative is evident from the detection of a caliche horizon within the vadose zone approximately one meter be low the surface. An extensive body of literature exists which describes the listed processes in relation to the formation of caliche zones under various climatic conditions, and which could aid in interpreting how these processes might impact groundwater composition and trace element distributions at the ZERT site. However, time limitations permit only a cursory examination of ch anges in calcite saturation in response to CO 2 

### injection during the test.

Table 5.8 shows the calculate d saturation indices of cal cite and calculated P(CO 2 

### ) for each groundwater sample based on a thermodynamic analysis using the code EQ3/6 (Wolery, 1992, 1993). The groundwater samples are those taken betw een the start of the sampling program and the first significant rainfa ll, i.e., between July 7 and July 17, 2008. The calculations incorporate an adjustment to the analytical value of HCO 3 

### to account for bicarbonate complexation, and entry of

the corrected HCO 3 - 

### as a “free” concentration to force the code to calculate the concentration of

dissolved CO 2 

### in equilibrium with the “free” bicar bonate concentration. Each calculation was conducted at the measured temperature for each groundwater ( ≈ 10oC). Also included in Table 5.8 is the calculated charge imbalance, as a percentage of total charge for each groundwater chemical analysis. The charge imbalances are, with the exception of samples 124 and 128, within the acceptable limits of ±5%, and might ge nerally be consider ed as excellent.

Figure 5.5 is a graphical representa tion of the same data. It is cl ear that the increase in ambient P(CO 2 

### ) leads to undersaturation of the groundwater with respect to calcite and the rate of

dissolution of the calcite is insufficient to mainta in equilibrium. The trend is irregular and shows considerable scatter, which would be expected of a system where samples are taken from different wells at different depths in a soil system that is vertically heter ogeneous, and probably also laterally heterogeneous. Also plotted on Figure 5.5 are three trend lines. Trend Line A assumes the approximate path taken with respec t to calcite resaturation as P(CO 2 

### ) increases. This path would vary depending on the rate of CO 2 

### saturation of the groundwater, the kinetics of calcite dissolution, calcite surface area and abundance, and groundwater fl ow rates. Path A is constrained by the other two trend lines. Trend Line B is the path taken un der static conditions wher e resaturation of calcite

occurs after instanta neous saturation of the groundwater with CO 2 

### at a pressure of 1 bar, the

### 53 approximate limit of saturation expected at the ZE RT site. Trendline C assumes that no dissolution of calcite takes place upon uptake of CO 2 

### . It is evident that about ha lf of the water samples taken between July 7 and July 17, 2008 reflect this state of affairs (close to trendline C), and that only after log P(CO 2 

### ) exceeded ≈ -0.7, equivalent to 0.2 bar, did signi ficant dissolution of calcite occur.

From the start of the test until the first significan t rainfall, the water table fell progressively, which suggests that the water table was still responding to some rainfall event preceding the start of the test, due either to drainage, or to evapotranspiration. D ilution by rainfall could explain the initial degree of undersaturation of calcite, and this observation is i ndirectly corroborated by the observed undersaturation with respect to opal-CT and am orphous hydroxy-aluminosilicates (HAS). For this reason, it is less likely that the initial undersaturation might be due to transient elevation of P(CO 2 

### ) due to aggressive mineralization of carbonaceous material by bacteria.

Table 5.8. Calculated calcite saturation index and CO 2 partial pressure for groundwaters taken at the ZERT site between July 7 and July 17, 2008

ZERT Sample

No. Charge

Imbalance.

% SI(Calcite) (-) Log P(CO 2 ) (bars)

102 -1.20 -0.377 -1.3113

103 -4.10 -0.342 -1.2739

104 -0.30 -0.156 -1.3567

105 -1.73 +0.691 -0.951

106 0.15 -0.301 -1.4203

107 0.65 -0.524 -1.2369

109 0.96 -0.090 -1.4708

112 0.74 -0.566 -0.9664

116 -2.00 -0.369 -0.3542

117 2.94 -1.237 -0.0961

118 0.78 -0.415 -0.7166

119 0.08 -0.178 -1.5106

124 -7.69 -0.492 -0.0881

126 0.63 -0.781 -0.7594

127 0.09 -0.856 -0.1881

128 8.11 -0.715 -0.1458

129 -2.95 -0.588 -0.1579

130 -1.20 -0.759 -0.7421

131 -1.30 -0.827 -0.1648

132 -2.39 -0.692 -0.09

### 54 Figure 5.5. Saturation index of calcite as a function of log P(CO 2 ) calculated for groundwater samples taken between July 7 and July 17, 2008 at the ZERT site. For an explanation of the trend lines, please refer to

the text.



### 5.2.4 Findings from Ge ochemical Evaluations

This section summarizes the prin cipal findings of the foregoing geochemical evaluations. These findings have been useful in setting up the ge ochemical model described in Section 5.3 below.

 Principal Component Analysis of the initial pertur bation (first 10 days of the test) in groundwater quality due to CO 2 

### uptake shows strong positive lin ear correlations between the transformed logarithmic concentr ations of many of th e measured species. The first principal component, PC1, correlates strongly with monov alent, divalent and trivalent cations, suggesting that ion exchange on smectite mi ght be the mechanism responsible for their release. The positive correlation between various anionic species: HCO 3 

### , SO 4 

### , Mo (as MO 4 

### ) and Se (as SeO 3 

### or SeO 4 

### ) and PC1 also suggest a common mechanism involving either ion exchange or adsorption. However, NO 3 

### and PO 4 

### are negatively correlated. The second principal component, PC2, correlate s strongly with F, and relati vely weakly with As, Se and Mo, and inversely, but weakly with H and K, suggesting some anionic exchange or adsorption process. The third principal compone nt, PC3, correlates i nversely with Mn, Co and Cl, and directly with As, NO 3 

### and PO 4 

### , suggesting precipitation control by a transition metal phosphate. PC4 correlates with Cl and SO 4 

### , which may be related to earlier infiltration of rain water from the overlying vadose zone . Not all correlation trendlines between the various species are readily expl ained, highlighting the need fo r more in-depth study of the geochemistry of shallow groundwater systems (Section 5.2.1).

 Thermodynamic evaluation of the groundwater chemistry before CO 2 

### injection reveals a species distribution that reflect s rapid weathering and the fo rmation of metastable hydroxy-

### 55 aluminum silicates. The initial groundwater is s lightly undersaturated wi th respect to calcite, opal-CT and barite, suggesting d ilution of the groundwater by ra inwater prior to the test. Supersaturation with respect to smectites, illite and kaolinite and gibbsite is indicative of saturation by persistent metastable h ydroxy-aluminum silicates (Section 5.2.2).

 Examination of the degree of supersaturation of redox sensitive minerals as a function of pe (or Eh) suggests that the re dox potential is most likely buffered by the coexistence of fougerite-CO 3 

### and Fe(OH) 3 

### (ferrihydrite) at a pe ≈ +9. However, there are also indications

that the concentrations of some trace elem ents, e.g., Cd, Cu, Pb, may be controlled by precipitation of reduced species such as CdSe, Cu 2 

### O and PbSe. This suggests that decaying organic matter might be exercising reductive co ntrol over these, and possibly other species, which implies that the system redox state is no t in equilibrium, and may also be spatially inhomogeneous. The distribution of species, th erefore, may be ki netically controlled (Section 5.2.2).

 Examination of the perturbation in calc ite saturation following injection of CO 2 

### shows that calcite immediately became si gnificantly undersaturated, but recovery of the saturation states was slow and incomplete during the first ten days of the test (Section 5.2.3).

5.3 Geochemical Modeling

Geochemical simulations were run to test some of the observations discussed earlier, namely that ion exchange and sorption processes, together with dissolution of carbona tes, could explain the water chemistry trends observed during the test (Zheng et al., 2009a). Additional modeling work incorporating transport processes was not warranted because of the relatively limited set of field data, as well as uncertainties regarding groundwater flow and CO 2 

### transport (e.g., Section 5.1), subsurface heterogeneity, and rainfall effects. It should be noted that geochemical changes in groundwater after the release of CO 2 

### were fairly rapid and even, o ccurring almost simultaneously in wells 1B, 2B and 5B (Section 3). Furthermor e, PCA analyses (Section 5.2) reveal similar groundwater chemistry responses to CO 2 

### injection at the different monitored locations. This suggests that “batch” simulations, focused solely on reactive processes indu ced by increasing levels

of CO 2 

### , are a good modeling approach in this case, with little more to gain from more sophisticated simulations including transport. The simulations were performed using, TOUGHREACT (Xu et al., 2006), which was enhanced through incorporatio n of surface complexation and ion exchange models to compute the distribution of adsorbed chemical species on mineral surfaces (Zheng et al., 2009b).

Rainfall was shown to significantly affect groundw ater chemistry during the test (Section 3).

Because of the limited amount of rainfall-related data (including composition, recharge rates, and oxidation/mixing effects in the subsurface), and because PCA analyses (Section 5.2.1) focused on the groundwater chemistry data before the first rainfall, the geochemical simulations discussed below were based solely on groundwater data co llected before the first rainfall.

