Drinking Water Treatment 

Considerations: Water Quality, 

Carbon Dioxide Concentration, 

and Geologic Sequestration 

Projects
Office of Water (4606M)

EPA 816‐R‐10‐014  

November 2010

http://water.epa.gov/drink/
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Drinking Water Treatment Considerations: Water Quality, Carbon Dioxide Concentration, and 

Geologic Sequestration Projects

1. Introduction

The Federal Requirements under the Underground Injection Control (UIC) Program for Carbon 

Dioxide Geologic Sequestration (GS) Wells (referred to in this document as the Class VI rule) are 

designed to minimize the risk of injected carbon dioxide contaminating an underground source of 

drinking water (USDW). The Class VI rule 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 

plugging, 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 into USDWs, there is a remote potential for the movement of carbon 

dioxide or formation 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, affect 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 responses 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 Chemistry, provides background on the 

chemistry of carbon dioxide-water interactions and discusses 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, discusses how the changes in water quality might affect technologies 

commonly used for drinking water treatment. 

• Chapter 4, Potential Ground Water Treatment Scenarios, presents several ground water 

quality scenarios which are likely to be representative of situations that might occur with a 

leak of carbon dioxide into a USDW, outlines 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 Carbon Dioxide Concentrations in 

Ground Water, discusses how a USDW of degraded water quality due to increased 

concentrations of carbon dioxide may be returned to compliance with drinking water 

standards.

• Chapter 6 presents a summary of the paper and conclusions. 
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Given the wide range of potential geochemical scenarios, this paper cannot describe every possible 

configuration of USDW and rock formation 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 

not intended to be an exhaustive description of all potential situations or responses; rather, it serves 

as a guideline to inform GS project planning.

2. Impacts of Carbon Dioxide on Water Chemistry 

To support an understanding of geochemical changes that may potentially occur in subsurface 

formations and what effects, if any, these 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 and 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 presents 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 pHaltered environments. Section 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. Some of the carbon dioxide will dissolve to form carbonic acid, and through further 

dissociation reactions, bicarbonate and carbonate, as illustrated in equations 1 through 3.
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CO2(g) + H2O H2CO3 (aq) (Equation 1) (carbonic acid)

 H2CO3 H+

 + HCO3

-

 (Equation 2) (bicarbonate)

HCO3

- 

 H+ + CO3

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 gaseous 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 paper assumes 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 various 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

pH

Log (Concentration (moles/L)

H2CO3

HCO3

CO3

Figure 1. Relationship of log concentration of different carbonate species to pH for an open system 

with a 1 atm partial pressure of carbon dioxide (adapted from Stumm and Morgan, 1981). 

Figure 2 further illustrates how relative concentrations of different carbonate species change with 

pH in a closed system. Carbonic acid and its dissociation products are not, in themselves, toxic to 

humans. It is, however, a weak acid and lowers the pH, which can affect water quality. 
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Figure 2. Relative concentrations of different carbonate species as pH changes (adapted from 

Stumm and Morgan, 1981).

2.1.1 pH Changes

As 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 modeled 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 water samples 

taken throughout the nation. They found that the pH dropped from a value of approximately

7.9 to 5.6 (Birkholzer et al., 2008). The pH did 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 resulted 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 that 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 found a steep drop in pH upon 
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injection of carbon dioxide. They also found that the pH rose as temperature and pressure (of 

the formation) declined. 

