# Drinking Water Treatment Considerations: Water Quality, Carbon Dioxide Concentration, and Geologic Sequestration ProjectsOffice of Water (4606 M)
EPA 816

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‐ 01 4 November 2010 http://water.epa.gov/drink/Page 1 

### Drinking Water Treatment Considerations: Water Quality, Carbon Dioxide Concentration, and G eologic S equestration Projects

1. 

### Introduction

The Federal Requirements under the Underground Injection Control (U I C) Program for Carbon
Dioxide Geologic Sequestration (GS) Well s (referred to in this document as the Class VI rule) are designed to minimize the risk of injected carbon dioxide contaminating an underground s ource of drinking water (USDW) . The Class VI ru le specifies multiple barriers to prevent endangerment of
USDWs from carbon dioxide injection for GS , including geologic site characterization; area of review (AoR) and corrective action; well construction and operation; testing and monitoring; and well pl ugging, post

- injection site care, and site closure.

Although the tailored requirements of the Class VI rule are designed to prevent the movement of injected carbon dioxide i nto USDW s , there is a remote potential for the movement of carbon
dioxide or for mation fluids into a USDW that supplies a drinking water system . Although carbon dioxide is not in itself toxic, it can change the water chemistry of a subsurface formation (including a
USDW ) which can

, in turn , a ffect drinking water treatment processes

. An understanding of the impacts of reactions between the carbon dioxide and water, particularly as they relate to water treatment technologies

, is necessary to inform the development of emergency and remedial response plans and allow expeditious response s to adverse events. To support owners or operators of Class VI wells and UIC program Directors in managing GS projects and preventing endangerment of USDWs, this paper presents the following information:

**•** Chapter 2

, Impacts of Carbon Dioxide on Water Chemi stry , p rovides background on the chemistry of carbon dioxide

- water interactions and discuss es the likely effects that mixing
with carbon dioxide would have on water quality within a USDW

.

**•** Chapter 3

, Impact of USDW Water Quality Changes on Drinking Water Treatment Technologies

, discuss es how the changes in water quality might affect technologies

common ly used for drinking water treatment.

**•** Chapter 4

, Potential Ground W ater Treatment Scenarios
, present s several ground water quality scenarios which are like ly to be representative of situations that might occur with a
leak of carbon dioxide into a USDW , outline s typical drinking water treatment trains which would be in place for each scenario, and discuss process changes that may be required to provide water which meets Safe Drinking Water Act (SDWA) requirements.

**•** Chapter 5, Remediation Options to Address Increased C arbon Dioxide Concentrations in
Ground W ater , discusses how a USDW of degraded water quality due to increased
concentrations of carbon dioxide m ay be returned to compliance with drinking water
standards.

**•** Chapter 6 presents a summary of the paper and conclusions

.Page 2 Given the wide range of potential geochemical scenarios, this paper cannot describe every possible configuration of USDW and rock format ion compositions, water qualities, and drinking water treatment and infrastructure configurations . Instead, this paper addresses ground water quality
conditions and drinking water treatment technologies that are typically encountered in the U.S. It is

no t intended to be an exhaustive description of all potential situations or responses ; rather

, it serves as a guideline to in for m GS project planning.

2. 

### Impacts of Carbon Dioxide on Water Chemistry

T o support an understand ing of geochemical changes that may potentially occur in subsurface
formation s and what effect s, if any, the se changes will have on potability and treatability of the

water , this section provides an overview of the impacts of carbon dioxide contact with formation
fluids and rock matrices an d subsurface fluid movement.

Very few field studies have been conducted to measure the effects carbon dioxide may have on a USDW, with the exception of a few field studies by the Department of Energy ’s (DOE’s) Regional Carbon Sequestration Partnerships throughout the United States (Litynski, 2008). Only one of these studies has released data on changes in water quality caused by injection of carbon dioxide (Kharaka et al., 2006). This study measured changes in the formation water in the injection zone

. While not directly applicable to the situation of carbon dioxide in a USDW, the results show the type of reactions that might occur, given the similarities with formation materials. Of the studies conducted, several used modeling or laboratory experiments to determine the effect of carbon dioxide exposure on water quality. One major study by the Lawrence Berkeley National Laboratory (LBNL) modeled the effects that carbon dioxide movement would likely have on a USDW (Birkholzer et al., 2008). The study analyzed 38,000 ground water samples from the United States Geological Survey’s National Water Quality Information System to estimate typical water quality and mineral equilibrium conditions of USDWs.

The following sections summarize the results of these studies. Section 2.1 present s background
information about how carbon dioxide reacts with water and resultant changes in pH/alkalinity.
Section 2.2 describes mineral reactions and impacts on metals chemistry that can occur in pH

altered environments. Sect ion 2.3 describes salinity changes associated with fluid movement in the
subsurface.

2.1 Geochemistry of Carbon Dioxide

- Water Interactions

Gaseous or supercritical carbon dioxide coming into contact with water will come into equilibrium
with the water

. So me of the carbon dioxide will dissolv e to form carbonic acid, and through further
dissociation reactions, bicarbonate and carbonate

, as illustrated in equations 1 through 3.Page 3 CO 2 (g) + H 2 O H 2 CO 3 (aq) ( Equation 1) (carbonic acid)

H 2 CO 3 H + + HCO 3 - ( Equation 2) (bicarbon ate)
HCO 3 - H + + CO 3 2- ( Equation 3) (carbonate)

There are two possible ways for an aqueous carbonate system to be controlled . In an open system, the water is in equilibrium with a gase ous carbon dioxide phase. The total amount of dissolved inorganic carbon changes as the pH changes. In a closed system

, there is no gaseous phase and the total inorganic carbon is fixed , and only the speciation of the carbon species changes with changes in pH. In carbon dioxide movement associated with carbon dioxide sequestration, a gaseous phase is likely to exist because the large amounts of carbon dioxide present are likely to exceed the
saturation level for carbon dioxide in water. Therefore this p aper assume s an open system is the most representative system for dealing with carbon dioxide leaking into a USDW

.

In an open system
, if the partial pressure of carbon dioxide is constant and known and if either the pH or the alkalinity of the system is known, concentrations of carbonate species can be calculated from the known equilibrium constants for equations 1 through 3. A diagram can then be constructed which shows the relationships of v arious carbonate species versus pH. Figure 1 shows a plot of the carbonate species for an aqueous solution in equilibrium with a partial pressure of carbon dioxide equal to 1 atmosphere (atm)

.

-20 -15 -10

-5 0 5 10 15

0 5 10 15



### pHLog (Concentration (moles/L)

H2CO3

HCO3

CO3

Figure 1. Relationship of log concentration of different carbonate species to pH for an open system
wit h a 1 atm partial pressure of carbon dioxide (adapted from Stumm and Morgan , 1981 ).

Figu re 2 further illustrates how relative concentrations of different carbonate species change with pH in a closed system

. Carbonic acid and its dissociation p roducts are not
, in themselves, toxic to humans. It is, however, a weak acid and lower s the pH , which can a ffect water quality.Page 4 Fig u re 2. Relative concentrations of different carbonate species as pH changes (adap t ed from Stumm and Morgan , 1981 ).

2.1.1 pH Changes


### A s stated above, carbon dioxide, when dissolved in water , forms carbonic acid
, a weak acid. Therefore it is logical to expect a decrease in pH on the introduction of carbon dioxide into a
USDW . The change in pH will depend on the original pH of the water, the amount of carbon dioxide introduced, and the buffering capacity of the water.