5.3.1 Model setup

A “batch” (single-gridblock) geochemical model was develo ped, with ini tial conditions representative of field conditions prior to CO 2 

### injection (Table 5.9). CO 2 

### was then (numerically)

### 56 injected into the modeled system, and model re sults were examined as a function of pH. As discussed previously, the CO 2 

### volume fractions measured in the head space of groundwater monitoring wells increased rapidl y (within a day to a few days) to levels up to 0.9 (90% by volume), corresponding to P(CO 2 

### ) values near 1 bar at the wate r table soon after injection. CO 2 

### was injected into the model gridbloc k until the computed pH value fell from an initial value ~7.1 to a value near 5.8, consistent with the field meas urements. At this point, the calculated maximum P(CO 2 

### ) was around 1.5 bar. Instantaneous equilibration of gaseous CO 2 

### with the aqueous phase was assumed. Other elements of the model are described below. Simulations were carried out assuming an initial temperature of 10°C and an init ial pressure close to atmospheric values.

Initial conditions

Seven groundwater samples were taken before the release of CO 2 

### . The initial chemical composition for the model (Table 5.9) was taken from one of these samples (i.e., 08ZERT-107), with initial pH adjusted slightly from 6.7 to 7.1 so that our calculation covers the pH range of all samples considered in this study. An initial aqueous sp eciation was conducted w ith EQ3 (Wolery, 1993) to verify charge balance and evaluate the water saturation with respect to several minerals. The initial partial pressure of CO 2 

### in the groundwater was calculated to be 2.5×10 -2 

### bars. Furthermore, calcite was close to saturation, i.e., SI(calcite) ≈ –0.1 (–0.3 at pH 6.7).

As pointed out in Section 5.2, redox is not likely to have reached equilibrium in this shallow, lowtemperature system. This complicates the use of ge ochemical models such as the one applied here, because these models rely on the assumption of re dox equilibrium to calculate the distribution of redox-sensitive species. Here, th e initial redox state of the groundwater is calculated from the Fe(II)/Fe(III) equilibrium using the measured (t otal) Fe concentration in water and assuming equilibrium with goethite, yielding a pe ~ 1.7. A ssuming equilibrium with ferrihydrite and fougerite (green rust) would result in a si gnificantly higher valu e (pe ~9) (Section 5.2.2) . This range of pe values suggests suboxic (but not strongly reducing) conditions pr evail at the site. Note that measured Eh values (150–200 mV, Section 3) co rrespond to pe ~ 2.5–3.5. However, as discussed in Section 5.2.2, measured potentials represent mixed potentials that are not representative of a specific redox couple and that may be biased by the redox sensor. Gi ven these uncertainties, and the fact that redox equilibrium is assumed in the simulations, mode l results regarding redox species should be viewed as qualitative ra ther than quantitative. Obviousl y, the choice of pe affects model results for redox-sensitive species. However, in this modeling case, the amount of solid Fe for sorption is unknown and adjusted to match observed groundwater compositions. Therefore, assuming a higher pe would likely be compensate d by this calibration such that, qualitatively, model results would not be expected to cha nge significantly if a di fferent pe was used.

### 57 Table 5.9. Primary species and their initial concentration

Species Concentration (mol/L) Species Concentration (mol/L)
pH 7.1 Ba +2 7.9E-07

Ca +2 1.9E-03 Cd +2 2.0E-09
Mg +2 9.8E-04 Co +2 4.2E-09

Na + 3.3E-04 Cu +2 2.4E-08

K + 1.4E-04 H 2 AsO 4 - 1.3E-08

Fe +2 1.8E-09 Mn +2 1.8E-08

AlO 2 - 1.2E-07 Sr +2 3.0E-06

SiO 2 (aq) 5.3E-04 Zn +2 4.6E-08

Cl

- 1.3E-04 HSeO 3 - 4.9E-08

HCO 3 - 6.3E-03 MO 4 -2 4.6E-09

SO 4 -2 7.7E-05 UO 2 +2 1.7E-08

NO 3 - 9.6E-06 Cr(OH) 2 + 2.2E-07

O 2 (aq) 1 1.2E-51 HPO 4 -2 1.8E-06

Pb +2 2.4E-10 Li + 7.7E-07

1 Corresponds to pe ~ 1.7, Eh ~ 100 mV (see text)



### Mineralogy

Section 4.2 describes the mineralogical analyses of cored sections to a depth of approximately 3 m. Sediments in the saturated zone are composed pr edominantly of a matrix of cobble-sized rounded fragments, consisting primarily of andesitic volca nic rocks with predominant plagioclase, quartzbiotite-amphibole gneiss, and red granite/gneiss. The mafic volcanic rock fragments are very high in magnetic iron. Minor limestone and occasional dolomite fragments are also observed in the matrix. Unspecified Fe oxides appear in a <0.25 mm fr action as “powdery orange blebs.” Secondary carbonate coatings are found through out the core, and a caliche zone is present in the overlying unsaturated zone. Small quantities of clay (~1 wt%), believed to consist of smectite and minor

kaolinite are also found in the sa turated zone. The list of minerals considered in the geochemical model (Table 5.10) was based on these observations , as well as on the disc ussions in Section 5.2. Arbitrary volume fractions are assigned to each mineral, which are roughly consistent with the mineralogy analyses. It should be noted that in the time frame of the simulation described in this report, most of the listed rock-for ming minerals are essentially inert, i.e., they are mostly unreactive during the short in jection period, as discussed later.

Table 5.10. Minerals considered in the model.

Primary Mineral Volume fraction Primary Mineral Volume fraction

Quartz 0.45 Ferrihydrite 0.01 K-feldspar 0.1 Magnetite 0.02 Oligoclase 0.34 Goethite 0.001

Smectite 0.01 Siderite 0.001

Dolomite 0.02 Calcite 0.05



### Chemical reactions and equilibrium/kinetic constraints

Chemical reactions considered in the model include aqueous complexation, cation exchange, adsorption/desorption, and mineral dissolution/pr ecipitation. Of these, cation exchange and

### 58 adsorption/desorption reactions are believed to predominate (see Section 5.2.1). Table 5.11 lists aqueous complexes included in the simulation an d their dissociation constants which are from database data0.ymp.R4 (Wolery and Jareck, 2003 ; Wolery and Jove-Colon, 2007). Table 5.12 lists the cation exchange reactions and their selectivity coefficients. The Gaines-Thomas convention is used here. The selectivity coefficient for exchangeable H + 

### is taken from Charlet and Tournassat (2005), and those for the remainder of the exchangeable cations are taken from Appelo and Postma (1994). The cation exchange capacity (CEC) is the key parameter for the calculation of cation exchange. Because the measured CEC for the sediment s at ZERT site is not currently available, we used a CEC value of around 12 meq/100g measured for Hanford sediments (Steefel et al., 2003), which has similar composition to the sediment s at the ZERT site. A nonelectrostatic surface complexation model calculates adsorption/deso rption, which involves most anions and H + 

### . Table 5.13 lists the surface complexation reactions. Su rface complexation constants are taken from Dzombak and Morel (1990). Hydrous ferric iron oxide (HFO) is assumed to be the only adsorbent. The site density for the strong site is assumed to be 1.76×10 -6 

### mol/m 2 

### , and that for weak sites is

assumed to be 3.22×10 -6 

### mol/m 2 

### (Muller and Sigg, 1991). The specific surface area is assumed to be

14.7 m 2 

### /g (Muller and Sigg, 1991). The amount of HFO for sorption was (loosely) calibrated to metal concentrations in gr oundwater using these data.

Table 5.11. Aqueous complexes and thei r dissociation constants

Species Logk (25 o C) Species Logk (25 o C) Species Logk (25 o C) Species Logk (25 o C)

OH - 13.991 Al(OH) 2 + -12.78 Cu(CO 3 ) 2 -2 10.476 SrOH + 13.292

CaHCO 3 + -1.043 AlOH +2 -17.868 NaHCO 3 (aq) -0.17 UO 2 (CO 3 ) 2 -2 3.747

FeHCO 3 + -2.044 Al 13 O 4 (OH) 24 +7 -198.749 NaSO 4 - -0.811 UO 2 (CO 3 ) 3 -4 9.43

FeCO 3 (aq) 4.879 CdCl + -1.966 NaCl(aq) 0.782 UO 2 (CO 3 ) (aq) 0.663

CO 2 (aq) -6.342 HCOO 2 - 21.243 PbCl + -1.456 UO 2 (OH) 2 (aq) 10.315

CO 3 -2 10.325 HAsO 4 -2 6.754 PbCl 2 (aq) -2.016 Cr(OH) 3 (aq) 8.3

MgHCO 3 + -1.033 MnCO 3 (aq) 5.809 SrSO 4 (aq) -2.271 Cr(OH) +2 -5.7

MgSO 4 (aq) -2.382 MnHCO 3 + -0.882 SrCO 3 (aq) 7.47 CaHPO 4 (aq) -2.69

CaSO 4 (aq) -2.1 BaCO 3 (aq) 7.691 SrCl + 0.253 MgHPO 4 (aq) -2.91

HAlO 2 (aq) -6.448 CuOH + 7.287 SrNO 3 + -0.796

### 59 Table 5.12. Cation exchange reactions and selectivity coefficients.

Cation exchange reaction K Na/M

Na + + X-H = X-Na + H + 1

Na + + X-K = X-Na + K + 0.2

Na + + 0.5X-Ca = X-Na + 0.5Ca +2 0.4

Na + + 0.5X-Mg = X-Na + 0.5Mg +2 0.45

Na + + 0.5X-Pb = X-Na + 0.5Pb +2 0.4

Na + + 0.5X-Ba = X-Na + 0.5Ba +2 0.35

Na + + 0.5X-Cd = X-Na + 0.5Cd +2 0.4

Na + + 0.5X-Co = X-Na + 0.5Co +2 0.6

Na + + 0.5X-Cu = X-Na + 0.5Cu +2 0.5

Na + + 0.5X-Mn = X-Na + 0.5Mn +2 0.55

Na + +0.5X-Sr = X-Na + 0.5Sr +2 0.35

Na + +0.5X-Zn = X-Na + 0.5Zn +2 0.4

Na + +0.5X-Fe = X-Na + 0.5Fe +2 0.6

Table 5.13. Surface complexation reactions and surface complexation constants (logK) on HFO

