User:Debra Tabron/sandbox

Acids are produced during Enhanced Reductive Dechlorination (ERD) which can cause pH to drop, inhibiting treatment. Alkaline materials including hydroxides and carbonates are sometimes added to the aquifer during ERD to maintain a pH of greater than 6, improving bioremediation performance. This article describes a simplified approach and a spreadsheet-based design tool for estimating the total amount of base required to achieve a specified target pH at the end of the treatment period, once all reactions have gone to completion. To prevent overshoot and excessively high pH, users typically add a fraction of the total base required in several increments spread over time.

Related Article(s):

• pH Buffering in Aquifers
• Low pH Inhibition of Reductive Dechlorination
• Bioremediation – Anaerobic
• Chlorinated Solvents

CONTRIBUTOR(S): Dr. Robert Borden, P.E.

Key Resource(s):

• Spreadsheet-based Design Tool -- Base Addition for ERD
• Post-Remediation Evaluation of EVO Treatment: How Can we improve performance

Introduction

Aquifer pH lower than 6 can reduce the efficiency of enhanced reductive dechlorination (ERD) and other in situ remediation processes. This article describes a simplified approach using a spreadsheet-based design tool for estimating the total amount of base required to achieve a specified target pH at the end of the treatment period, once all reactions have gone to completion. This approach requires simplification of several important processes controlling subsurface pH, and is only appropriate for developing initial estimates of the total amount of base required. In practice, aquifer pH should be monitored during ERD and periodically adjusted to maintain conditions for optimal microbial growth and contaminant degradation. Early in the bioremediation process, reactions will not have gone to completion and less base will be required to maintain the target pH. To prevent overshoot and excessively high pH, users typically add a fraction of the total base required in several increments spread over time.

Background Aquifer pH

In humid areas, rainfall combined with carbonic acid produced in the soil leaches out base cations (Na⁺, K⁺, Ca²⁺, Mg²⁺) gradually acidifying the soil. Figure 1 shows soil pH in the contiguous United States.

(FIGURE 1 HERE) NEED TO REDUCE FILE SIZE

Groundwater pH is influenced by soil pH, but also by other factors. As acidic water infiltrates through the soil profile, some of the acidity may be neutralized by dissolution of soil and aquifer minerals. Silicate minerals including feldspars and micas can hydrolyze over time, consuming acid, and releasing dissolved cations (Na⁺, K⁺, Ca²⁺, Mg²⁺). However, these weathering reactions are slow, and are often not sufficient to prevent pH declines during ERD. Figure 2 shows a cumulative frequency distribution of groundwater pH at a site in eastern North Carolina. Average soil pH at this site is approximately 5. The vadose zone and surficial aquifer at the site are predominantly quartz sand and weathered clays, so weathering processes provide minimal buffering capacity. 90% of the groundwater pH measurements were between 4.5 and 6.0 indicating low soil pH at this site was a reasonable predictor of acidic groundwater. At other sites containing carbonate minerals or relatively young rocks, mineral dissolution often results in greater buffering.

Figure 2. Cumulative frequency distribution of groundwater pH measurements at site in eastern North Carolina where soil pH is approximately 5.

Groundwater Acidity

Acidity is the amount of base required to neutralize acids present in a water sample. Since various acids can disassociate to different extents, two solutions with the same pH can have different acidities (see pH Buffering in Aquifers). In most cases, the acidity of the background groundwater is the sum of acidity from strong mineral acids (phosphoric, nitric, sulfuric, and hydrochloric acids), organic acids, and dissolved carbon dioxide. Mineral acidity (also referred to as methyl orange acidity) is determined by titrating a sample with a strong base to pH=3.7 (Standard Method 2310, AWPA 2016). If the initial pH is greater than 3.7, mineral acidity is zero. The spreadsheet-based design tool presented here uses this measurement to account for the base demand of any strong acids already present in the aquifer. Total acidity (commonly referred to as phenolphthalein acidity) should not be used to calculate the amount of base required because the titration endpoint of this parameter is pH 8.3, which will over-estimate the base requirement for optimal ERD.