LBNL performed a modeling study which examined the effects of movement of carbon dioxide into shallow ground water. They model ed a flow of about 2.4 tons per year and an aquifer with representative water quality and geology based on an analysis of 38,000 ground w ater samples taken throughout the nation. They found that the pH dropped from a value of approximately 7.9 to 5.6 (Birkholzer et al

. , 200 8 ). The pH d id not drop further once the water was saturated
with carbon dioxide

, which happened fairly early on in the process. Kharaka et al. (2006) measured the pH change in a Texas carbon sequestration test site

. They found that injection of
carbon dioxide resul ted in a pH reduction of the formation from 6.5 to 5.7. Several other
authors have modeled the pH changes that occur in carbon dioxide injection zones . A modeling study by Xu et al. (2007) predicted th at pH may drop to 3.4 in the vicinity of an injection well. Another modeling study by Pruess et al. (2002) predicted a pH of 3 to 5 near the injection well. Kaszuba et al. (2005) performed laboratory experiments exposing a rock

- brine mixture to carbon dioxide at elevated temperature and pressure . They fou nd a steep drop in pH uponPage 5 injection of carbon dioxide. They also found that the pH rose as temperature and pressure ( of
the formation ) declined

.

The findings of the se studies indicate that ground water pH will d ecrease with introduction of
carbon dioxid e. While a decrease in pH may be extreme near the site of an injection well, it is likely to be less pronounced with exposure to lower concentrations of carbon dioxide, lower pressures, and lower temperatures. The studies most applicable to USDW exposur e to carbon
dioxide suggest that a pH drop to a value between 5.5 and 6.0 may reasonably be expected.

2.1.2 Changes in Alkalinity Alkalinity is a measure of water’s capacity to neutralize acids. It is measured by titrating a water sample with strong acid to a previously established point , traditionally the end point for carbonate is around pH 4.3. Mathematically

, alkalinity is the sum of the equivalents of hydrogen ion neutralizing capacity of each base present. In most naturally occurring waters the carbon ate species are typically the dominant factors in determining alkalinity. Carbonate alkalinity is defined as

C alk = 2[CO 3 2- ] + [HCO 3 - ] ( Equation 4)

w here [CO 2- 3 ] is the concentration of the carbonate ion and [HCO - 3 ] is the concentration of the
bic arbona te ion. T he carbonate alkalinity in an open system in equilibrium with gaseous carbon dioxide will depend on the partial pressure of carbon dioxide and the initial alkalinity before the
introduction of the carbon dioxide. C arbonate alkalinity can be expr essed as

C alk = 2P CO2 2K H K a1 K a2 /[H + ] 2 + P CO2 K H K a1 /[H + ] ( Equation 5)

w here :

P CO2 is the partial pressure of carbon dioxide in atm
K H is the Henry’s law constant for carbon dioxide in mol/L*atm

K a1 is the equilibrium constant for equation (2) in (mol/L)

K a2 is the equilibrium constant for equation (3) in (mol/L) [H + ] is the concentration of hydrogen ion in (mol/L)

The tot al alkalinity of the system is a conservative quantity and will not change. The overall equation is :

Alk 0 = C alk + [OH - ] – [H + ] + **Σ** (C WB – C WA ) ( Equation 6)

w her e: Alk 0 is the initial alkalinity in ( mol/L ) [OH - ] is the concentration of hydroxide ion in (mol/L)
**Σ** (C WB – C WA ) is the sum of the equivalents of weak base minus the equivalents of weak acid in (mol/L)Page 6 In many natural s ystems , the weak acid and weak base terms are insignificant, leaving only the initial alkalinity, total alkalinity, the hydrogen ion concentration, and the hydroxide ion
concentration. If the initial alkalinity is known, equation 6 can be solved to determi ne the pH.

Equation ( 5 ) shows that the carbonate alkalinity will increase when carbon dioxide comes into contact with ground water . The extent of that increase will depend on the amount of carbon
dioxide migrating into the aquifer. The increase in carbonate alkalinity , however, will be offset by an increase in hydrogen ions and there may be no change in the overall alkalinity. Kharaka et al. (2006) measured the carbonate alkalinity change at the Frio carbon sequestration test site in
Texas and did fi nd an increase in carbonate alkalinity from 100 to 300 meq/L.

2.2 Impacts of Geochemical Changes on Mineralogy and Metals Concentration in USDWs

This section discusses how changes in ground water chemistry resulting from exposure to carbon
dioxide can affect minerals and metals in the rock matrix and potentially cause leaching into USDWs.

2.2.1 Minerals Dissolution The lower pH and higher concentrations of carbonate species caused by introduction of carbon dioxide into ground water can affect equilibrium relationships of minerals in the aquifer material. Different mineral phases may become more stable with the changed water quality
conditions
, and minerals may either dissolve or precipitate depending on the water quality and minerals present. Dissolution of minerals will release the components of that mineral into
solution. O ther compounds bound to mineral surfaces can also be released upon dissolution of
the mineral.

The LBNL study modeled the movement of carbon dioxide into shallow ground water (Birkho lzer
et al

. , 200 8 ). Based on these typical ground w ater compositions and the mineral constituents likely to be in equilibrium under those conditions, the new equilibrium concentrations of various metals were calculated for carbon dioxide partial pressures ranging from 0.0001 to 1 0 bar. They
found that lead, zinc, and arsenic were greatly affected by carbon dioxide concentration

. Barium and cadmium were mobilized but to a lesser degree. Mercury, selenium, and uranium
were insensitive to carbon dioxide co ncentrations. They found that arsenic could be mobilized to concentrations greater than EPA’s maximum contaminant level (MCL) of 10 ug/L for drinking water at higher carbon dioxide concentrations. Lead was found to approach the drinking water regulatory action level (15 ug/L) in some circumstances
, but did not exceed it. The other metals were mobilized at concentrations less than the drinking water MCL. The authors also found that once water was saturated with carbon dioxide, the concentration of metals released did not
increase further.

Kharaka et al. (2006) measured water quality changes within the injection zone. They found
that minerals containing carbonate and iron near the well were dissolved. The iron concentration near the injection well incre ased from 30 mg/L to 100 mg/L. The authors notedPage 7 mobilization of arsenic and other metals in addition to iron . Several studies have modeled the effect that injected carbon dioxide w ould have on minerals in the injection zone

. M odeling studies by Xu (Xu et al.

, 2005; Xu et al.

, 2004) found that iron minerals such as goethite and
hematite would dissolve as well as minerals such as glauconite and oligoclase. Minerals such as ankerite and siderite were precipitated. Pruess et al

. (2002) found the dissoluti on of annite and
calcite

, followed by the precipitation of feldspar and siderite . Kaszuba et al

. (2003, 2005)
conducted laboratory experiments exposing a mineral brine combination to carbon dioxide at elevated temperature and pressure and found the dissol ution of silicate minerals and the precipitation of magnesite

.

The introduction of carbon dioxide into a USDW can lead to dissolution of minerals in the rock matrix of the aquifer. The minerals that dissolve will depend on aquifer conditions, but iron

ba sed minerals seem likely candidates. Any dissolution of minerals will lead to increased metal
concentrations in the water in that aquifer

. Based on the studies mentioned in this paper, i ron , arsenic, zinc, and lead have been identified as metals that cou ld be released from minerals at levels of concern to drinking water systems.

2.2.2 Change in Speciation of Metals Changes in pH can also change the speciation of metal complexes already in solution. As the pH drops, more protonated species will form at the expense of less protonated species. For example, arsenic (V) changes from AsO 4 3- at high pH to the HAsO 4 2- and H 2 AsO 4 - species at lower pH. The significance of these changes can be in the behavior of the metals both with in the
USDW and during treatment. Species with less negative charges may bind less strongly to native aquifer materials and can react differently through treatment processes.