The findings of these studies indicate that ground water pH will decrease with introduction of 

carbon dioxide. 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 exposure 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 carbonate 

species are typically the dominant factors in determining alkalinity. Carbonate alkalinity is 

defined as

Calk = 2[CO3

2-] + [HCO3-

] (Equation 4)

where [CO 2-

3 ] is the concentration of the carbonate ion and [HCO -3 ] is the concentration of the 

bicarbonate ion. The 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. Carbonate alkalinity can be expressed as

 Calk = 2PCO22KHKa1Ka2/[H+

]

2

 + PCO2KHKa1/[H+] (Equation 5)

where :

PCO2 is the partial pressure of carbon dioxide in atm

KH is the Henry’s law constant for carbon dioxide in mol/L*atm

Ka1 is the equilibrium constant for equation (2) in (mol/L)

Ka2 is the equilibrium constant for equation (3) in (mol/L)

[H+

] is the concentration of hydrogen ion in (mol/L)

The total alkalinity of the system is a conservative quantity and will not change. The overall 

equation is: 

 Alk0 = Calk + [OH-

] – [H+] + Σ(CWB – CWA) (Equation 6)

where: 

Alk0 is the initial alkalinity in (mol/L) 

 [OH-

] is the concentration of hydroxide ion in (mol/L)

Σ(CWB – CWA) is the sum of the equivalents of weak base minus the equivalents of weak 

acid in (mol/L)
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In many natural systems, 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 determine 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 find 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. Other 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 (Birkholzer 

et al., 2008). Based on these typical ground water 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 10 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 concentrations. 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 increased from 30 mg/L to 100 mg/L. The authors noted 
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mobilization of arsenic and other metals in addition to iron. Several studies have modeled the 

effect that injected carbon dioxide would have on minerals in the injection zone. Modeling 

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 dissolution 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 dissolution 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 ironbased 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, iron, 

arsenic, zinc, and lead have been identified as metals that could 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 AsO4

3- at high pH to the HAsO42- and H2AsO4-

species at lower 

pH. The significance of these changes can be in the behavior of the metals both within 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 affect 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. Ionic 

strength will affect charge-based reactions including complexation reactions. Generally, a 

higher ionic strength will lower 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. 
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The following sections discuss metals commonly found in ground water and the likely effect that 

increased carbonate concentrations will have on metal speciation.

Iron 

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., 2008), 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 (FeCO3) 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 Fe2+ 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 dioxide 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.

 Source: Snoeyink and Jenkins, 1980
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Manganese

While manganese 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 significant at higher pH 

values. Figure 4 shows a pE-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, while 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 MnCO3 will increase, while it will decrease 

with lower concentrations of carbon dioxide. 

Figure 4: Manganese species in a 10-5 M manganese solution with total carbon equal to 10-3 M.

 Source: Stumm and Morgan, 1981

Arsenic 

Arsenic 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 waters. Figure 5 shows a plot of arsenic 

species varying with pH and redox potential. It can be seen that under reducing conditions the

arsenite species H2AsO -

3 and H3AsO3 dominate. Solid forms of arsenic are typically iron-arsenic 

minerals such as arseno-pyrite (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. 
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Figure 5: Arsenic speciation with pH and pE.

 Source: Welch and Stollenwerk, 2003

Lead

Lead can 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 Pb2+

and 50 percent Pb4(OH) 4+

4 . Phosphate complexes can also be important at middle to low pH if 

significant phosphate is present. 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 (PbCO3) can become the dominant solid under high carbonate 

concentrations.
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Figure 6: Distribution of Pb species in a 0.04 M solution of Pb(II).

 Source: Snoeyink and Jenkins, 1980

Zinc

Zinc exists in the II oxidation state. It can form complexes with 1, 3, or 4 hydroxyl groups. The 

primary 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 (ZnCO3) solid 

may precipitate at low pH or hydrozincite (Zn5(CO3)2(OH)6) at intermediate pH. In the presence 

of significant silica concentrations, hemimorphite (Zn4Si2O7(OH)2H2O) may also become 

significant. 

Figure 7: Zinc solubility with 1 mg/L total carbon species.

 Source: Stumm and Morgan, 1981
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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 (CdCO3) 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 HSeO3

- or SeO32-. The –II 

species is likely to be present as HSe- or H2Se. Solid species are likely to be analogues of sulfide 

materials with Se substituted for S.