Under low pH conditions, organic volatile fatty acids (acetic, propionic, etc.) can accumulate and contribute to acidity. However, in background groundwater, organic acid concentrations are generally low. In the design tool, acidity from background organic acids is assumed to be negligible.

In theory, acidity from dissolved carbon dioxide can be measured using a modification of Standard Method 2310 in which the sample is titrated to a target pH using a strong base. When using this approach, special precautions should be taken to prevent any pH shift due to degassing of dissolved carbon dioxide from the sample during transport to the laboratory or during analysis of the sample. In the design tool, acidity from background dissolved carbon dioxide is instead calculated from the dissolved inorganic carbon concentration (DIC) and background pH. This eliminates the problem of pH shift due to dissolved carbon dioxide degassing. Commercial laboratories can measure DIC using a modification of the Total Organic Carbon (TOC) procedure.

Aquifer Buffering Capacity

An aquifer’s buffering capacity is the resistance to pH change and is primarily due to two processes: (1) buffering by dissolved and solid carbonates; and (2) surface complexation and/or ion exchange reactions on mineral surfaces. While mineral weathering processes do influence long-term pH changes, weathering reactions are often too slow to prevent pH declines due to HCl and CO₂ production during ERD.

Figure 3. Distribution of H₂CO₃*, HCO₃^- and CO₃^2- as a function of pH.

Buffering by Dissolved and Solid Carbonates In natural aqueous systems, pH buffers are predominantly weak acid anions that easily bind and release hydrogen ions. The most common are the weak acid anions produced by dissolved CO₂ . When CO₂ dissolves in water, some CO₂ combines with H₂O forming carbonic acid (H₂CO₃). For convenience, the sum of dissolved CO₂ and H₂CO₃ is often written as H₂CO₃*. H₂CO₃*can then disassociate releasing a bicarbonate ion (HCO₃^-) and one H⁺ or a carbonate ion (CO₃^2-) and two H⁺ by the following:

H₂CO₃* ⇔ H⁺ + HCO₃^- ⇔ 2 H⁺ + CO₃^2-          Reactions 1 and 2

Figure 3 shows the relative distribution of these solutes as a function of pH. The reactions are reversible so that (a) an influx of acid will cause the HCO₃^- and CO₃^2- ions to protonate, consuming the acid, or (b) an influx of base will cause dissociation (deprotonation) of H₂CO₃* and the bicarbonate ion to consume the base. The maximum resistance to pH change (buffering capacity) occurs when the pH is equal to the dissociation constant of either carbonic acid (pH = 6.3 @ 25°C) or of the bicarbonate ion (pH = 10.3 @ 25°C). The buffering capacity of groundwater is measured with an alkalinity titration. The units of alkalinity can be given as milliequivalents (meq) per liter, but are often reported as mg CaCO₃/L (milligrams of CaCO₃ per liter) using the following equation:

Alkalinity (mg CaCO₃/L) = Alkalinity (meq/L) × 50 (mg CaCO₃/meq)                  Equation 1

For a system open to a large reservoir of carbon dioxide (e.g. the atmosphere), the buffering capacity (assuming no soluble minerals are present) is a function of the partial pressure of carbon dioxide in the reservoir. Near the water table, groundwater is open to the atmosphere and can release excess dissolved CO₂ gas if its partial pressure is greater than the reservoir’s, essentially stripping some acid out of the groundwater by partially reversing the above reactions.