An increase in carbonate alkalinity will also a ffect metal speciation. An increase in bicarbonate
and carbonate ions in the USDW can lead to the formation of metal

- carbonate complexes, changing the speciation of the metal. If carbonate concentrations are high enough, carbonate minerals may begin to precipitate. Anions liberated from dissolving minerals may also be free to form metal complexes. For example, if pyrite were to dissolve, liberated sulfide molecules could form complexes with metals. Complexation will change the charge of a metal and may form
colloidal material

, both of which can affect the behavior of the metal through treatment

processes.

Another factor that may impact metal speciation is the ionic strength of the solution. I onic strength will affect charge

- based reactions including complexation reactions. Generally

, a higher ionic strength will l ower the concentrations of more highly charged species relative to the same species at lower ionic strengths. Although the ionic strength effect is small in normal
water conditions
, it could become significant if the intrusion of brine or the dissolution of minerals results in total ionic concentrations in the range of 1,000 ppm or greater.Page 8 The following sections discuss metals common ly found in ground water and the likely effect that increased carbonate concentrations will have on metal speciation.

Iro n Iron can exist in two oxidation states, either with a positive charge of two or three. As most ground waters tend to be reducing (Birkholzer et al

. , 200 8 ), ferrous iron ( Fe(II) ) is most likely to be the predominant species. The bare Fe(II) tends to be the dominant species under most
conditions. Hydroxyl complexes can also be significant at higher pH values . Other anions that can complex with iron include phosphate and sulfide. The Fe(OH) 2 solid tends to control the solubility of iron at higher pH. At lower pH

, siderite (FeCO 3 ) and pyrite (FeS) can become the controlling solids. Figure 3 shows a solubility diagram for a system containing 0.001 molar total dissolved carbon species and iron(II). The pH is along the x axis and the concentration of iron along the y axis. In this particular system

, the dissolved species are either Fe 2+ or Fe(OH) + . The Fe(OH) 3 - complex can become predominant at very high pH. Siderite is , however, the controlling solid over most of the pH range. At greater carbon dioxid e concentrations
, the area of stability
for siderite would increase.

Figure 3 : Solubility diagram for Fe(II) in a solution with 0.001 M total inorganic carbon.
S ource: Snoeyink and Jenkins , 1980Page 9 Manganese While manganes e can exist in the II, III, IV, and VII oxidation states, only II and IV are common in
natural environments. In a reducing environment, Mn(II) will be the dominant form. Most manganese will be present as the bare ion. Hydroxyl complexes can be significa nt at higher pH
values. Figure 4 shows a p E- pH diagram of the manganese carbonate system. The pE

- pH diagram plots the speciation of the metal versus redox potential and pH. The y

- axis represents
the oxidation reduction ( redox ) potential of the system, wh ile the x

- axis represents the pH. The figure represents the situation where there is 10 micromolar manganese and 1 mg/L of total inorganic carbon. The Mn(OH) 2 solid is dominant at high pH, while rhodocrosite becomes dominant at mid to low pH. At low pH

, the Mn (II) ion becomes the dominant species. As carbon
dioxide concentrations increase
, the stability area of MnCO 3 will increase, while it will decrease
with lower concentrations of carbon dioxide

.

Figur e 4 : Manganese species in a 10 -5 M manganes e solution with total carbon equal to 10 -3 M.
Source : Stumm and Morgan , 1981

Arsenic
Ars enic can exist in either the III or V oxidation state. The III oxidation state is most likely to be the common oxidation state present in reduced ground w aters. Figure 5 shows a plot of arsenic species varying wit h pH and redox potential. It can be seen that under reducing conditions the arsenite species H 2 AsO - 3 and H 3 AsO 3 dominate

. Solid forms of arsenic are typically iron

- arsenic
minerals such as arseno

- pyrit e (FeAsS). Carbonate is not likely to form any complexes directly with arsenic and is likely to only have an influence through its influence on changes in pH.Page 10 Fig ur e 5 : Arsenic speciation with pH and pE.
So u rce: Welch and Stollenwerk , 2003

Lead
Lead ca n exist in either the II or IV oxidation state. In a reduced water

, the II oxidation state is more likely to be common. Figure 6 shows the behavior of lead and hydroxide species in water. The diagram is for 0.04 M total lead concentration. The y axis represents the fraction of a species represented by a given line. For example

, at around pH 6 lead exists as 50 percent Pb 2+ and 50 percent Pb 4 (OH) 4+ 4 . Phosphate complexes can also be important at middle to low pH if significant phosphate is prese nt. Galena (PbS) is typically the predominant lead solid. As carbonate species are introduced, complexes with carbonate or bicarbonate can be significant in
the mid

- pH range. Cerussite (PbCO 3 ) can become the dominant solid under high carbonate

concentrat ions.Page 11 Fig ur e 6 : Distribution of Pb species in a 0.04 M solution of Pb(II).
S o urce: Snoeyink and Jenkins , 1980

Zinc
Zi nc exists in the II oxidation state. It can form complexes with 1, 3, or 4 hydroxyl groups. The pri mary solid formed by zinc is typically sphalerite (ZnS). Figure 7 shows a solubility diagram in a 1 mg/L carbonate system. With increased carbonate concentrations, smithsonite (ZnCO 3 ) solid may precipitate at low pH or hydrozincite (Zn 5 ( CO 3 ) 2 (OH) 6 ) at in termediate pH. In the presence of significant silica concentrations

, hemimorphite (Zn 4 Si 2 O 7 (OH) 2 H 2 O) may also become significant.

Fig ur e 7 : Zinc solubility with 1 mg/L total carbon species. Source: Stumm and Morgan , 1981Page 12 Cadmium
Cadmium exists in the II oxidation state and is present as the bare ion. The solid formed by cadmium under reducing conditions is generally greenockite (CdS). It can also be found as an impurity in sphalerite. At increased carbonate concentrations

, otavite (CdCO 3 ) can become the predominant solid form.

Selenium
Selenium can exist in the

– II, IV, and VI oxidation states. In reducing waters
, the

– II and IV oxidation states are more likely. The IV species will likely be present as HSeO 3 - or SeO 3 2- . The

– II species is likely to be present as HSe - or H 2 Se. Solid species are likely to be analogues of sulfide
materials with Se substituted for S.

Bari um Barium is present in the II oxidation state. It is generally present as the bare ion. Usually the mineral barite (BaSO 4 ) controls the solubility. At high carbonate concentrations

, witherite (BaCO 3 ) may control solubility.

Mercury Mercury can be present in the I and II oxidation states. The I state is more likely in a reducing
environment. In the solid state

, mercury m ay be present as HgS, HgSe, or as a trace element in pyrite.

Uranium Uranium is present in the IV and VI oxidation states. The IV oxidation state is most likely in a reducing environment. It is most likely present as the bare cation. It can be present in water as
the UO 2 2+ complex in the oxidized form. It can also form complexes with hydroxyl anions and
carbonate. Solid forms of uranium can be UO 2 or USiO 4 .

Antimony Antimony is present in the III oxidation state. Solid phases are most likely subst itutions in
sulfides.