Surface complexation logK

HFO_sOH 2 + = HFO_sOH + H + -7.29
HFO_wOH 2 + = HFO_wOH + H + -7.29

HFO_sO

- + H + = HFO_sOH 8.93

HFO_wO

- + H + = HFO_wOH 8.93
HFO_sOHSO 4 -2 = HFO_sOH + SO 4 -2 -0.79
HFO_wOHSO 4 -2 = HFO_wOH + SO 4 -2 -0.79
HFO_sSO 4 - + H 2 O = HFO_sOH + SO 4 -2 + H + -7.78
HFO_wSO 4 - + H 2 O = HFO_sOH + SO 4 -2 + H + -7.78
HFO_sSeO 3 - + H 2 O = HFO_sOH +HSeO 3 - -4.29
HFO_wSeO 3 - + H 2 O = HFO_wOH +HSeO 3 - -4.29
HFO_sMoO 4 - + H 2 O = HFO_sOH + MoO 4 - + H + -9.5
HFO_wMoO 4 - + H 2 O = HFO_wOH + MoO 4 - + H + -9.5
HFO_sOHMoO 4 -2 = HFO_sOH + MoO 4 -2 -2.4
HFO_wOHMoO 4 -2 = HFO_wOH + MoO 4 -2 -2.4
HFO_sH 2 AsO 4 + H 2 O = HFO_sOH + H 2 AsO 4 - + H + -10.17
HFO_wH 2 AsO 4 + H 2 O = HFO_wOH + H 2 AsO 4 - + H + -10.17
HFO_sHAsO 4 - + H 2 O = HFO_sOH + H 2 AsO 4 - 0.35
HFO_wHAsO 4 - + H 2 O = HFO_wOH + H 2 AsO 4 - 0.35
HFO_sH 2 PO 4 + H 2 O = HFO_sOH + HPO 4 -2 + 2H + -18.9
HFO_w H 2 PO 4 + H 2 O = HFO_wOH + HPO 4 -2 + 2H + -18.9
HFO_sPO 4 -2 + H 2 O = HFO_sOH + HPO 4 -2 -5.4
HFO_wPO 4 -2 + H 2 O = HFO_wOH + HPO 4 -2 -5.4
HFO_wCO 2 - + H 2 O = HFO_sOH + HCO 3 - -2.45
HFO_wCO 2 H + H 2 O = HFO_sOH + HCO 3 - + H + -2.45

### 60 All minerals are set to react under kinetic constraints. Currently, TOUGHREACT (Xu et al., 2006) uses a general form of kinetic rate law based on the transition state theory (TST) (Lasaga et al., 1994; Steefel and Lasaga, 1994):



##   1 / r kA Q K    

      

### (5-16)

where r is the kinetic rate (positiv e values indicate dissolution, negative values precipitation), k is the rate constant (moles per unit mineral su rface area and unit time), which is temperature dependent, A is the specific reac tive surface area per kg H 2 

### O, K is the equilibrium constant for the mineral–water reaction, and Q is the reaction quotient. Here, the parameters θ and η are assumed equal to unity, and surface areas are estimated ba sed on Xu et al. (2006). For further details concerning the kinetics of minera l dissolution/precipitation, the r eader is referred to Xu et al. (2006).

Tables 5.14 and 5.15 list the equi librium constants (as log( K

)) and rate constants ( k ), respectively, for minerals included in the simulations. Except for siderite, equilibrium constants are from database data0.dat.ymp.R4 (Wolery and Ja reck, 2003; Wolery and Jove-Colon, 2007). The equilibrium constant for siderite was calculated from more recent data fr om Preis and Gamsjager (2002), and differs slightly from that adopted in data0.ymp.R4. The value of Q/K in Equation 5-16 dictates whether dissolution or pr ecipitation occurs. In the presen t case, all minerals dissolve. However, except for calcite and siderite, the calcul ated dissolution rates are very small (relative to the length of the te st) and do not affect model results.

The dissolution rate of calcite is critical, because the increase of Ca +2 

### resulting from calcite dissolution affects cation exchange, which may be re sponsible for the increases in major cations and most trace metals except for Fe and possibly Mn and Co (Section 5.3.2). Depending on the level of undersaturation, according to Kaufmann and Dreybrodt (2007), the dissolution of calcite can be described by a fast and a slow rate law. The fast and slow rate laws are discriminate d by the ratio of

calcium concentration in the solution to the calc ium concentration in equilibrium with calcite (C/Ceq) at a given CO 2 

### partial pressure. Kaufmann and Dre ybrodt (2007) reported that when the ratio (C/Ceq) is <0.3, the fast rate law applies, wh ereas when C/Ceq > 0.3, the slow rate law applies. In the current simulations, C/Ceq is about 0.4–0.5, and thus the slow rate law of Kaufmann and Dreybrodt (2007) is adopted, with the rate constant k (Equation 5-16) equal to 1.6×10 -6 

### mol/m 2 

### /s. Svensson and Dreybrodt (1992) investigated the dissolution kinetics of calcite in CO 2 

### -water systems approaching calcite equili brium and gave a rate constant k of 1.6×10 -6 

### – 2.2×10 -6 

### mol/m 2 

### /s, although they indicated that C/Ceq should be around 0.6–0.8. Gledhill and Morse (2006) reported

that the dissolution rate de pends on brine composition, P(CO 2 

### ) (0.1–1 bar), and temperature and the rate constant (k) can be estimated from a mu ltiple regression model. Applying their regression model to our simulation conditions, the rate constant is from 1.6×10 -6 

### to 2.2×10 -5 

### mol/m 2 

### /s. Jordan and Rammensee (1998) gave a dissolution rate of 1.6×10 -6 

### mol/m 2 

### /s on the calcite surface obtained by scanning force microscopy, which corresponds well to the rates obtained from batch experiments. Pokrovsky et al. (2005) measured the dissolu tion rate of calcite at high pressure, which is more relevant to deep CO 2 

### geological storage than CO 2 

### releases at the ZERT site. Therefore, these various published dissolution rates for calcite at lower levels of unders aturation are all quite consistent with the value selected for this study.

### 61 The baseline model considers only cation exchange for iron. However, sensitivity analyses were

conducted to consider the dissolutio n of green rust (fougerite) and si derite. Note that siderite is initially undersaturated (SI ≈ -1.5) and thus unlikely to be pres ent, except possibly as a detrital phase. For siderite, the kinetic rate reported by Duckworth and Martin (2004) is adopted. These investigators studied the dissoluti on of siderite under bot h oxic and anoxic conditions at 298 K, and fitted their measured data to a mechanistic kinetic equation similar to that employed in

TOUGHREACT. The rate constant is given for both neutral and acid dissolution mechanisms:

+

2 

### [H O] [H ]
n H 

### kk k

  

### (5-17)

with k

n 

### = 10 -8.65 

### mol/m 2 

### /s and k H 

### = 10 -4.6 

### mol/m 2 

### /s, respectively. The di ssolution rate is dominated by the neutral mechanism when pH > 5.5, whereas the acid mechanism becomes important when pH < 5.5. Testemale et al. (2009) confirmed that the dissolution rate constant of siderite at acid condition (pH < 4) can be expressed as:

+ 

### [H ]
H 

### k k

 

### (5-18)

with k H 

### = 10 -9.3 

### mol/cm 2 

### /s ( k H 

### = 10 -5.3 

### mol/m 2 

### /s) which is close to the k H 

### given by Duckworth and Martin (2004). Testemale et al. (2009) compared their findings with other data (Braun, 1991; Golubev et al., in press) and concluded that publi shed kinetic rates for si derite dissolution are generally in good agreement. Duckworth and Martin (2004) also mentioned that the dissolution of

siderite is retarded under oxic conditions, due to armoring by Fe (hydr)oxides. At the ZERT site, the proximity of the water table to ground surface and presence of secondary ferric oxy-hydroxides suggest oxic to suboxic conditions may prevail, at least during rainfall events. Therefore, the

siderite dissolution rate adopted for the model could be underestimated.