Barium 

Barium is present in the II oxidation state. It is generally present as the bare ion. Usually the 

mineral barite (BaSO4) controls the solubility. At high carbonate concentrations, witherite 

(BaCO3) 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 may 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 UO2

2+ complex in the oxidized form. It can also form complexes with hydroxyl anions and 

carbonate. Solid forms of uranium can be UO2 or USiO4. 

Antimony

Antimony is present in the III oxidation state. Solid phases are most likely substitutions 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, copper 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 shows pE-pH plots of the copper 

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.
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Figure 8. EMF-pH 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 an injection formation for GS will displace some of the formation fluids. 

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 

the carbon dioxide plume or its pressure front into the USDW. The 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 concentrations are unpalatable. The 

US EPA secondary standard for drinking water is 500 ppm TDS. Table 1 shows TDS concentrations of 

interest.
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Table 1. Noteworthy total dissolved solids (TDS) concentrations

TDS Level Concentration (ppm)

US EPA Drinking Water Secondary Standard 500

WHO Standard 1,000

Typical Ground water Concentrations 200 – 3,000

Limit for Aquifer to be considered USDW (per 

SDWA)

10,000

Seawater 35,500

In addition to causing an unpleasant taste, high TDS increases the ionic strength of the 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 concentrations in the brine as a result of mineral 

dissolution.

 

3. Impact 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 disinfectants 

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 water 

disinfection, is dependent on pH. In aqueous conditions, gaseous chlorine dissolves into water 

under the following equilibria:
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Cl2(g) + H2O  2HOCl- + 2H+

 (Equation 7)

HOCl-

 H+ + OCl2- (Equation 8)

The HOClmolecule is a more effective disinfectant than the OCl2- 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 achieve 3 log Giardia inactivation using 

chlorine at 10o

C for various pH levels. 

Table 2. 3 Log Inactivation of Giardia by Chlorine 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. 

Source: EPA 1991

pH can also affect disinfection byproducts (DBP) formation which may result from the use of

chlorine. Several studies have found that lowering the pH can decrease the formation of 

trihalomethanes (Edwards and Reiber, 1997; Reckhow and Singer, 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; lower pH will decrease formation of 

bromate by ozone but increase 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 scales 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. If 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 achieved by one of two closely related mechanisms: (1) 

adsorption and (2) precipitation and filtration. Adsorption involves adsorbing iron and manganese 

onto media which favorably binds the metals. In precipitation and filtration, the iron and 

manganese are precipitated 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 adsorbed 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 values between 6.2 and 6.5. 

Adsorption is also affected by pH. Adsorption typically depends on the 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 introduced, 

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 require 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

Many 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. Many corrosion 

control programs involve maintaining a set pH and occasionally adding phosphate or silicate based 

corrosion inhibitors. Decreased 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 

metals may also consume corrosion inhibitors and lessen their effectiveness, or require higher doses 

of inhibitors to control metals release. 

3.4 Water Softening

Water softening involves the precipitation 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 require 
Page 17

greater 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. alumina) 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, these 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 air. 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 membrane 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. Nanofiltration 

membranes (NF) have 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 many larger molecules, including metals.

Some membrane materials 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 cause 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. Therefore 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 boron 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 Water Treatment Scenarios 

This chapter examines 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 depends 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 USDW. 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 scenarios 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 water 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

This scenario was chosen because it represents 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 

mineral constituents such as iron. It is assumed that liquid hypochlorite disinfection and an 

iron removal technology such as greensand filtration are in place. The drinking water 

treatment 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. Therefore, any chemical treatment process used by a 

drinking water system will be affected by the lower pH caused by contact with carbon dioxide. 

The treatment effects of lower pH will vary, with some being positive and others detrimental. 

Positive effects include 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 hydroxide (caustic soda) and lime. 

4.1.2. Shorter Run Times

Exposure to carbon dioxide could significantly increase the concentrations of iron and 

manganese and other metals in the water. Increased 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 frequent 

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 new media or media regeneration. 
Page 20

4.1.3. Membrane Fouling

Although membrane treatment is not commonly 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, certain 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 fouling 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. TDS 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. 