Copper Copper typically exists in water in either the I (cuprous) or II (cupric) oxidation states, either of which can form insoluble precipitates and soluble complexes. Under strong oxidizing conditions and high temperatures, cop per can exist in Cu III and Cu IV oxidation states (Schock et al

., 1995,
AwwaRF – DVGW
, 1996). Copper can be found in pure metallic form or as a component of many minerals, including sulfides, oxides and carbonates. Figure 8 show s pE- pH plots of the coppe r carbonate system. Copper will typically be present as a solid, but as the pH drops, release of the bare copper(II) ion becomes likely. This is more likely in situations where dissolved oxygen is
present.Page 13 Figur e 8. EMF

- p H diagram for copper in water containing carbonate, assuming formation of cupric and cuprous hydroxide at DIC = 4.8 mg C/L, I = 0, copper species concentrations of 1.3 mg/L, and 25 ° C Source: Schock et al

., 1995

2.3 Salinity Changes

Carbon dioxide injected into a n injection formation for GS will displace some of the formation fluid s

. Typically these fluids are brines with high total dissolved solids (TDS) concentrations. If carbon
dioxide migrates toward a USDW , some of the high salinity formation fluids may also be displaced by
th e carbon dioxide plume or its pressure front into the USDW. T he dissolution of minerals in the USDW by the reduced pH caused by the carbon dioxide can also increase TDS in the water. TDS
does not have any effect on health but waters with high TDS concentr ations are unpalatable. The
US EPA secondary standard for drinking water is 500 ppm TDS . Table 1 shows TDS concentrations of
interest.Page 14 Table 1

. Noteworthy total dissolved solids (TDS) concentrations

TDS Level Concentration (ppm)

US EPA Drinking Water S econd ar y Standard 500

WHO Standard 1,000

Typ i cal Ground w ater C oncentrations 200 – 3,000

Limit for Aquifer to be considered USDW (per
SDWA ) 10, 000

Seawater 35,500

In addition to causing an unp leasant taste , high TDS increases the ionic strength of th e water. Higher ionic strength can affect the complexation reactions of metals, lowering the charge resistance between molecules and favoring lower charged particles. Very high TDS levels can
lead to the precipitation of metals at lower concentrations than would normally occur.

Kaszuba et al. (2003) performed laboratory experiments exposing a sandstone and brine material representing a saline formation to carbon dioxide at high pressure and temperature.
The researchers found increased chloride concen trations in the brine as a result of mineral
dissolution.

3. 

### Impa ct of USDW Water Quality Changes on Drinking Water Treatment Technologies

The chemical effects of carbon dioxide mixing with the source water of drinking water systems will
affect drinking water treatment processes. The following sections describe how treatment technologies commonly used at ground water

- based drinking water systems ( e.g., disinfection, iron and manganese removal) would be affected if source water quality changes. Note that this chapter

can not cover all conceivable technologies and is limited to technologies that are commonly used in drinking water systems

.

3.1 Disinfection

Disinfection is used to inactivate microbial pathogens, rendering them harmless. Most disinfecta nts are chemical oxidants which destroy microbes by oxidation. Common chemical disinfectants include chlorine, chloramines, chlorine dioxide

, and ozone. UV light can also be used as a disinfectant.

The disinfection effectiveness of chlorine, the most common disinfectant used for ground w ater disinfection, is dependent on pH

. In aqueous conditions , gaseous chlorine dissolves into water under the following equilibria:Page 15 Cl 2 (g) + H 2 O  2HOCl - + 2H + ( Equation 7)
HOCl -  H + + OCl 2- ( Equation 8)

The H OCl molecule is a more effective disinfectant than the OCl 2- molecule. A decrease in pH will shift equation 8 to the left forming more HOCl - and resulting in a higher disinfection efficiency. Table 2 shows the changes in the required contact time to ach ieve 3 log Giardia inactivation using
chlorine at 10 o C for various pH levels

.

Table 2. 3 Log Inactivation of Giardia by Chl orine at Varying pH. pH 6.0 6.5 7 .0 7.5 8.0

Contact Time 79 94 112 134 162

Contact time is based on a residual of 1mg/L

. S ourc e: EPA 1991

pH c an also affect disinfection byproducts (DBP ) formatio n which may result from the use of chlorine. Several studies have found that lowering the pH can decrease the formation of
trihalomethane s (Edwards and Reiber

, 1997; Reckhow and Si nger , 1985). Studies have indicated

, however, that lower pH may increase the formation of haloacetic acids (Reckhow and Singer

, 1985). Lower pH may also affect the production of DBPs by ozone ; l ower pH will decrease formation of bromate by ozone but inc rease the production of brominated organic compounds.

In addition to effects caused by pH, increased metals concentrations can affect disinfection. Reduced metals can consume disinfectants leaving less disinfectant for inactivating microbial pathogens. If the increased metals are not removed before disinfection, water systems will need to
increase their disinfectant concentration in order to achieve the same disinfection efficiency. If UV disinfection is used, metals such as iron and calcium may form s cales on the lamp sleeve reducing the efficiency of the lamp. This will require more frequent cleaning of the sleeve or may require pretreatment to remove the metals if the problem is significant

. Many metals are also less soluble
in their oxidized forms

. I f filtration is not practiced, an increase in metals concentration may
increase the extent to which the metals may precipitate out after the disinfectant addition point.

3.2 Iron and/or Manganese Removal

Iron and manganese removal are typically achieve d by one of two closely related mechanisms : (1) adsorption and (2) precipitation and filtration. Adsorption involves adsorbing i ron and manganese onto media which favorably binds the metals. In precipitation and filtration

, the iron and

manganese are pre cipitated by oxidation. Oxidants include chlorine, permanganate, oxygen, and ozone. Often a ferric based coagulant is used to aid coagulation of the precipitated metals. The precipitated metals are then removed using a filter. Manganese greensand is a combination of the (min)Page 16 two processes. It involves a filter with a coating of manganese oxide on sand. Iron and manganese

are oxidized by the manganese oxide coating and are then adsorb ed to the media

. The media is regenerated by adding permanganate to the filter media.

Both the oxidation and adsorption mechanisms commonly used for iron and manganese removal are affected by pH. Manganese oxidation is rapid above pH 8 and slows at lower pH . Iron oxidation is also slower at lower pH. Greensand filters do not typically regenerate properly at a pH below 6.2. Systems using an alum coagulant for filtration may also experience aluminum carryover problems at pH value s between 6.2 and 6.5.

Adsorption is also affected by pH. Adso r ption typically depend s on t he charge of the surface of the media and the charge of the adsorbing species. As pH drops

, the iron and manganese species become more positively charged. If the surface of the media is also positively charged, this will result in a lower adsorption rate and require more frequent replenishing of the media. Likewise, changes in metal speciation ( described in Section 2.4 ) can affect adsorption rates. If the new species is more positively charged than the species present before the carbon dioxide is introd uced
, adsorption rates may be lowered and treatment may become less effective

.

Increased metals concentrations may also have an effect on iron and manganese removal. The increased concentrations will increase the loading rate upon the process and requi re shorter run times before the process must be renewed by regeneration and backwashing. (Run time refers to the length of time the process can be run before the contaminant begins to break through and the media needs to be regenerated.) Materials such as vanadium, silica, and phosphate have been shown to interfere with some adsorptive media. If these materials are released by introduction of carbon dioxide, run times will decrease and media will need to be replaced more frequently.

3.3 Corrosion Control

Ma ny systems have corrosion control programs in place in order to comply with the Lead and Copper Rule (LCR) and to provide extended life for distribution system components. M any corrosion control programs involve maintaining a set pH and occasionally add in g phosphate or silicate based
corrosion inhibitors. D ecreased pH from exposure to carbon dioxide may require increasing the amount of base fed into the system in order to maintain the desired pH to minimize release of lead and copper from distribution and premise piping materials

. Increased concentrations of these
metal s may also consume corrosion inhibitors and lessen their effectiveness , or requir e higher doses
of inhibitors to control metal s release

.

3.4 Water Softening

Water s oftening involves the pre cipitation of calcium and magnesium from solution at high pH. Softening is usually accomplished by raising the pH of the water with chemicals such as lime or sodium hydroxide. Therefore the lower pH caused by exposure to carbon dioxide would requirePage 17 grea ter addition of chemicals in order to obtain the proper pH. In addition

, dissolved carbon dioxide
in the water may precipitate out as calcium carbonate, consuming additional base.