In a closed system (e.g. below the water table), the buffering capacity is a function of the total dissolved carbonates (H₂CO₃*, HCO₃^-, and CO₃^-2). When solid carbonate minerals are present (CaCO₃, MgCO₃, CaMg(CO₃)₂, FeCO₃), carbonate mineral dissolution can limit pH declines caused by strong acids (e.g. HCl). For example, in an aquifer containing solid CaCO₃, the ambient groundwater is saturated with CaCO₃ (s). When HCl is produced by ERD, groundwater pH declines, and dissolved CO₃^2- also declines as the carbonate ion is protonated by the added H⁺ ions causing an increase in HCO₃^- and/or H₂CO₃* (see Figure 3). The groundwater is now under-saturated and CaCO₃ will dissolve, consuming H⁺ by the following reaction:

CaCO₃(s) + H⁺ ⇒ Ca²⁺ + HCO₃^-                   Reaction 3

Figure 4. Buffering capacity of aquifer solids from SA-17 and OU-2 measured in DI water. The legend shows depth below ground surface in feet for individual samples.

However, the solubility of carbonate minerals is relatively low, so solid carbonates are much less effective in limiting pH declines due to CO₂ production in the saturated zone. Below the water table, CO₂ may not degas, causing a buildup of HCO₃^-. Geochemical modeling indicates that the aquifer can become supersaturated with CaCO₃ during ERD, causing carbonate dissolution to stop. As additional CO₂ is produced, CaCO₃ (s) will precipitate producing H⁺ by the following reaction:

H₂CO₃* + Ca²⁺ + CO₃^2- ⇒ H⁺ + HCO₃^- + CaCO₃(s)                   Reaction 4

Surface Complexation and Ion Exchange Reactions H⁺ sorption to Fe and Al oxyhydroxides and clay minerals through surface complexation and ion exchange reactions can have a major impact on pH. H⁺ adsorbs strongly to some mineral surfaces, which can accumulate large amounts of H⁺ as the solution pH declines and release the H⁺ back to solution as the pH rises. This strong buffer can reduce the pH decline in many systems, but can also greatly increase the amount of base required to increase aquifer pH.

Figure 5. Aquifer buffering capacity (pHBC) from multiple ERD sites.

Estimating Aquifer Buffering Capacity

When carbonate minerals are present, geochemical models are needed to determine the amount of buffering provided by carbonate mineral dissolution. However, in many naturally low pH aquifers, carbonate minerals are absent and the extent of buffering can be estimated by adding a strong base to the aquifer solids, equilibrating for several days, and measuring the resulting pH. These buffering curves are typically linear in the pH range of 4.5 to 6.5 recommend carrying out the titrations in a uniform ionic strength solution, most commonly 0.01 M CaCl₂.

Figure 4 shows the measured final pH versus meq of OH- added per kg of dry aquifer material. Incremental base addition results in a nearly linear increase in pH. The slopes of these curves were estimated by linear regression. pH buffer capacity (pHBC) is calculated as the inverse of the slope and is reported in units of milliequivalents per kilogram (meq/kg) per pH unit.

Figure 5 shows cumulative frequency distributions for pHBC in 6 different contaminated aquifers in the eastern U.S. Within a single unit, pHBC values are reasonably constant, with the middle 50% of measurements varying by a factor of two. However, median pHBC values can vary by a factor of ten between sites and even between sandy versus clayey units on the same site.

Acid and Base Production

Multiple processes can produce or consume acidity during ERD including redox, precipitation, and hydrolysis reactions. Table 1 shows the amount of H⁺ released from some important redox reactions.