The initial water is supersaturated with respect to green rust (SI ≈ 0.3, normalized to 1 Fe), although this mineral readily dissolves upon pH decrease. Its dissolution rate is taken from Williams and Scherer (2001) with k

n 

### = 10 -8.82 

### mol/m 2 

### /s and k H 

### = 10 -2.6 

### mol/m 2 

### /s (Equation 5-16).

Table 5.14. Equilibrium constants for minerals 1

Primary

Mineral log( K

) Primary Mineral log( K )

Quartz -3.75 Ferrihydrite -63.24 K-feldspar -22.91 Magnetite -6.505 Oligoclase -97.78 Goethite -8.12

Smectite -39.51 Siderite -0.251

Dolomite 2.524 Calcite 1.85 Fougerite 24.24

1 For dissolution in terms of species shown in Table 5.9

### 62 Table 5.15. Initial volume fraction and kinetic parameters for minerals considered in the model 1

Parameters for Kinetic Rate Law

Neutral Mechanism Acid Mechanism Base Mechanism Mineral Volume

fraction A

(cm 2 /g) k 25

(mol/m 2 /s) k 25
(mol/m 2 /s) n(H + ) k 25
(mol/m 2 /s) n(H + )

Quartz 0.45 9.8 1.023  10 -14

– – – –

K-feldspar 0.1 9.8 3.89  10 -13 8.71  10 -11 0.5 6.31  10 -12 -0.823 Oligoclase 0.338 9.8 1.44  10 -12 2.13  10 -10 0.457 – –

Smectite 0.01 151.6 1.66  10 -13 1.05  10 -11 0.34 3.02  10 -17 -0.4

Dolomite 0.02 9.8 2.95  10 -8 6.46  10 -4 0.5 – –

Ferrihydrite 0.01 12.9 2.34  10 -7 1 – –

Magnetite 0.02 9.8 4.57  10 -10 4.17  10 -7 1 – –

Goethite 0.001 12.9 2.52  10 -12

– – – – Siderite 0.001 22.3 2.24  10 -9 2.5  10 -5 0.75 –

– Calcite 0.05 3.05 1.6  10 -6

– – – – 1 Kinetic data from Xu et al. (2006), except for calcite (Kaufm ann and Dreybrodt 2007) and siderite (Duckworth and Martin, 2004 ), see text.



### 5.3.2 Results and Discussion

Earlier generic studies of CO 2 

### leakage into potable groundwater aquifers (Birkholzer et al., 2008; Wang and Jaffe, 2004; Zheng et al., 2009b) indicated that the dissolution of carbonate minerals (mainly calcite) can significantly buffer the pH decrease accompanying the dissolution of CO 2 

### in water (from the dissociation of car bonic acid). Clay ion exchange rs, when present in significant amounts in an aquifer host rock or sediment, can also strongly buffer pH through exchange of H +


### with the groundwater (Bradbury and Baeyens, 1997; Samper et al., 2008; Zheng and Samper, 2009; Zheng et al., 2009c). In the present model, the di ssolution of calcite, dolomite and siderite, and H +


### buffering through cation exchange and surface comple xation could all contribute to buffering pH. Figure 5.6 shows the calculated pH by different co mbinations of these chemical equilibria. A

simulation without carbonate minerals leads to a lower pH than the base case where the presence of carbonates is assumed, but is only marginally higher than the simulation without any buffering. This suggests that at the ZERT site, dissolution of carbonate minerals is the principal pH buffering process, whereas the H + 

### exchange by clays does not affect pH significantly, because the shallow

saturated zone at the ZERT site soil does not contain large amounts of clay minerals, and the pH decrease is only moderate.

5 5.5

6 6.5

7 7.5

0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.6

Fugacity of CO 2 (bar)pH base case

no carbonate minerals

no carbonate mineral + no protonation

Figure 5.6. Calculated pH with different combination of chemical reactions

### 63 The groundwater at the ZERT site res ponds promptly to the release of CO 2 

### . The pH falls within 1 or 2 days from the start of inject ion, and a significant increase in major cations and trace metals is observed before the first rainfall occurs, about 10 da ys later. In such a shor t-duration test, only the dissolution kinetic rates for rapidl y dissolving minerals such as calcite are important. As noted earlier, the dissolution rate of calcite is critical, because the resulting increase of Ca +2 

### induces cation exchange, which may be responsible for the increase s in major cations and most trace metals except for Fe and possibly Mn and Co. The value adopted in this study falls within the lower range of published values (Section 5.3.1) and yields results consistent with measured changes in calcium

concentration observed in the field (Figure 5.7). If a higher calcite di ssolution rate is used, such as that given by Gledhill and Morse (2006), i.e., 2.2×10 -5 

### mol/m 2 

### /s, the computed calcium

concentration overestimates the measured data (Figure 5.7).

As pH decreases, calcite dissolves and releases Ca +2 

### into the aqueous phase. Aqueous Ca +2 

### in turn exchanges with the exchangeable cations on the cation exchanger, tentatively assumed to be smectite. As a result, the concentrations of all the major cations increase in solution. Figures 5.8– 5.10 show the calculated sodium, potassium, and ma gnesium concentration versus pH. The close agreement between calculated and measured c oncentrations supports the hypothesis that cation exchange may indeed dominate the evoluti on of major cations in response to the CO 2 

### injection at

the ZERT site.

0E+00 1E-03 2E-03 3E-03 4E-03 5E-03 6E-03 7E-03 8E-03 9E-03

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHCa (mol/L) measured data

base model k-calcite = 1.6e-6

sensitivity run k-calcite = 2.2e-5

Figure 5.7. Calculated calcium concentration vs pH with different calcite dissolution rate.

0.E+00 1.E-04 2.E-04 3.E-04 4.E-04 5.E-04 6.E-04

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHNa (mol/L)

measured data

model results

Figure 5.8. Calculated sodium concentration vs pH.

### 64 0.E+00 5.E-05 1.E-04 2.E-04 2.E-04 3.E-04

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHK (mol/L)

measured data

model results

Figure 5.9. Calculated potassium concentration vs pH.

0.E+00 5.E-04 1.E-03 2.E-03 2.E-03 3.E-03 3.E-03 4.E-03

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHMg (mol/L)

measured data

model results

Figure 5.10. Calculated magnesium concentration vs pH.



### Figures 5.11-5.14 show the calculate d trace metal concentrations for lead, cadmium, copper, and zinc. With the exception of Zn, which PCA s hows to behave anomalously, reasonable matches

between calculated and measured concentrations again suggest that the calcium-driven cation exchange could explain the observed trace metal response to CO 2 

### injection.

0.E+00 1.E-10 2.E-10 3.E-10 4.E-10 5.E-10 6.E-10 7.E-10

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHPb (mol/L)

measured data

model results

Figure 5.11. Calculated lead concentration vs pH

### 65 0.E+00 1.E-09 2.E-09 3.E-09 4.E-09 5.E-09 6.E-09

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHCd (mol/L)

measured data

model results

Figure 5.12. Calculated cadmium concentration vs pH

0.0E+00 5.0E-09 1.0E-08 1.5E-08 2.0E-08 2.5E-08 3.0E-08 3.5E-08 4.0E-08 4.5E-08 5.0E-08

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHCu (mol/L)

measured data

model results

Figure 5.13. Calculated copper concentration vs pH

0.0E+00 2.0E-08 4.0E-08 6.0E-08 8.0E-08 1.0E-07 1.2E-07 1.4E-07 1.6E-07

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHZn (mol/L)

measured data

model results

Figure 5.14. Calculated zinc concentration vs pH



### The baseline model, which considers only cati on exchange for iron, l eads to a significant

underestimation of measured iron co ncentrations (see Figure 5.15). Th is finding is consistent with the PCA, which indicates that its behavior is not identified with the component associated with ion exchange. If siderite dissolution is assumed to contribute to the increase in aqueous iron concentration, a better fit is obt ained, although the predicted iron con centrations are still lower than

### 66 those actually measured. Note that the calculated iron concentration is highly sensitive to the rate of

siderite dissolution.

Another hypothesis that has been tested to expl ain the behavior of ir on is the dissolution of Fougerite (green rust, Fe 6.7 

### Mg 5.3 

### O 34 

### H 26.7 

### C 3 

### ) (see Section 5.2.2). The dissol ution of green rust leads to a better match between computed and measur ed iron concentration, although this process is highly redox dependent, and it must be remembered that a large uncertainty exists regarding the

iron data and redox conditions at the site. Furt her determinations of the redox condition of groundwater at the ZERT site would be useful to validate the hypothesis that dissolution green rust

controls the evolution of iron.

It should be noted that the desorption of Fe +2 

### sorbed onto HFO (which is known to take place under suboxic but not necessarily strongly reducing cond itions, e.g., Hiemstra and van Riemsdijk, 2007) could also contribute to increasing Fe concentr ations in groundwater. Preliminary simulations including this process showed a larger than obser ved Fe concentration incr ease in groundwater upon pH decrease. Given the limited amount of field da ta and uncertainties regarding redox conditions at the site, further modeling work including th is process was not attempted.

0.E+00 1.E-06 2.E-06 3.E-06 4.E-06 5.E-06 6.E-06 7.E-06 8.E-06 9.E-06 1.E-05

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHFe (mol/L) measured data

base model: cation exchange only

sensitivity run: cation exchange + siderite dissolution

sensitivity run: cation exchange + green rust dissolution

Figure 5.15. Calculated iron concentration vs pH in the base model, in which only cation exchange is considered and a sensitivity run in which both cation exchange and siderite dissolution are considered



### Another effect from the dissolution of CO 2 

### in groundwater is the increase of aqueous bicarbonate (HCO 3 - 

### ) concentration (Figure 5.16). This increase m odifies the distribution of adsorbed anionic species on HFO. Table 5.16 lists the adsorbed concentr ations associated with those ions considered in the surface complexation reactions on HFO. Deso rption and adsorption cons titutes the net effects of decreasing and increasing the co ncentrations for all relevant su rface complexes. For example, the desorption/adsorption of As(V) can be represen ted by all relevant su rface complexes including HFO_sH 2 

### AsO 4 

### , HFO_wH 2 

### AsO 4 

### , HFO_sHAsO 4 - 

### and HFO_wHAsO 4 - 

### (see Table 5.13). The total adsorbed concentration (represented as s_ H 2 

### AsO 4 - 

### in Table 5.16) is the summation of the concentrations for all four comple xes. Table 5.16 s hows the initial (C i 

### ) and final (C f 

### ) adsorbed

concentrations and their difference C f 

### –C i 

### . Negative values of C f 

### –C i 

### indicate desorption, while positive values indicate adsorption. Initially, adsorb ed carbonate occupies a large portion of the sorption sites (about 55%). The increase in HCO 3 - 

### due to the dissolution of CO 2 

### leads to increased adsorption of carbonate on the HFO surface. As a resu lt, some anionic species are displaced into the aqueous phase.