3.5 Ion Exchange

Ion exchange treatment can be either cation or anion exchange. During ion exchange treatment, undesired ions such as metals are exchanged for ions such as sodium. Cation exchange is frequently used to soften water by removing calcium ions and replacing them with sodium ions. Some ion exchange resins (e.g. alum ina) are pH sensitive and their capacity depends on the pH of the solution. Therefore changes in pH could alter the run time between regeneration cycles. If increased ions are
in the water because of exposure to carbon dioxide , t hese increased concentrations can also reduce
the run time requiring more frequent regeneration to prevent contaminant breakthrough. Increased concentrations of ions (e.g. barium in a cation exchanger) that bind more strongly than the target ion can also displace the target ion causing a spike in its concentration in the effluent.

3.6 Aeration

In aeration, air is contacted with water either through injection of air through diffusers or by flowing air and water over physical structures designed to enhance the mixing of water and ai r. The air contacting the water can remove volatile contaminants from ground water. While pH and metal changes associated with a carbon dioxide plume are unlikely to affect the efficiency of the process,
metals added to the water as a result of exposure to carbon dioxide may be oxidized by the aeration. Oxidized metals such as iron and manganese can precipitate out and

, if filtration is not provided , may carry over and show up as sediments in the distribution system.

3.7 Membranes

Several types of membra ne technology are available. They are characterized by the size of the pore openings in the membrane. Microfiltration membranes and ultrafiltration membranes have larger pores and can remove particles, some microbes, and larger organic molecules. Nanofil tration
membranes (NF) ha ve smaller pores and can remove all particulates and microbes along with colloids and larger molecules. Reverse osmosis membranes (RO) do not have pores and can
eliminate m any larger molecules

, including metals.

Some membrane mat erials are pH sensitive and will degrade at low pH. Scaling is also a significant problem with most membranes. Metals such as calcium, strontium, barium, silica and iron are known to cause fouling of membranes. If exposure to carbon dioxide were to caus e an increased concentration of scaling metals

, it would require more frequent cleanings of the membranes with
associated costs for cleaning materials. It would also likely result in shorter membrane life and necessitate more frequent membrane replacement

.Page 18 If increased metals are released as a result of exposure to carbon dioxide , the effect on membrane

treatment will depend on the type of membrane used. Most metals will pass through the pore openings of membranes used for treatments other than RO . Th erefore additional treatment may be required to remove the metals. RO membranes will remove most metals, so the most likely effect

may be a more concentrated waste stream requiring disposal and the possibility of scaling.

Some smaller metals such as bo ron can pass through standard RO membranes. If there is increased
boron as a result of exposure to carbon dioxide , special membranes can be used or a pretreatment process such as adsorption can be employed.

4. 

### Potential Ground W ater Treatment Scenarios

Th is chapter examine s how the water quality changes discussed in Chapter 2 will likely affect
treatment processes in place at ground water

- based drinking water systems and new treatment that might be necessary in order for a drinking water system to continue to deliver water that meets all SDWA MCLs and secondary standards. Selection of appropriate treatment depend s on the salinity of the USDW or the presence of other treatment in place, e.g., to address high TDS levels or other water quality issues in the U SDW. The treatment discussed includes both the modification of existing treatment at the drinking water system and/or the installation of new treatment where appropriate.

As noted above, it is not possible to cover every potential combination of USDW water quality and treatment technologies that may be encountered. Instead

, this chapter considers three scenarios
that are based on representative ground water quality and drinking water treatment technologies that might occur in the U.S. The three scenari os are:

**•** Ground water below 500 mg/L TDS ( the secondary standard ) with existing drinking water treatment, e.g., disinfection and iron removal processes (section 4.1) ;

**•** Ground water below 500 mg/L TDS with no existing treatment (section 4.2); and

**•** Ground w ater above 2,000 mg/L TDS with RO treatment (section 4.3)

.

4.1. Scenario 1: Ground Water Below 500 mg/L TDS with Existing Drinking Water Treatment

Th is scenario was chosen because it represent s the situation that is most common among larger
water systems , i.e. a good quality aquifer with a TDS below the secondary drinking water standard (< 500 mg/L) and all contaminants at levels below the MCL. Most large drinking water systems that use ground water sources at least use disinfection technologies, and many are required to remove iron and manganese.Page 19 Treatment Options

In this scenario

, the carbon dioxide is assumed to move into a low TDS aquifer that is subject

to basic treatment, which would typically include disinfection along with treatment for any
min eral constituents such as iron. I t is assumed that liquid hypochlorite disinfection and an iron removal technology such as greensand filtration are in place. The drinking water

t reatment changes necessary to continue to meet water quality standards will include chemical and operational adjustments to the existing treatment train and possibly implementation of additional treatment techniques in order to maintain finished water quality treatment goals

.

4.1.1. pH Adjustment pH affects most aqueous reactions. Ther efore

, any chemical treatment process used by a drinking water system will be affected by t he lower pH caused by contact with carbon dioxide

. The treatment effects of lower pH will vary

, with some being positive and other s detrimental.
Positive effects inc lude increased chlorine disinfection efficiency and decreased trihalomethane
formation. Potentially negative effects of lower pH include increased corrosion of metal piping and components used in the treatment and distribution of water , increased lead and copper release, increased production of haloacetic acids, lower efficiency of iron and manganese removal processes, and carryover of aluminum coagulants (EPA, 2007).

Because the lower pH caused by contact with carbon dioxide can have impacts on treatment
systems , many of which are negative , the best option is probably to raise the pH of the water to its level before exposure to carbon dioxide prior to its entering the treatment plant

. This would require the addition of chemical storage tanks and a chemical feed system. Typical chemicals which could be added include sodium hydrox ide (caustic soda) and lime.

4.1.2. Shorter Run Times Exposure to carbon dioxide could significantly increase the concentrations of iron and

manganese a nd other metals in the water

. Inc reased metals concentrations will increase the loading rate on filters and adsorptive media. Increased filter loading can lead to premature breakthrough of particles.

The solution to increased loading rates is generally to increase coagulant dose and/or decrease filter run times. Decreased filter run times require more frequent backwashes. More fre quent backwashes will lead to more solids which require disposal and will lead to higher loss of water which can be an issue if the water system is already producing its well near capacity. Water loss can be mitigated by recycling backwash water. If adsorptive media have higher loading rates, the media will need to be regenerated more frequently to prevent breakthrough. This will result
in additional costs for ne w media or media regeneration.Page 20 4.1.3. Membrane Fouling Although membrane treatment is not common ly used for iron and manganese treatment, it can
be used in some situations. If exposure to carbon dioxide results in the release of minerals into the water, certai n minerals can precipitate on the concentrate side of the membrane and result in fouling of the membranes. Minerals such as calcium, barium, strontium, and silicon can cause precipitates which result in membrane fouling.

The solution for increased foul ing is to clean the membranes more frequently. Alternatively

, pretreatment could be added prior to the membranes. Pretreatment options vary depending on the metals present but could include cartridge filters, softening, or nanofiltration membranes.

4.1.4. TD S Removal High TDS water could be moved into a USDW by a carbon dioxide plume or pressure front. Dissolving minerals could also contribute to a rise in TDS. TDS could exceed the secondary
standard for TDS and may impact the taste of the water

.