Table 1. Net Acid Production from Important Redox Reactions

e- Acceptor	e- Donor	Product	Reaction	H+ Produced per mole e- donor
PCE	H2	TCE	C2Cl4 + H2 → C2H3Cl3 + H+ + Cl-	1
TCE	H2	cDCE	C2HCl3 + H2 → C2H2Cl2 + H+ + Cl-	1
cDCE	H2	VC	C2H2Cl2 + H2 → C2H3Cl + H+ + Cl-	1
VC	H2	Ethene	C2H3Cl + H2 → C2H4 + H+ + Cl-	1
H2O	Acetic Acid	H2, HCO3-	C2H4O2 + 4 H2O → 2 H2CO3* + 4 H2	2 (1 - α)
H2O	Lactic Acid	H2, HCO3-	C3H6O3 + 3 H2O → 3 H2CO3* + 6 H2	3 (1 - α)
H2O	Glucose	H2, HCO3-	C6H12O6 + 12 H2O → 6 H2CO3* + 12 H2	6 (1 - α)
H2O	Soybean Oil	H2, HCO3-	C56H100O6 + 162 H2O → 56 H2CO3* + 156 H2	56 (1 - α)
Oxygen	H2	H2O	O2 + 2 H2 → 2 H2O	0
Nitrate	H2	N2, OH-	NO3- + 2 ½ H2 → 2 H2O + ½ N2 + OH-	-1
Goethite	H2	Fe2+, OH-	FeO(OH) + ½ H2 → Fe2+ + 2 OH-	-2
Sulfate	H2	HS-	SO42- + 4 H2 + Fe2+ → FeS + 4 H2O	0
H2CO3*	H2	CH4	H2CO3* + 4 H2 → CH4 + 2 H2O	α - 1

During reductive dechlorination of PCE (C₂Cl₄) to TCE (C₂HCl₃), a chlorine atom (Cl) is replaced with hydrogen (H₂) releasing one proton (H⁺) and one chloride (Cl^-).

C₂Cl₄ + H₂ ⇒ C₂HCl₃ + H⁺ + Cl^-                  Reaction 5

As TCE is further dechlorinated to cDCE, VC and ethene, an additional proton is released for each chlorine removed. As a result, ERD can release large amounts of acid. For example, complete reduction of one kilogram of PCE to ethene produces 0.9 kilograms of HCl.

Organic substrates added as electron donors ferment in oxygen-poor environments, releasing H₂ and acetic acid. Acetic acid rarely accumulates during ERD, because it can be used by various cohorts of organisms for reduction of PCE and TCE, for reduction of background electron acceptors, or for fermentation to methane (CH₄). Consumption of acetic acid by all of these processes produces H₂CO₃*. The amount of H₂CO₃* produced during fermentation of different substrates can be estimated with the following formula:

C_θH_βO_γ + (3θ - γ)H₂O ⇒ θH₂CO₃* + (2θ + β/2 - γ)H₂                  Reaction 6

where:
θ	is the number of carbon atoms per mole of substrate
β	is the number of hydrogen atoms per mole of substrate
γ	is the number of oxygen atoms per mole of substrate

Figure 6. Fraction of H₂CO₃* not ionized (α) versus pH.

As shown in Figure 3, CO₃^2- is only present in significant concentrations at pH > 8. For pH ranges relevant for ERD (5>pH>8), H⁺ release from H₂CO₃* can be represented by the following reaction.

H₂CO₃* ⇒ α H₂CO₃* + (1 - α) H⁺ + (1 - α) HCO₃^-                   Reaction 7

where α is the fraction of H₂CO₃* that does NOT ionize at the target pH with α = (1+ 10^-6.352 / [H⁺])^-1. Figure 6 shows the variation in α as a function of pH. At pH 6, α = 0.69, so 0.31 moles of H⁺ are released for every mole of CO₂ produced. However at pH 7, α = 0.18, so 0.82 moles of H⁺ are released for every mole of CO₂ produced.

Reduction of some electron acceptors produces OH^-, counter-acting acid produced from dechlorination and CO₂ production. Representative reactions for reduction of nitrate (NO₃^-), iron oxides, sulfate (SO₄^2-) and bicarbonate (HCO₃^-) are included in Table 1. Goethite [FeO(OH)] is used as a typical iron oxide for these calculations.

Base Addition

A variety of different bases have been used to raise the aquifer pH to stimulate ERD including soluble and solid carbonates, soluble and solid hydroxides, phosphates and silicate minerals. The carbonates and hydroxides are used most commonly because of their relatively low cost and easy availability. Table 2 summarizes the physical properties of common basic salts used to raise pH.