### 67 Table 5.16. Initial (C i ) and final (C f ) adsorbed concentration and their difference C f –C i

Adsorbed species C i C f C f - C i

s_H + 1.7E-03 2.0E-03 3.0E-04
s_ SO 4 -2 5.2E-07 2.1E-07 -3.1E-07
s_HSeO 3 - 9.4E-07 5.2E-07 -4.2E-07
s_MoO 4 -2 1.7E-09 1.0E-09 -7.1E-10
S_HCO 3 - 2.3E-03 3.0E-03 7.2E-04
s_ H 2 AsO 4 - 8.5E-08 6.7E-08 -1.9E-08
s_ HPO 4 - 2.3E-04 2.3E-04 7.4E-07

0.0E+00 5.0E-03 1.0E-02 1.5E-02 2.0E-02 2.5E-02

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHHCO3 (mol/L) measured data

model results

Figure 5.16. Calculated bicarbonate concentration vs pH



### Figure 5.17 shows the aqueous concentr ation of arsenic. Even though sca tter in the As da ta is large, the similarities between computed and measured As concentrations suggests carbonate-initiated desorption could be a controlling process for As, as observed in other systems (e.g., Appelo et al., 2002; Frau et al., 2008). It shoul d be noted, however, that our m odel has ignored adsorption of As(III) species. This could affect model results. However, because As(III) does not sorb as much as As(V) and because significant As(III) is not expected in suboxic to oxic conditions at the site, it is unlikely that inclusion of As(III) in the system would significantly affect model results.

0.0E+00 5.0E-09 1.0E-08 1.5E-08 2.0E-08 2.5E-08 3.0E-08 3.5E-08

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHH2AsO4-2(mol/L) measured data

model results

Figure 5.17 Calculated arsenic concentration vs pH

### 68 The computed aqueous concentrations of seleni um overestimate the measured concentrations (Figure 5.18), possibly attributable either to uncertainties in th e selenite surface-complexation constant, or to the omission of adsorbed selenate or selenide species. Pr eliminary EQ3 distribution of species modeling suggests that oligomeric neutral selenide complexes of Fe and Zn may predominate in solution, although such a distribution is highly depe ndent on a correct estimate of the redox state of the groundwater, and the unlikel y assumption that homogeneous equilibrium with respect participating redox species has been attained. The concentrati on of aqueous phosphate (Figure 5.19) decreases with the decrease in pH , because anionic adsorption occurs. Although the concentrations of adsorbed HFO_sPO 4 -2 

### and HFO_wPO 4 -2 

### tend to decrease as carbonate species adsorb, the concentrations of HFO_sH 2 

### PO 4 

### and HFO_wH 2 

### PO 4 

### increase as the pH falls. This increase is greater than the decrease in HFO_sPO 4 -2 

### and HFO_wPO 4 -2 

### , thus the concentration of phosphate in the groundwater falls.

0E+00 1E-07 2E-07 3E-07 4E-07 5E-07

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHHSeO3-(mol/L) measured data

model results

Figure 5.18 Calculated selenium concentration vs pH

0.0E+00 5.0E-07 1.0E-06 1.5E-06 2.0E-06 2.5E-06

5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2

pHHPO4-2(mol/L) measured data

model results

Figure 5.19 Calculated phosphate concentration vs pH



### In summary, the concentration in crease of major and trace elemen ts in the groundwater at the MSUZERT field test has been inte rpreted with a batch geochemical model in which hypothesized reactive processes were tested, including the dissol ution of calcite and Fe minerals, ion exchange, and sorption/desorption. These processes were pr esumed to take place on the basis of observed groundwater quality trends (Sect ion 3) as well as statistica l and thermodynamic evaluations

### 69 (Section 5.2). Only data before the first rainfall (a fter the start of the test) were used for comparing model results with observed trends, to avoid comp lexities resulting from the mixing of infiltrating rainwater with groundwater. Reasonable agreemen t between measured data and model results suggest that (1) calcite dissoluti on could be the primary process buffering pH, (2) the increase in the concentrations of major cations and trace metals except Fe could be explained by Ca +2 

### -driven exchange reactions, (3) the release of anions from adsorption sites due to competing adsorption of bicarbonate could explain the concen tration trends of most anions and (4) the dissolu tion of reactive Fe minerals could explain the incr ease in total Fe concentration.



## 6. SUMMARY AND CONCLUDING REMARKS

The primary purpose of the MSU ZERT CO 2 

### release test conducted in summer 2008 was to allow

controlled studies of near-surface CO 2 

### transport and detection technolog ies. However, this test also provided a unique opportunity to evaluate the effects of CO 2 

### release on the chemistry of groundwater at this site, which is the focus of this study.

When CO 2 

### dissolves in water, carbonic acid is produced, decreasing the pH and increasing the

dissolved carbonate content of the water reacting with CO 2 

### . These processes and their effects were monitored at the ZERT site, through geochemical investigations conducted prior to, during, and after the 2008 injection te st. During this test, CO 2 

### was injected for 30 days into a shallow horizontal pipe installed at a depth near 1.8 m below ground surface, and about 0.5 m to 1 m below the water table. Groundwater chemistry was monitored from five locations, one upgradient and four downgradient of the injection pipe. Over 60 gro undwater samples were collected. Most of these samples were obtained from shallow wells screened at depths close the soil-g roundwater interface.

The samples showed a fast response to the CO 2 

### release, as early as a day after start of injection in nearby wells. The response was marked primarily by a pH decrease from ambient values ~ 7 to minimum values ~5.8 at wells closest to the inje ction well (within 1–2 m from it), and a significant increase in alkalinity and elec trical conductivity. The concentrations of many dissolved species, including trace metals, also increased. Volatile or ganics were also detect ed in groundwater as the

result of CO 2 

### injection, but were found to be linked to the CO 2 

### source and not to CO 2 

### -sediment

interactions. Concentrations of constituents listed on the EPA prio rity pollutant risks (US EPA, 2004), however, did not rise above their established maximum contam inant levels (MCL). Several samples were also collected from deeper wells screened at levels below the level of the injection pipe. These samples, however, ha ve not been affected by the CO 2 

### release. Several rainfall events occurred during the test. Because of the thin vadose zone, water levels and water chemistry were quick to respond to rainfall (within a day or more). In particular, th e concentrations of Fe and Mn in groundwater were observed to drop at the onse t of rainfall, presumably as the result of groundwater mixing with more oxi c recharge water. For these reasons, the interpretation of geochemical system response to CO 2 

### injection concentrated on the first week of CO 2 

### injection, during which time the test was not affected by rainfall.

Standard geochemical evaluations , statistical analyses, and geoc hemical modeling suggest that:

 CO 2 

### dissolution in the groundwater causes calcite to dissolve. Calcula tions of the calcite saturation index indicate increasing undersaturation with increasing CO 2 

### uptake, with slow calcite dissolution and incomplete resaturati on during the first ten days of the test.

### 70  Observed increases in the concentration of dissolved trace metals result likely from Ca +2



### - driven ion exchange with clays (smectit es) and sorption/desorption reactions likely involving Fe (hydr)oxides.

 Bicarbonate from CO 2 

### dissolution appears to compete for sorption with anionic species such

as HAsO 4 -2 

### , potentially increasing dissolved As levels in groundwater.

In particular, principal component analyses (PCA) for the first 10 days of the test, using the collected groundwater qual ity data show that:

 A strong correlation exists between alkali metals (Li, Na, K), al kali earths (Mg, Ca, Sr, Ba), B, Al, Cr, SiO 2 

### (aq), Cd and Pb, and an inverse relati onship to pH, suggesting ion exchange as a release mechanism for these cationic species.

 F, As, Se, and Mo are correlated, suggesting po ssible adsorption of these anionic species on secondary iron (hydr)oxides or anionic clays.

 PO 4 

### is negatively correlated with Mn, Fe, and Co, suggesting some precipitation/dissolution mechanism may be operative.

 Cl and SO 4 -2 

### are correlated and suggest perturbati on by earlier rainfall events, prior to monitoring.

Thermodynamic analyses indicate that:

 The initial groundwater is slightly undersaturated with respect to calcite, opal-CT and barite, suggesting dilution of the groundwater by rainwater prior to the test.

 The calculated redox potential is at pe ≈ +9 assuming fougerite-CO 3 

### and Fe(OH) 3 

### (ferrihydrite) as the cont rolling phases, and pe ≈ 2 if assuming goethite. There are indications that the system redox state may not be in equilibrium or spatially homogeneous.

Geochemical simulations were run to test whether ion exchange and sorpti on processes, together

with dissolution of car bonates, could explain the water chemis try trends observed during the test. Results show reasonable agreement with field da ta, suggesting these proc esses were triggered by

the CO 2 

### release. Because of the lack of site-specific data for key model parameters such as the cation exchange capacity, smectite and iron oxyhydr oxide content, sorption data, mineral surface areas, and redox state, the simulatio ns relied on assumed and/or typi cal but not site-s pecific values for these critical parameters. In addition, no unique solution may exist for a system as complex as at the ZERT site, for which multiple combinations of input parameters and modeled processes could yield similar results. Consequently, the modeli ng effort should be regarded as providing more qualitative than quantitative results.