Solut ions to increased TDS include blending with lower TDS waters and RO treatment

. Blending requires another source of water with relatively low TDS and is only effective for relatively small reductions in TDS. If substantial TDS reduction is required

, RO memb ranes are the only treatment option . If a system installs RO
, it will need to find a way to handle the treatment residual, which will have very high TDS concentrations along with elevated concentrations of any metals found in the water. Disposal options may include surface water disposal through a National Pollutant Discharge Elimination System ( NPDES ) permit, disposal to the ocean, injection, disposal to a sewer, or treating the treatment residual and recovering the minerals.

4.1.5. Additional Treatment for Bo ron Removal
Boron was not examined in the LBNL study on exposure to carbon dioxide ; however

, it is possible it might be released as the result of such an event. It is not currently regulated by EPA, but the WHO has set a provisional guideline of 0.5 mg/L. It can be a problem as it passes through many treatment processes. Treatment o ptions for boron removal include specially designed RO membranes (Taniguchi et al

. , 2004) and adsorption (Ferreira et al

. , 2006; Okay et
al

. , 1985).

4.1.6. Additional Treatment for Silica Removal Silica might be released from silicate minerals during exposure of a USDW to carbon dioxide

. Silica is not regulated but can be problematic at high concentrations because it can tend to precipitate out and form scales. It can limit the abi lity to use technologies such as membranes due to the scaling nature of the silica.

Treatment options for silica depend on the nature of the silica and its concentration. Colloidal or particulate silica may be amenable to particle removal technologies s uch as coagulation/Page 21 filtration. Dissolved silica will require other technologies such as l ime softening, ion exchange,
and membranes.

4.1.7. Additional Treatment for Arsenic Removal
If arsenic is liberated from the rock matrix of the USDW due to exposure to ca rbon dioxide

, the system will need to treat to remove the arsenic if levels rise above the MCL which is 10 **μ** g/L.

Treatment for arsenic is well known and includes adsorption technologies such as activated aluminum, zero valent iron, several proprietary a dsorption media, ion exchange, and oxidation followed by coagulation and precipitation. NF membranes may also be used (Chinn et al

. , 2002). An option for in

- situ treatment of arsenic has also been developed where oxygen is injected into
the USDW to oxidi ze the arsenic to arsenic(V) and cause it to precipitate out with iron (Miller

, 2006).

4.2. Scenario 2: Ground Water Below 500 mg/L TDS with No Existing Drinking Water Treatment

Although most large drinking water systems that use a ground water source a t least use disinfection technologies , many small ground water

- based systems do not treat the water at all. EPA figures show that less than 60 percent of ground water systems serving less than 100 people disinfect their water and only 20 to 30 percent of non

- community systems disinfect ( US EPA
, 1996). Systems that are not treating their water are assumed to have high quality water with a TDS under the secondary standard (500 ppm) and all other constituents under the MCL.

This scenario was chosen becaus e systems that are not currently treating may be the most
vulnerable to carbon dioxide or fluid movement into the USDW . Because there are no additional treatment barriers, any changes in water quality would be directly experienced by the consumers. If

an y water quality parameters drastically change, the system will likely need to add treatment or seek another source. These systems would likely face the highest costs in the event of a carbon
dioxide leak

, and this scenario would likely represent a worst c ase scenario.

Treatment Options

In this scenario

, the carbon dioxide is assumed to move into a low TDS aquifer which requires no additional treatment prior to distribution. A ddition of new treatment would likely be required

to remove contaminants that m ight be released to the water after exposure to carbon dioxide, to remove TDS that might increase, or to increase pH to minimize release of lead and copper from plumbing materials. Treatment options to address TDS, arsenic, boron and silica are
described in Scenario 1 (in section 4.1); additional treatment that may be needed under this
scenario is described below.Page 22 4.2.1. Uranium Removal
If uranium is released due to exposure to carbon dioxide

, it will need to be removed if it rises
above the MCL of 30 **μ** g/L. Although Birkholzer et a l. (2008) found that uranium was not likely to be mobilized under conditions representative of most ground waters, it is possible that other geologi c formations exist where uranium could be mobilized.

Options for uranium treatment include ion exchange (Clifford and Zhang, 1994; Jelinek and Sorg

, 1988), softening, and coagulation/ precipitation (Sorg

, 1988 ; White and Bondietti, 1983). For softening and filtration

, it is important to control the pH in order to remove uranium. Bec ause carbon dioxide is also likely to lower the pH, pH adjustment will likely be required.

4.2.2. Barium Removal
The LBNL study revealed that barium is slightly sensitive to release by exposure to carbon
dioxide . It seems unlikely the MCL of 2 mg/L would be e xceeded as a result of exposure to carbon dioxide

.

If the barium level did increase to the point where treatment were necessary, options for treatment include anion exchange, softening, and precipitation and filtration (Krause and
Stover , 1982).

4.2.3. Se lenium Removal Selenium was identified as being unlikely to be released to a significant degree from exposure to carbon dioxide (Birkholzer et al

. , 2008). It is possible ; however

, that selenium might be released
under different conditions than those studie d. If selenium should exceed the MCL of 50 **μ** g/L, adsorption could be the most effective treatment.

4.2.4. Iron and Manganese Removal Iron and manganese are common minerals and are likely to be released by exposure to carbon
dioxide . Neither one is a health threat, but both have secondary standards bec ause of aesthetic concerns. The standards are 0.3 and 0.05 mg/L for iron and manganese

, respectively.

Treatment for iron and manganese consists of oxidizing the metals, precipitating them out

, and removal by filtration. Greensand filtration is one means of doing this which uses a special manganese coated filter media to improve removal. Oxidation of iron and manganese can be accomplished using aeration, chlorination, permanganate, or ozone.Page 23 4.2.5. Lead Removal There are two ways by which carbon dioxide co uld lead to the release of lead. The first is through the dissolution of lead containing minerals within the USDW formation . The second

case is the lower ed pH by carbon dioxide could cause release of lead from lead plumbing in the distribution system.

Lea d released into a USDW from the formation rocks can be removed by softening, adsorption, ion exchange, RO , or coagulation and filtration. Options for preventing lead release from plumbing include adjusting the pH, adding corrosion inhibitors such as orthop hosphate, or replacing lead plumbing.

4.3. Scenario 3 : Ground Water above 2,000 mg/L TDS and RO Treatment

While some private supplies may use water with TDS as high as 2,000 ppm (North Dakota State University
, 2009) without treatment, most municipal syst ems will treat water that is greater than 1,000 ppm or look for alternative sources (USGS

, 2005). In this scenario

, the carbon dioxide is
assumed to move into a high TDS aquifer that already requires treatment to reduce the TDS to make the water palatable

.

This scenario was chosen to represent the case of systems with limited water supplies where saline ground waters are now being examined as drinking water sources. Sandia National L aboratory predicted that

, by 2020

, desalinated saline ground water or se awater will provide a significant portion of drinking water in the United States (Sandia National Laboratory

, 2003).

Treatment Options

This scenario examine s a USDW with a TDS higher than 2,000 ppm. The existing treatment is a RO membrane system

, as t hat is the most feasible treatment for removing chloride and other common components of TDS. The calcium and sulfate concentrations are assumed to be low enough that gypsum scale formation is not a problem, since utilities will generally try to avoid
such water if possible. It is assumed that the treatment residual is disposed of by means of a sewer connection. Some sort of pretreatment is also necessary to remove solids and other
materials that could scale and clog the membranes. It is assumed that the system has installed a cartridge filter to remove colloidal and particulate matter prior to the membranes.