Table 2. Physical Properties of Common Basic Salts used for pH Control

Base	Formula	Molecular Weight (g/mole)	OH- (eq/mole)	OH- (eq/kg Base)	Solubility (g/L)	Saturated solution pH
Caustic Soda	NaOH	40.0	1	25.0	1,100	>14
Caustic Potash	KOH	56.1	1	17.8	1,200	>14
Soda Ash	Na2CO3	106	1 + α	11.2	300	~11.7
Baking Soda	NaHCO3	84	α	2.2	78	~8.3
Hydrated Lime	Ca(OH)2	74.1	2	27.0	1.85	~12.4
Magnesium Hydroxide	Mg(OH)2	58.3	2	34.3	<0.01	~10.3

NaOH and KOH provide a large number of OH- equivalents (eq) per kg and are very soluble so only small volumes of base are required to raise the aquifer pH. However, concentrated solutions of NaOH and KOH have pH > 14 which is inhibitory to bacteria, would expose workers to safety hazards, and can partially dissolve aluminosilicates.

As described above, solid calcium carbonate (CaCO₃) is generally not effective in raising pH during ERD because of its low aqueous solubility. Na₂CO₃ and NaHCO₃ are much more soluble and can be effective in raising aquifer pH. The amount of H⁺ consumed per mole varies as a function of pH between pH 5 and 8. For closed conditions (below water table where CO₂ cannot degas) and pH < 8, NaHCO₃ and Na₂CO₃ disassociate to H₂CO₃* and HCO₃^-, consuming H⁺ by the following reactions.

NaHCO₃ + α H⁺ ⇒ Na⁺ + α H₂CO₃* + (1 – α) HCO₃^-                   Reaction 8

Na₂CO₃ + (1 + α) H⁺ ⇒ 2 Na⁺ + α H₂CO₃* + (1 – α) HCO₃^-                   Reaction 9

At pH = 6, α = 0.69 so 0.69 moles of H⁺ are consumed per mole of NaHCO₃ and 1.69 moles of H⁺ consumed per mole of Na₂CO₃ (Figure 6). However at pH = 7, α = 0.18 so only 0.18 moles of H⁺ are consumed per mole of NaHCO₃ and 1.18 moles of H⁺ consumed per mole of Na2CO₃. As a result, bicarbonates and carbonates are relatively effective at raising the pH to 6. However, these materials provide less alkalinity per unit mass at pH =7, increasing the amount of material required.

Ca(OH) ₂ and Mg(OH) ₂ provide large amounts of OH- per kg. However, these materials have a low aqueous solubility, making them more difficult to distribute in the subsurface. describe the use of a colloidal form of Mg(OH) ₂ with improved transport properties.

Design Tool

An MS Excel based design tool was developed to aid in estimating the amount of base required to achieve a specified target pH at the end of the treatment period, once all reactions have gone to completion. The general approach and calculations were presented in the previous sections. The design tool is specifically focused on pH adjustment for ERD and calculation procedures include the following assumptions.

1. The target pH is between 5 and 8.
2. The carbonate system is the primary aqueous pH buffer.
3. Organic acids including volatile fatty acids (acetic, propionic, etc.) are consumed and do not accumulate.
4. Any H₂ or acetate produced by substrate fermentation that is not consumed through reduction of chlorinated solvents and background electron acceptors (O₂, NO₃, iron oxides, and SO₄) is consumed by methanogenesis.
5. All processes and reactions have gone to completion. The calculated total base requirement is for the end of the treatment period once all acids (HCl and H₂CO₃) have been produced.

In general, the most important factors controlling base requirement are: a) initial and target pH; b) amount of substrate consumed; c) aquifer buffering capacity (pHBC); and d) amount of iron oxides reduced. To calculate base requirement, design tool users must enter the following information.