It is evident that further detail ed characterization of sediments and groundwater samples, including redox-sensitive sampling and analysis, would be re quired to test some of the hypotheses made in this study, and to provide a full understanding of metal behavior and speciation in this complex system. Such detailed analysis, however, is not within the scop e of this study. Ongoing investigations, including the analyses of metals in sequential extractions of sediment samples, may prove quite useful to determine th e association of metals with so lid phases in the sediment. Such associations are critical as they are expected to dictate the magnitude and type (dissolution, desorption, exchange) of groundw ater quality response to CO 2 

### injection.

### 71 It should be noted that this study was conducted in a shallow wa ter-table aquifer under conditions that are not representative of deeper USDW. However, the collect ed geochemical data during this study were useful to evaluate key processes (dissolution, exchange, and sorption/desorption reactions) that would be e xpected to take place in any aquifer impacted by CO 2 

### . It is anticipated that higher pressures in deep USDW would result in increased CO 2 

### dissolution in groundwater, relative to observations at the ZERT site, and thus coul d result in a more pronounced pH decrease and possibly a more significant release of contamin ants. Such stronger res ponse has indeed been observed during CO 2 

### injection tests in reservoirs at much greater depths, such at the Frio test site (Kharaka et al., 2009b). Furthermor e, the controlled-release period of about four weeks may have been too short to test the importance of slower release mechanisms for trace metals, such as the

dissolution minerals containi ng trace amounts of metals. We recommend conducting further controlled-release experiments to improve our understanding of po ssible groundwater contamination from CO 2 

### intrusion, ideally in deeper aquifers more representative of the majority of our groundwater resources, and with l onger injection period to monitor contaminant release from shortterm and long-term geochemical processes.

The possible impact of increased CO 2 

### dissolution in groundwaters at gr eater depth could be partly mitigated when impacted waters migrate upward from deeper formations, releasing CO 2 

### as the hydrostatic pressure decreases. It has long been known, for exampl e, that the exsolution of CO 2 

### from naturally CO 2 

### -rich and metal-rich thermal waters upon ascension to shallow depths is accompanied by a pH increase that causes the precipitation of calcite and other metal-bearing phases (e.g., Drummond and Ohmoto, 1985; S pycher and Reed, 1989). The effect of CO 2 

### exsolution along smaller temperature gradients but higher pressure gradient s in the case of CCS (relative to thermal waters) has not been evaluate d, but it would be anticipated to at least partly reverse the trends predicted in this study. Therefore, the full assessment of CCS impacts on the quality of a USDW may require assessing geochemical processes not only at the source of a CO 2 

### leak, but also along the path of impacted (carbon ated) waters as they migrate towards potential

receptors.

### 72 

## 7. ACKNOWLEDGMENTS and Disclaimer

This work was conducted as part of the ZERT pr oject, supported by the A ssistant Secretary for Fossil Energy, Office of Sequestration, Hydrogen, and Clean Coal Fuels, through the National Energy Technology Laboratory, U.S. Departme nt of Energy, under Contract No. DE-AC02-

05CH11231. This research was conducted at the ZERT field facility directed by Lee Spangler and managed by Laura Dobeck, MSU, Bozeman, MT. http://www.montana.edu/zert. We thank the entire ZERT team and participating organizations for creating a supportive and exciting research environment, and gratefully acknowledge. This research was funded primarily by the Electric Power Research Institute, EPRI; additional funds were also obtained from EPA, DOE, LBNL, and

USGS.

The USGS contributions to this re port have not undergone formal internal review and approval for publication by the USGS; therefore, these contributions are subject to change and alternative interpretation.