A ddition of new treatment would not be likely, since RO already in place w ould remove nearly any contaminant that might be introduced. Most of the changes necessary will be adjustments to the existing train in order to prevent poor treatment performance of the RO system and other operational issues.Page 24 4.3.1. pH

- Resistant Membranes Some membrane materials degrade under acidic conditions. Specifically, cel lulose based membrane materials must operate between pH of 4 and 8 (US EPA
, 2005). Polyamide membranes do not have this restriction but are very sensitive to strong oxidants such as
chlorine. If contact with carbon dioxide lowers the pH below 4

, cellulos e membranes may be damaged.

If the pH becomes inappropriate for the membranes being used

, options are either to replace the membranes with a more resistant material or to adjust the pH back into the acceptable range before the water passes through the me mbranes.

4.3.2. Membrane Fouling
If exposure to carbon dioxide results in the release of minerals into the water, certain minerals can precipitate on the concentrate side of the membrane and result in fouling of the membrane. Minerals such as calcium, barium, strontium, and silicon can cause precipitates which result in membrane fouling. Increased fouling would result in higher costs and more down time for membrane cleaning as well as more frequent membrane replacement.

The simplest treatment option to co unter fouling problems is more frequent membrane cleaning. If membrane cleaning becomes too frequent or results in too much down time , the system might need to add treatment to remove the minerals prior to the RO membranes. Cartridge filters typically use d for pretreatment are not likely to remove dissolved constituents such as calcium. Softening or NF membranes could remove these minerals (EPA 2005)

.

4.3.3. Additional Treatment for Boron and Silica
If additional boron is liberated the rock matrix of the USDW

, this can cause problems for
traditional RO membranes. The rejection rate of boron from traditional RO membranes is less than of other salts and boron may pass through RO membranes (Magara et al

. , 1998).

Treatment o ptions to remove additional boron include: a second RO unit in series (Magara et al

. , 1998), adsorption (Polat et al

. , 2004; Okay et al

. , 1985; Ferreira et al

. , 2006), or special membranes designed for boron (Taniguchi et al

. , 2004). Silica can also foul membranes and may require pretreatmen t or additional membrane cleaning.

4.3.4. TDS Removal Since the system is assumed to already be removing TDS, additional increases caused by fluid
movement associated with GS will not require additional treatment. If concentrations drastically increase , then t he removal efficiency of the RO unit may decrease.

Decreases in removal efficiency can be countered by increasing the trans

- membrane pressure of
the RO unit.Page 25 4.3.5. Potential Waste Disposal Issues
RO produces a treatment residual which contains elevated concen trations of the contaminants removed from the water stream. Generally

, the TDS concentration and concentration of other contaminants is too high for disposal to surface water sources without further treatment. Systems near an ocean can dispose of the concentrate in the ocean. Other systems will generally either dispose of the concentrate to a sewer line, an injection well, or evaporation ponds. While in this case the system will already have disposal in place, increases in
concentrations of TDS and meta l concentrations will further increase the concentrations in the concentrate to be disposed. This may result in additional fees for disposal if a sewer is used. Depending on the metals which are liberated as a result of exposure to carbon dioxide

, it is also possible that the concentrate will contain concentrations of metals which exceed MCLs or other threshold concentrations even if the concentrations are lower in the aquifer itself. Such elevated concentrations could limit disposal options necessitatin g treatment of the concentrate before disposal.

If the waste stream needs to be treated before disposal, t reatment options include a second RO unit, adsorption, and evaporation ponds.

4.4. Other Issues

4.4 .1 Corrosion

- Resistant Components
The lower pH of water saturated with carbon dioxide can cause increased corrosion of drinking water system components. If high TDS water intrudes along with the carbon dioxide, high chloride concentrations can further accelerate corrosion reactions. Components whi ch might be more susceptible to corrosion with lower pH and higher chloride water include: steel pipes and fittings, pumps, well grouting, cement basins, and any metal parts or fittings. Neoprene and nitrile fittings are also susceptible specifically to carbonic acid (McCoulou et al., 1985).

Systems would need to replace corrosion

- susceptible components with corrosion resistant
materials

, such as stainless steel and Teflon

, which would increase costs to the operator of the drinking water system . For ab ove ground components
, the system could choose to add caustic or lime to raise the pH instead of replacing parts.

4.4.2 D istribution System Corrosion Control Lower pH and increased chloride concentrations can increase corrosion in the distribution
syste m. Both metal and concrete pipes will be susceptible to corrosion under these conditions.

A system is unlikely to replace its entire distribution system with corrosion resistant materials, so systems will more likely choose to either raise the pH of th e water prior to the distribution system or to add corrosion control chemicals to the water. Systems which already use chemicalPage 26 addition for corrosion control for the lead and copper rule may need to adjust their chemical doses upward.

5. 

### Rem ediation Option s to Address Increased C arbon Dioxide Concentrations in G round
W ater

This section examines technologies available to address carbon dioxide that has migrated to a USDW
or other formation that is not currently being used as a drinking water source. In so me cases , it may be desirable to treat ground water that has been exposed to carbon dioxide at a location other than an existing drinking water well. There could be many reasons for such treatment: the formation may not be currently used for drinking wate r, there may be so many wells that installing individual treatment at each one is not cost efficient, or it may be desirable to treat a carbon dioxide plume before it reaches existing drinking water wells. In this case

, remediation of the carbon dioxide

p lume within the formation is desirable . Since there has never been a known large scale release of
carbon dioxide in to a USDW
, not much is understood about such a situation. The following paragraphs describe analogous situations and discuss potential remediation strategies for carbon
dioxide.

T he first step in any remediation process is characterization (i.e., determining the location, size, and characteristics of the plume ), which is necessary in deciding upon a remediation strategy. If stopping the lea k is possible , it will minimize the scope of the remediation effort. If the leak is through a well, the well c an be repaired or plugged. Flows through faults may be able to be plugged by injecting low permeability materials (e.g.

, cement) into the fault (B urton et al., 2008). Often the first step in remediation is to contain the plume by establishing treatment along its front edge. Pumping water into the zone above the confining zone to increase the pressure and/or decreasing the pressure in the injection zone may eliminate the driving force for plume migration and stop plume movement out of the injection zone (Benson and Hepple, 2005).

Plume characterization is also necessary to design a proper remediation plan. In addition to the large size of the plum e, information such as the phase of the carbon dioxide and the pH and metals concentrations of any aqueous phases will be necessary to properly design a treatment scheme.

M ethods for tracking and locating a carbon dioxide plume will be discussed in guidan ce documents that will support the Class VI rule

. Although these techniques are intended to monitor the position
of the plume in the injection zone

, they are also useful for monitoring in shallower formation s

. Methods include Light Detection a nd Ranging ( LIDAR ) , seismic profiles, and electromagnetic techniques. Ground water models may also be helpful to predict future plume movement and speed.

Once the plume has been properly characterized

, a treatment strategy must be designed. In traditional grou nd water remediation, treatment is generally done through pump and treat techniques or by in

- situ treatment. Pump and treat involves drilling wells to pump thePage 27 contaminated water out of the USDW/formation , treating the water, and then returning the water either through injection wells or possibly through infiltration basins for shallow unconfined aquifers. Treatment wells would typically be located at the front edge of the plume in order to contain plume
movement. In

- situ treatment involves treating the ground water without pumping it from the aquifer. It would most likely involve injection of chemicals into the aquifer with the goal of treating the acidity and metals liberated by the carbon dioxide plume.