1. Treatment zone dimensions and design period for this phase of remediation.
2. Site characteristics including average hydraulic conductivity (K), porosity, hydraulic gradient, contaminant concentrations in aquifer material and groundwater, and amount of electron acceptors to be produced (i.e. CH₄) or consumed (O₂, NO₃, Fe, SO₄).
3. Background pH, total inorganic carbon, mineral acidity, and pH buffering capacity (pHBC). A database of pHBC measurements is provided to aid users in selecting design values when laboratory measurements are not available.
4. Mass of organic substrate and base to be injected.
5. When vegetable oil is used as a substrate, the fraction of injected oil that is consumed during the design period.
6. Target pH

Users are reminded that the design tool calculates the total amount of base required once all reactions go to completion. Early in the bioremediation process, reactions will not have gone to completion and less base will be required to maintain the target pH. To prevent overshoot and excessively high pH, users typically add a fraction of the total base required in several increments spread over time.

Figure 7. Required base addition for PRB for different target pH values.

Case Study

The design tool was used to estimate the amount of base required to stimulate ERD in a 435 ft long emulsified vegetable oil (EVO) biobarrier at Operable Unit 2 (OU2) on the former Naval Training Center (NTC) in Orlando, FL. The biobarrier consists of two rows of injection wells spaced approximately 30 ft on center which were treated with a total of 20,224 lb of vegetable oil. The vertical treatment thickness was 10 ft, influent groundwater pH ~5, and pHBC = 8 meq/kg/pH. To provide guidance on amounts of base required, the design tool was used to estimate the amount of NaOH, Na₂CO₃, NaHCO₃, or Mg(OH) ₂ required to achieve different pH values at the end of the five year treatment period. Results presented in Figure 7 indicate that the amount of base required is very sensitive to the target pH. For target pH values < 5.4, little or no base is required. Increasing the target pH to 6 or 7, requires progressively greater amounts of base because of the soil acidity and carbonic acid released from substrate fermentation.

To achieve a pH of approximately 7, 10,200 lb of NaOH or 7,400 lb of Mg(OH) ₂ are required, which is equivalent to 50% or 37% respectively (by weight) of the vegetable oil injected. Multiple NaOH injections would be required since a single injection would result in an excessively high initial pH and the base would migrate out of the treatment zone over time with flowing groundwater. The total mass of Mg(OH) ₂ required is less than NaOH, since two moles of OH- are released per mole of Mg(OH) _2. However, Mg(OH) ₂ has a very low aqueous solubility, so the material must be injected in a colloidal form. Since the material is a solid and would not migrate downgradient, 100% of the required Mg(OH) ₂ could be injected simultaneously with the EVO. The pH of a pure slurry of Mg(OH)₂ is relatively high (~10.3). However once injected, CO₃^2- precipitates on the surface of the Mg(OH) ₂ particles forming a MgCO₃ coating that maintains the aquifer pH between 7 and 8. In general, it is not practical to raise the pH to near 7 with NaHCO₃ due to the small amount of H⁺ consumed by this material at near neutral pH. Na₂CO₃ could be effective, but would require multiple injections due to the high pH (~11.7) and downgradient migration with groundwater flow.

Summary

Alkaline materials are commonly added during ERD to maintain an aquifer pH of greater than 6. A simplified approach and spreadsheet-based design tool are presented for estimating the total amount of base required to achieve a specified target pH at the end of the treatment period, once all reactions have gone to completion. To prevent overshoot and excessively high pH, users typically add a fraction of the total base required in several increments spread over time. Users should be aware that results are very sensitive to variations in the target pH since the amount of H⁺ released from carbon dioxide and consumed by NaHCO₃ and Na₂CO₃ is a function of pH.

See Also

• BUCHLORAC: Program for predicting the bicarbonate requirement for anaerobic bioremediation of chlorinated solvents