Solutions 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 membranes 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 Boron 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 options 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 ability 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 such as coagulation/
Page 21

filtration. Dissolved silica will require other technologies such as lime 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 carbon dioxide, the 

system will need to treat to remove the arsenic if levels rise above the MCL which is 10 μg/L. 

Treatment for arsenic is well known and includes adsorption technologies such as activated 

aluminum, zero valent iron, several proprietary adsorption 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 oxidize 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 at 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 because 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 

any 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 case 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. Addition of new treatment would likely be required 

to remove contaminants that might 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 μg/L. Although Birkholzer et al. (2008) found that uranium was not likely 

to be mobilized under conditions representative of most ground waters, it is possible that other 

geologic 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. Because 

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 exceeded 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. Selenium 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 studied. If selenium should exceed the MCL of 50 μg/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 because 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 could 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 lowered pH by carbon dioxide could cause release of lead from lead plumbing in the 

distribution system.

Lead 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 orthophosphate, 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 systems 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 Laboratory 

predicted that, by 2020, desalinated saline ground water or seawater will provide a significant 

portion of drinking water in the United States (Sandia National Laboratory, 2003). 

Treatment Options

This scenario examines a USDW with a TDS higher than 2,000 ppm. The existing treatment is a 

RO membrane system, as that 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.

Addition of new treatment would not be likely, since RO already in place would 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, cellulose 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, cellulose 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 membranes. 

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 counter 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 used 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 options 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 pretreatment 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 the 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 concentrations 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 metal 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 necessitating treatment of the concentrate 

before disposal. 

If the waste stream needs to be treated before disposal, treatment 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 which 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 above ground components, the system could choose to add caustic 

or lime to raise the pH instead of replacing parts. 

4.4.2 Distribution System Corrosion Control

Lower pH and increased chloride concentrations can increase corrosion in the distribution 

system. 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 the water prior to the distribution 

system or to add corrosion control chemicals to the water. Systems which already use chemical 
Page 26

addition for corrosion control for the lead and copper rule may need to adjust their chemical 

doses upward.

5. Remediation Options to Address Increased Carbon Dioxide Concentrations in Ground

Water 

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 some 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 water, 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 

plume within the formation is desirable. Since there has never been a known large scale release of 

carbon dioxide into a USDW, not much is understood about such a situation. The following 

paragraphs describe analogous situations and discuss potential remediation strategies for carbon 

dioxide. 

The 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 leak is possible, it will minimize the scope of the remediation effort. If the leak is 

through a well, the well can be repaired or plugged. Flows through faults may be able to be plugged 

by injecting low permeability materials (e.g., cement) into the fault (Burton 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 plume, 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. 

Methods for tracking and locating a carbon dioxide plume will be discussed in guidance 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 formations. 

Methods include Light Detection and 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 ground water remediation, treatment is generally done through pump and treat 

techniques or by in-situ treatment. Pump and treat involves drilling wells to pump the 
Page 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 reinjected 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. Treated 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 require 

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 treatment 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 better 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., other 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.)

If 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 materials, 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 reactive 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 oxidation 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 wide variety of metals are present. 

5.2.In-situ Treatment

In-situ treatment is meant to treat contaminated ground water within the aquifer to save the cost of 

pumping the water out of the aquifer. Historically most in-situ treatments have focused on 

biological 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 precipitation. 

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 precipitate 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. Metals precipitation and adsorption

Removing metals from the drinking water in the aquifer will require either precipitating the 

metals out of solution 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 precipitates but

there 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 add chemicals into an aquifer to bring about precipitation and/or 

adsorption. Choosing 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 precipitated 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, one study did find success 

in removing arsenic from a drinking water wells by injecting air into the aquifer in the area 
Page 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 increase 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 ground 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. 

Alternatively, it 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 precipitation. 
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, facilitate emergency and remedial response to an adverse 

event. 

7. References

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