5.1. Pump and Treat

In a pump and treat scenario, the carbon dioxide plume and affected ground water are pumped from the aquifer, treated

, and returned to the aquifer . If the carbon dioxide plume is still in the gaseous phase, the carbon dioxide may be able to be produced through wells, captured, and re

i njected into a more secure geologic formation (Benson and Hepple, 2005). If the plume is aqueous or has an aqueous component

, the water is pumped out and then treated and returned to the aquifer, although other uses for the treated water are possible. Trea ted water can be returned through percolation basins if it is a shallow, unconfined aquifer ; otherwise injection wells are necessary. Wells to withdraw water from the USDW are typically placed in such a way as to contain the plume when the pump and treat system is operating. In a pump and treat situation, typically one or more wells are placed in such a configuration so that the entire plume is drawn toward the capture zone of the wells (US EPA
, 1996). Determining the capture zone of the well will requir e ground water hydrological modeling. Injection wells are also sometimes placed to form pressure ridges in the ground water which can contain the plume and push it toward the capture zone.

Once the water is pumped from the aquifer, it will require tr eatment before being returned to the aquifer. Treatment will depend upon the characteristics of the water in the aquifer or the nature of the contamination. The water will more than likely have a low pH and unless the selected treatment option performs b etter at low pH, pH adjustment will be necessary. Aeration may also lower the carbon dioxide content of the water and raise the pH (Benson and Hepple, 2005). If high TDS is a problem, then RO treatment will be necessary. As discussed in section 4 .1 .3., o ther considerations may need to be taken into account with the RO system design. pH adjustment or acid resistant membrane material may be necessary. RO systems will generally require pretreatment to reduce solids content and scaling minerals. Typically one or more cartridge filters and an ion exchange unit are used as pretreatment. Silica or boron may require special pre

- treatment such as an adsorption or precipitation step. (See Chapter 4 for a detailed description of these treatment technologies.)

I f TDS is not a problem, then only metals removal and pH adjustment are likely to be the main

concerns . There are many options for metals removal. Section 5.2 discusses some of the treatment options. Raising the pH using materials such as lime, limestone , caustic soda, soda ash, and other
alkaline m aterials , with or without oxidation has proven successful for treating acid mine drainage (AMD) (Skonsen et al

., 1998). Another method that has worked with AMD is reducing the metals by passing them through re active organic materials and then precipitating out the metals as sulfides (Brenner et al

., 2007) . This would require a significant presence of sulfur compounds in the aquifer.Page 28 If the metals cannot be coagulated and precipitated using pH adjustment and o xidation or reduction, some sort of adsorption process may be a good choice. There are numerous media that can be used for metals adsorption ; the selection of media will depend largely on the metals present. Cation exchange resins may be an option if a w ide variety of metals are present.

5.2. In

- situ Treatment

In

- situ treatment is meant to treat contaminated ground water with in the aquifer to save the cost of pumping the water out of the aquifer. Historically most in

- situ treatments have focused on biolog ical treatment of organic compounds ( US EPA 1996). Unlike organics, the metals liberated by a carbon dioxide plume cannot be destroyed. Treatment, instead

, will aim to render the metals immobile by either adsorption to the aquifer material or precipitati on.

5.2.1. pH adjustment
Metals released as a result of exposure to carbon dioxide are likely released because of the acidic nature of the plume. Such reactions are generally reversible, so raising the pH to the original aquifer pH may cause the metals to pre cipitate out and become immobilized

. Chemicals such as sodium hydroxide, potassium hydroxide, and sodium carbonate could be used to raise the pH by injecting them into the aquifer. Calcium and magnesium salts could also work but may be more difficult to handle . Trenches filled with such materials have been used successfully
in shallow contamination events.

5.2.2. M etals precipitation and adsorption

R emov ing metals from the drinking water in the aquifer will require either precipitating the
metals out of so lution or having the metals adsorb to the aquifer material. Raising the pH as
described in section 5.2 .1 is one method of achieving that. Other methods include oxidizing the metals, reducing the metals, or adding chemicals which will form insoluble preci pitates but

t here is little experience with such treatments. Reactive barriers

, such as zero valent iron

, have
been used to react and adsorb metals to material in the barrier. Reactive barriers with organics used as reductants have been used to precipitate sulfides out of AMD (Brenner et al

., 2007). Reactive barriers, however, would be difficult to install in deep aquifers. An alternative to
reactive barriers would be to a dd chemicals into an aquifer to bring about precipitation and/or adsorption. Choosi ng the proper chemical would require knowledge of the chemistry of the metals to be precipitated or adsorbed. For example, in general

, iron, manganese, and arsenic are less mobile in their oxidized form. Many sulfides are insoluble and can be precipitate d out under reducing conditions. Before undertaking such chemical additions, a careful understanding of the geochemistry of the aquifer is necessary to prevent undesirable changes. For example, while oxidizing metals may lead to precipitation of iron, it could lead to release of other minerals
such as uranium. Aquifer clogging can also be an issue if precipitation reactions are initiated. While chemical additions have not been conducted on a large scale, o ne study did find success in removing arsenic from a drinking water wells by injecting air into the aquifer in the areaPage 29 surrounding the drinking water well (Miller, 2006). The air oxidized ferrous iron in the aquifer causing it to precipitate and caused arsenic to adsorb onto the precipitated iron.

6. Summary/ Conclusions

The regulations for Class VI injection wells provide many barriers to prevent changes in the existing water quality of USDWs . However
, it is still possible that unforeseen circumstances could lead to
the movement of carbon dioxide or formation fluids into a USDW that supplies a drinking water system and alter the water quality . While carbon dioxide is not toxic, it can change the chemistry of a USDW, including a drop in pH ; release of metals such as iron, lead, arsenic, and zinc ; an i ncrease in
carbonate alkalinity ; an increase in TDS ; and potential changes in metal speciation. This can affect drinking water treatment processes or necessitate additional treatment to address mobilized
minerals and metals.

The chemical changes in a g round water contacting carbon dioxide can require additional treatment or changes in existing treatment processes to maintain compliance with drinking water MCLs . These potential additions and changes include:

**•** Raising the pH through chemical addition ;

**•** Adding a metals removal process, such as adsorption or oxidation/filtration, or more frequent backwashing/regeneration and chemical use in existing metals removal processes ; and

**•** Addition of RO units to treat high TDS water migrating from other formations or additional pretreatment if RO units are already in place.

In some cases
, it may be desirable to treat the carbon dioxide plume in the aquifer before it reaches drinking water wells. This can be done by pumping the water out of the aquifer, treating it, and returning it to the aquifer through injection wells. Treatment processes would be similar to those applied at a drinking water well and would include :

**•** pH adjustment ;

**•** metals adsorption or precipitation ; and

**•** RO if high TDS is involved.

Alterna tively, i t may be possible to treat the aquifer in

- situ . Although this has never been done in
association with GS

, analogues from acid mine drainage and arsenic problems may point to some possible solutions. Injections of caustic soda, soda ash, or lime into the aquifer can be used to raise the pH and potentially precipitate the metals. In some cases

, adding air or oxygen may oxidize
metals and result in their precipitat ion

.Page 30 Class VI permitting authorities and owners or operators of wells should be aware of the potential impacts of reactions between carbon dioxide and water, particularly as they relate to water treatment technologies. This can inform improved management of a GS project to reduce the potential for such events or, if necessary, facilita te emergency and remedial response to an adverse

event.

7. References

AWWARF

- DVGW. 1996. Internal Corrosion of Water Distribution Systems 2 nd Edition. AWWARF, Denver, CO

. Benson, S., and R. Hepple. 2005. Detection and Options for Remediation of Leakage f rom Underground CO2 Storage Projects. Presented at the 7 th annual Greenhouse Gas Technology Conference.

Birkholzer, J., Q. Zhou, K. Zhang, P. Jordan, J. Rutqvist, and C.F. Tsang. 2008. Research Project on CO2 Geological Storage and Groundwater Resources: Large

- Scale Hydrological Evaluation and Modeling of the Impact on Groundwater Systems. NETL Project Annual Report, October 1, 2007 to September 30, 2008.

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