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Geochimica et Cosmochimica Acta 70, 4600–4616.79APPENDIX ATable A1Chemical composition of water samples from ZERT monitoring wellsMW1Major Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1037/7/0817:006.8104329.25.524950.310.120.001<0.010.103.70.0420.790.175.8273721.00.0090.090.30.41.50.0030.3<0.02<24.94.0MW2Major Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1027/7/0815:306.8103967.56.224850.270.130.002<0.010.096.50.0510.120.208.3274220.80.0120.07<0.2<0.41.40.0030.4<0.02<24.33.41AMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1057/8/089:155886.983867.85.623840.270.110.001<0.010.104.70.0400.550.177.5263920.90.0120.08<0.2<0.41.20.0030.40.02<24.02.81BMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1127/10/0815:156506.5134388.46.026940.290.110.003<0.010.144.90.0490.330.167.1344631.20.0180.280.3151.70.0060.40.0654.2408ZERT-1177/12/089:158475.795519.46.8341240.390.150.003<0.010.095.00.046<0.010.147.3374941.20.0220.380.3301.90.0070.40.0654.03.608ZERT-1277/15/0810:0010155.997249.87.6431550.490.190.005<0.0130.115.20.0460.310.177.7415451.30.0210.430.4512.10.0081.00.0864.63.208ZERT-1317/17/0810:3010415.997359.77.7441590.510.200.0060.0140.115.30.0480.280.187.9354840.90.0230.400.4132.00.0070.60.0744.83.508ZERT-1357/19/0810:1511246.1108219.98.1461740.550.220.0170.0170.125.50.0550.260.198.1375141.20.0170.210.552.10.0050.30.0535.02.608ZERT-1397/21/089:5011816.02110009.47.7441730.530.210.021<0.0250.155.60.0460.230.188.4365051.10.0160.33<0.25182.20.0070.40.06<35.22.608ZERT-1447/23/0811:459086.0158618.77.7401580.490.190.012<0.0130.125.60.0520.230.178.4344931.00.0130.200.422.20.0050.40.0534.84.008ZERT-1477/26/089:306675.9135657.66.2261090.330.120.005<0.010.135.60.0460.230.177.7284121.00.0110.110.2<0.41.90.0040.50.05<24.33.408ZERT-1517/29/089:006465.7154808.26.125990.310.110.0080.0140.155.60.0470.260.197.6274120.90.0110.110.2<0.41.50.0040.30.03<24.22.908ZERT-1608/4/0811:159615.9156939.37.5381470.460.180.0060.0180.135.60.0500.380.198.0334740.90.0210.330.3<0.51.90.0060.40.0643.92.608ZERT-1638/7/089:009845.9157029.27.7381490.470.190.0060.040.105.60.0560.480.208.1385231.10.0170.150.30.91.80.0050.30.0334.02.508ZERT-1678/8/089:009545.8136689.07.7361470.460.180.0060.0300.115.60.0590.500.208.0385241.00.0190.200.3161.80.0060.30.04<34.02.208ZERT-1718/11/089:407056.0135297.96.8291180.360.140.004<0.010.115.80.0620.510.188.0344830.80.0160.190.3121.90.0060.30.04<23.72.108ZERT-1768/13/0811:206256.4264027.46.121910.280.110.002<0.0060.125.90.0620.680.207.8304531.00.0160.190.271.70.0050.50.0234.11.62AMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1077/8/0811:455856.783877.55.324840.260.110.001<0.010.124.80.0420.600.187.4324531.00.0180.230.2121.60.0050.40.05<44.13.1802BMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1097/9/089:006517.0124349.15.428920.300.100.028<0.010.175.40.0990.260.107.2324531.30.0170.290.4122.50.0070.70.0654.33.808ZERT-1187/12/0810:009526.496649.77.1411420.450.190.190.080.145.30.0480.120.057.4405351.00.0220.451.2542.30.0080.50.0863.89.008ZERT-1287/15/0811:0012705.9981610.18.3542040.630.270.130.390.075.50.0490.200.027.7445771.30.0240.611.2722.60.0100.50.1074.53.208ZERT-1327/17/0811:1511935.999249.58.0491910.570.260.140.530.065.60.0490.20<0.0157.8375160.40.0200.431.2212.20.0070.50.06<34.32.308ZERT-1367/19/0811:1513426.01011159.97.4552230.690.230.011<0.0250.055.50.0560.250.248.4385260.90.0170.220.572.40.007<0.50.0534.12.808ZERT-1407/21/0810:1013356.01611279.66.9512270.670.210.008<0.0250.095.50.0470.300.268.4375181.10.0210.40<0.25262.40.0080.40.0734.43.708ZERT-1467/23/0813:1514246.015114710.27.4552390.730.240.014<0.0250.135.60.0510.350.268.6395481.60.0180.220.612.40.0070.70.0644.43.508ZERT-1507/26/0811:3513395.91310549.67.1472410.680.220.015<0.025<0.055.60.0470.410.028.5304571.5<0.0150.19<0.5<1.02.00.007<0.75<0.05<54.04.008ZERT-1547/29/0813:4512355.8219678.87.6402160.580.250.130.780.065.70.0550.46<0.0158.0294451.2<0.0150.181.2<11.50.0060.2<0.05<54.04.408ZERT-1618/4/0812:1511955.8199168.57.7362190.520.270.191.10.105.80.0520.64<0.0157.83853100.40.0200.321.3411.90.0060.4<0.05<34.03.508ZERT-1658/7/0810:4512015.7228848.87.7352120.500.260.211.190.076.10.0620.77<0.0158.0385380.50.0160.231.572.00.0050.30.0334.14.208ZERT-1698/8/0811:257326.0215117.15.9201250.290.150.130.870.186.30.0651.0<0.0158.0314640.40.0180.190.8121.60.0050.60.03<24.12.808ZERT-1728/11/0810:206155.8144517.25.4171060.250.120.0900.350.246.50.0701.10.0237.9294530.50.0140.140.6131.70.0040.50.02<24.42.508ZERT-1778/13/0812:306066.4253897.85.216940.230.100.0520.150.276.90.0731.40.067.8294620.70.0170.150.571.60.0050.70.0434.35.93BMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1748/11/0811:3014306.216106911.39.1632240.750.310.21<0.025<0.056.30.0630.720.079.9436061.40.0240.610.5472.80.0100.80.0965.44.64AMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1067/8/0810:006096.983988.15.625860.280.100.31<0.010.115.10.0440.190.137.4263920.80.0130.100.6<0.41.10.0040.30.02<24.03.24BMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1047/7/0817:406446.9154319.75.627900.300.100.110.050.175.30.0460.100.097.3334631.00.0170.260.8102.40.0061.30.08<44.52.408ZERT-1197/12/0811:006247.094108.25.825870.280.121.020.050.145.40.044<0.010.037.6314430.90.0190.321.8151.50.0060.70.0754.14.208ZERT-1267/15/089:157066.394688.76.029990.330.141.20.080.135.40.045<0.010.027.4344730.90.0190.372.0281.60.0060.60.0654.43.308ZERT-1307/17/089:307206.394759.16.1301010.330.151.20.080.115.40.045<0.010.0247.5324530.60.0210.302.071.60.0050.50.0534.03.108ZERT-1347/19/089:008486.51053310.36.7351230.410.160.510.0300.135.40.0500.070.057.7344830.60.0190.181.232.00.0040.60.0424.62.208ZERT-1417/21/0811:308826.61558910.36.6371300.420.160.37<0.0130.135.50.0460.100.057.9334740.60.0220.350.7132.10.0060.50.0534.22.508ZERT-1437/23/0811:009976.21670811.67.0421430.490.170.280.0340.125.40.0540.140.058.2375130.60.0180.190.842.30.0060.50.0424.84.408ZERT-1487/26/0810:1510856.41474212.17.6471630.550.190.340.0200.115.50.0490.200.048.5304530.70.0140.180.8<0.52.00.0050.40.04<35.02.808ZERT-1527/29/089:509416.41464110.77.4411410.480.191.000.110.115.60.0540.130.028.2324620.80.0130.191.6<0.51.40.0050.40.16<34.53.008ZERT-1598/4/0810:008076.1145289.17.0331150.380.181.440.210.096.10.0600.13<0.0158.0354940.70.0220.442.5201.70.0060.60.0643.82.908ZERT-1668/7/0812:157866.2165429.17.0331110.380.171.40.140.096.30.0660.21<0.0158.0365130.60.0150.232.551.60.0050.50.0334.06.108ZERT-1688/8/089:457946.3155469.17.0331130.380.181.410.180.096.40.0680.24<0.0158.0355030.50.0190.272.5101.60.0060.50.04<24.02.308ZERT-1738/11/0811:057026.3165819.57.0351220.410.170.650.0680.116.70.0740.660.0168.3355130.50.0200.261.2111.80.0070.50.04<24.12.208ZERT-1788/13/0813:507276.5244848.56.6291000.350.140.770.0890.137.00.0760.870.0208.2334930.70.0160.221.451.70.0060.60.0344.12.8815BMajor Solutes (mg/L)Trace Solutes (μg/L)SAMPLEDATEtimeECpHTHCO3NaKMgCaSrBaMnFeFClBrNO3PO4SO4SiO2TDSAlAsBCdCoCrCuLiMoPbSeUZn08ZERT-1167/11/0814:4512576.2999412.76.5661900.670.250.0460.190.155.30.051<0.01<0.0158.146608<20.0260.411.0462.70.0111.00.07<104.73.908ZERT-1247/14/0813:1512876.09117411.76.8682010.710.270.140.630.145.40.0540.018<0.0158.550648<20.0280.592.5872.80.0130.80.10<104.95.908ZERT-1297/15/0812:3012646.01098211.36.8641810.630.260.170.930.065.40.0530.049<0.0158.347617<20.0240.563.1762.40.0120.70.08<104.73.608ZERT-1337/17/0812:0013215.99106011.26.8701990.710.270.190.720.095.50.0520.028<0.0158.7415570.60.0250.582.4202.30.0100.70.0945.24.208ZERT-1377/19/0813:0014346.010115210.96.8702270.830.290.150.060.065.50.0560.050.079.1425750.60.0190.311.082.90.0080.90.0735.73.908ZERT-1427/21/0811:5014156.023108110.56.6662280.800.280.10<0.0250.055.60.0510.140.079.5395570.60.0260.550.8273.00.0101.00.0745.95.408ZERT-1457/23/0812:4014955.915113611.46.9682330.830.290.110.0260.105.60.0550.320.0910.2425851.10.0160.321.193.50.0090.70.0746.26.608ZERT-1497/26/0811:0012856.01510239.96.4552130.740.250.058<0.025<0.055.60.0480.550.089.4294441.1<0.0150.21<0.5<1.02.90.007<0.75<0.05<54.64.808ZERT-1537/29/0811:1511116.0178819.26.0461850.620.220.0830.100.165.60.0520.340.038.5294430.8<0.0120.200.8<0.81.90.007<0.6<0.04<44.14.408ZERT-1628/4/0813:0011135.9178439.56.5481870.630.240.140.310.135.90.0560.50<0.0158.5385350.50.0210.541.4402.00.0070.70.0654.22.908ZERT-1648/7/0810:006.07619.46.5441600.550.210.200.850.156.20.0680.68<0.0158.5395540.60.0170.272.181.80.0060.50.0334.23.008ZERT-1708/8/0812:3011306.1208329.56.8471760.620.240.230.770.156.40.0690.54<0.0158.737535<0.50.0210.412.3121.80.0090.70.05<34.24.308ZERT-1758/11/0812:4510376.1167278.46.0381670.550.210.0720.0550.116.70.0751.50.0199.234524<0.50.0170.350.7142.50.0080.60.04<34.12.808ZERT-1798/13/0815:047596.2205067.05.2251140.380.140.0330.0710.167.00.0821.90.0188.9294630.80.0130.220.512.20.0060.80.0343.52.9

### 82 

## APPENDIX B

Table B1 Well log for well 6

Core (ft) (cm) % re covery Depth (cm) Notes BTEX Description

0-5cm grass and air
dark brown, clay rich, plastic soil with roots througout. Slightly fizzes with HCl
11-15cm large rock m ost likely broken by geoprobe

30-39cm plastic clay grading from dk brown to lt brown indicating a
gradation into caliche

39-57cm 46-48.5cm 10YR 5/4. Caliche zone. Highly reactive with HCl, very fine
grained and plastic with some roots. Few larger, rounded
grains between 1-3cm.

57-64cm Sandy gravels (with som e caliche smearing on kevlar sleeve).

0-15cm 11-15cm Slough? with a large grain cored through by the geoprobe at

11-15cm.

15-76cm 30cm 54-58cm 10YR 3/3 dark brown. Coarse sandy gravel, possibly higher
percentage of gravels than other cores. Grains as large as 4cm. Gravels appear to have been rounded then broken,
possible geoprobe damage? Saturated below 30cm. Lightly
fizzes with acid throug

8-10' 61.0 67% 0-41cm 33-37cm 10YR 3/3 dark brown. Coarse sandy gravel, few fines.
Grains up to 4cm, possibly damaged by geoprobe. Some plastic from a broken core catcher is present. Few orange
clay streaks near 37cm. 3.5-8' Well 6 Log Date 4/29/09 - 4/30/09

0-3.5' 5-30cm

107

137 60%

55%

### 83 

## APPENDIX C

Estimation of the Solubility Product of Fougerite

Over the past twenty years, considerable st udy has been devoted to the characterization of

metastable minerals that form in soil zones at or near the water table, an d where the redox state is affected by decomposition of organic matter derived from vegetative cover. In particular, much attention has been paid to the formation of so called ”green rusts”, which consist of stacked positively charged tri-octahedral sheets of ferrous hydroxide in which some of the Fe(II) has been oxidized to Fe(III). The octahedral sheets are sepa rated by interlayers containing water and chargebalancing anions (see Genin et al ., 2006, and references cited therein) . At the time of their initial characterization (Bernal et al ., 1959; Dasgupta and Mackay, 1959) , it was found that green rusts could be characterized by two distinct structur es; (1) a rhombohedral fo rm designated as GR-I, which contained simple charge-balancing anions such as OH - 

### or Cl - 

### , or planar oxy-anionic charge balancing species such as HCO 3 - 

### and/or CO 3 -2 

### , (2) a hexagonal form containing large oxy-anionic species such as SO 4 -2 

### and SeO 4 -2



### .

Earlier work concluded that green rusts occupy a stability field in Eh-pH space defined by neutral to weakly acid pH and moderately re ducing (anoxic) conditions (Refait and Genin, 1993; Drissi et al., 1995; Genin et al., 1996, 1998; Ruby et al., 2006). Other minerals that might also be associated with green rusts, depending on the concentration of Fe(T) and SO 4 -2 

### in the system, and fluctuations in oxidations state include hydroxy-magnetite or ma gnetite (Olowe and Genin, 1991), schwertmannite and ferrihydrite (Majzlan et al., 2004), these latter phases principally occupying stability fields at higher Eh, and where the phase boundary between schw ertmannite and ferrihydrite falls within a pH range close to 7.0. Ruby et al. (2006) also noted that th e initial oxidation of green rust causes a characteristic color change from blue-green to orange-yellow while pres erving the original trioctahedral structure. It is notewor thy that some of the recovered fi ne material from cores taken at the ZERT site also contained unspecified iron oxi des characterized as “powdery orange blebs” (Thordsen, 2009), lending credence to the belief that green rust may actually form in shallow groundwater at the ZERT site. Green rust decom poses exceedingly rapidly upon exposure to air, and therefore it would not be expected to surv ive coring and recovery at the ZERT site unless extreme precautions had been taken. Elsewhere (S chwertmann and Fechter, 1994) has observed that green rust eventually transforms to lepidocroci te under oxidizing conditions . However, whether or not green rust transforms to fe rrihydrite as an intermediate st ep before recrystallization to lepidocrocite is not clear from the literatur e, although the following thermodynamic analysis suggests that this could be the case at the ZERT site.

Characterization of green rusts in a gleysol horizon at Fougeres in Normandy, France (Trolard et al., 1996; Abdelmoula et al., 1998) lead to its r ecognition as a mineral species; fougerite, and

furthermore revealed that the tri-oc tahedral sheets are in fact a mixtur e of Fe(II), Fe(III) and Mg(II). Bourrie et al. (2004) created a solid solution mode l of the hydroxide form of fougerite in which Fe(II)-Mg(II) substitution was treated as an id eal solution, whereas Fe(II)-Fe(III) substitution was treated as a regular solution with the constraint that adjacent Fe(III) sites are prohibited. Information on the thermodynamic properties of the end member binary solid solutions in the three-component system was obtained from prior solu bility and calorimetric studies re ported in the literature. This model was subsequently expanded by Feder et al. (2005) to cons ider anion occupancies by (HCO 3 -



### + CO 3 -2 

### ), Cl - 

### and SO 4 -2 

### in place of OH -

- 

### in the interlayer positions, without, however, taking into

### 84 account interlayer ion exchange. Re cently, Christiansen et al. (2009) found that at least some green rusts also contain interlayer cati ons in addition to anions. Therefor e, a complete characterization of green rust thermodynamic properties must await further study. Feder et al. (2005) also examined the possibility that these forms of fougerite might achieve thermodynamic equilibrium in the saturated soil zone at Fougeres. They found that saturation was achieved with respect to certain carbonate and sulfate fougerite compositions, but that chloride forms would always be undersaturated, whereas the hydroxide forms would always be grossly supersaturated.

A preliminary comparison of groundwater compositions at Fougeres (Feder et al., 2005) with those at the ZERT site indicated that although temperature and pH were comparable, Mg +2 

### , Fe(T), Cl - 

### and

SO 4 -2 

### concentrations were lower, and the alkalinity (as HCO 3 - 

### ) was higher. These observations suggest that a carbonate form of fougerite might be st able at the ZERT site.

The solubility products of all forms of fougerite were calculated assuming the stoichiometry to be ([2/9Fe II 

### (OH) 2 

### .1/3Fe III 

### (OH) 2 

### .4/9Mg(OH) 2 ]+1/3 

### [1/3A] -1/3 

### ), where A = OH-, Cl-, (2/3HCO 3 - 

### .1/6CO 3 -2 

### ),

or 1/2SO 4 -2 

### . The solubility products were determined by direct comparison with the reported field saturation states at Fougeres, rather than by calcu lating the solubility products from the reported thermodynamic data given by Bourrie et al. (2003) and Feder et al. (2005), as the reference states

used in their derivations differe d from those we currently use in EQ3/6. Incorporation of the solubility products of the fougerite forms in th e EQ3/6 thermodynamic database and testing using the baseline groundwater composition before CO 2 

### injection indicated that the carbonate form of fougerite could be stable at the ZERT site, depending on the oxidation state of the system.

### 85 

## APPENDIX D

The Saturation State of Redox-Sensitive Ch romium, Copper, and Selenium Minerals as Function of pe.

In this appendix, we pick up on Section 5.2.2, whic h considers potential solid phase redox buffers that might influence the redox state, or be infl uenced by the redox state, in the coexisting groundwater in a shallow unconfined aquifer such as th at represented by the ZERT site. We start with a

discussion of the stabilities of a suite of chromium oxide minera ls, most of which are spinels.

Figure D.1. Normalized saturation indices of chromite, magnesiochromite, and the chromium spinels of Cd, Co, Cu and Zn, Eskolaite and Cr-ferrihydrite, plotted as a function of pe.



### The phase behavior of these chromium oxide minera ls as a function of pe is displayed in Figure

D.1. The saturation indices of chromium spinels, MCr 2 

### O 4 

### , where M = Mg(II), Fe (II), Cu(II), Co(II) and Zn(II), and eskolaite (Cr 2 

### O 3 

### ) varies systematically with pe, a nd at intermediate pe values, i.e., pe +2 - +8, are, with the exception of the Zn spinel, closely correlated. The phase boundary between Cr-ferrihydrite (Fe 4 

### (CrO 4 

### )(OH) 10 

### ), and chromite (FeCr 2 

### O 4 

### ) occurs at pe ≈ +8.7 at SI ≈ +2, where the pe value is similar to that between fougerite-CO 3 

### and Fe(OH) 3 

### . The significance of this correlation is not clear. It is prob able that Cr(III) substi tutes as a minor component in these phases at the ZERT site. However, this substitution is constrained at higher redox potentials by the oxidation of Cr(III) to Cr(VI) in solution with concom itant destabilization of chromium spinels, as is evident in the figure. The close correlation between the saturation indices of the chromium spinels between pe +2 - +8, invites speculation that the concentrations of Mg (II), Fe(II), Cu(II), and Co(II) in solution might be determined by solid so lution equilibrium with respect to some common iron, or Fe(II)/Mg(II) oxide mineral in the system at th e ZERT site. However, the principal component analysis is equivocal on this point. Further study of th e behavior of Cr in shallow

### 86 groundwater systems is certainly merited, and th e thermodynamic properties of the participating

minerals should be verified.

Figure D.2. Normalized saturation indices of copper, bornite, covellite, chalcocite, cuprite and tenorite, plotted as a function of pe.



### The normalized saturation indices of various copper minerals as a function of pe are plotted in Figure D.2. The redox behavior of copper is of inte rest because phase relations between tenorite (CuO) and cuprite (Cu 2 

### O) marks the redox transition between the Cu(II) and Cu(I) oxidation states, which occurs under moderately reducing conditions. The concentration of Cu in solution at the

ZERT site is insufficient to cause saturation with respect to teno rite. However, both native copper and cuprite would achieve saturation at pe ≈ +1.5, whereas the chalcocite and cuprite phase boundary is intersected at pe ≈ -0.5.

It is interesting to note that at pe ≤ +2, cuprite remains close to saturation, i.e., SI ≈ 0. A possible explanation is that the redox pair Cu +2 

### /Cu + 

### forms one of the cluster of redox pairs in local partial equilibrium with respect to organic matter, as observed by Washington et al. (2004). However, the average crustal abundance of Cu is such that this redox pair is un likely to exert any redox control over the system. Similar cont rol by organic matter on the SO 4 -2 

### /HS - 

### redox pair might be responsible for saturation with respect to chalcocite.

### 87 Figure D.3. Normalized saturation indices of selenium, moly bdenum selenide, cadmoselite and clausthalite, plotted as a function of pe.



### Figure D.3 illustrates the behavior of the saturation indice s of selenium minerals with variation in pe. Saturation is achieved with re spect to native selenium at pe ≈ +2.7, and with respect to

cadmoselite and clausthalite at about +2.4. MoSe 2 

### , a phase that has not been identified in nature, would saturate at pe ≈ 0. Given that the participating trace elements Cd, Mo, Pb and Se are all present in exceedingly low concentrations, none of the saturating phases are likely to exert a significant redox buffering capacity on the system. Howe ver, as noted with re spect to the potential saturation of cuprite and chalcocite in the copper system, the concentr ations of Cd and Pb could be controlled by saturation with respect to their seleni des or sulfides as a re sult of reduction of Se(IV) to Se(0) and Se(-II), and S(VI) to S(-II) by organic matter.

Both Cu and Se are trace element with very lo w natural crustal abundan ce, and it is therefore unlikely that either play a significa nt role in controlling the oxidat ion state of the system. However, the impact of organic matter in exercising local partial equilibrium control over some redox pairs cannot be discounted, especially in relation to Se, S and possibly U. This would imply that, although the redox state of MSU-ZERT groundwaters are controlled primarily by iron minerals and their interaction with infiltrating nitrate, organic ma tter could also be playing a role with respect to some species. In other words, the MSU-ZERT groundw aters are behaving in a manner similar to shallow groundwaters investigated by Washington et al. (2004) with formation of two redox cluster in local partial equilibrium, one dominated by ni trate and the other dominated by organic matter.

An intriguing possibility, which merits further stu dy, is that a redox potential gradient exists at the site, with more reducing conditions at greater dept h, where transient ingress of infiltrating oxidizing groundwater is less likely to have an impact. Th e more reducing condition s induce local saturation with respect to Cu(I) mineral such as cuprite or chalcocite, and control th e activity of copper in

### 88 solution, which is established in the overlying oxidizing horizons through diffusive transport. Similar behavior is observed with respect to the selenides cadmoselite (CdSe) and clausthalite (PbSe) with respect to seleni um minerals. The foregoing statement is highly speculative and requires further study before it can be advanced as a serious hypothesis.