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. 2023 Jan 31;57(47):18433–18442. doi: 10.1021/acs.est.2c05324

Inhibition of Hexavalent Chromium Release from Drinking Water Distribution Systems: Effects of Water Chemistry-Based Corrosion Control Strategies

Cheng Tan 1, Haizhou Liu 1,*
PMCID: PMC10690716  PMID: 36719710

Abstract

graphic file with name es2c05324_0008.jpg

In drinking water distribution systems, the oxidation of zerovalent chromium, Cr(0), in iron corrosion scales by chlorine residual disinfectant is the dominant reaction to form carcinogenic hexavalent chromium, Cr(VI). This study investigates inhibitive corrosion control strategies through adjustments of chemical water parameters (i.e., pH, silicate, phosphate, calcium, and alkalinity) on Cr(VI) formation through oxidation of Cr(0)(s) by free chlorine under drinking water conditions. The results show that an increase in pH, silicate, alkalinity, and calcium suppressed Cr(VI) formation that was mainly attributed to in situ surface precipitation of new Cr(III) solids on the surface of Cr(0)(s), including Cr(OH)3(s), Cr2(SiO3)3(s), CrPO4(s), Cr2(CO3)3(s), and Cr10Ca(CO3)16(s). The Cr(III) surface precipitates were much less reactive with chlorine than Cr(0)(s) and suppressed the Cr redox reactivity. The concentration of surface Cr(III) solids was inversely correlated with the rate constant of Cr(VI) formation. Adding phosphate either promoted or inhibited the Cr(VI) formation, depending on the phosphate concentration. This study provides fundamental insight into the Cr(VI) formation mechanisms via Cr(0) oxidation by chlorine and the importance of surface precipitation of Cr(III) solids with different corrosion control strategies and suggests that increasing the pH/alkalinity and addition of phosphate or silicate can be effective control strategies to minimize Cr(VI) formation.

Keywords: chromium, carcinogen, free chlorine, corrosion, redox reactivity, drinking water

Short abstract

This study reports practical methods to reduce the carcinogenic hexavalent chromium release in drinking water distribution systems.

Introduction

Chromium exists in one of three stable oxidation states in the environment, including zerovalent chromium, Cr(0), trivalent chromium, Cr(III), and hexavalent chromium, Cr(VI).1 In drinking water with a pH above 6.5, Cr(VI) in the form of carcinogenic CrO42– is the dominant dissolved Cr species (chromic acid: pKa1 = 0.7 and pKa2 = 6.5).2−4 The current maximum contaminant level (MCL) for total Cr established by the U.S. EPA is 100 μg/L.5 A Public Health Goal (PHG) for Cr(VI) was set at 0.02 μg/L by the Office of Environmental Health Hazard Assessment in California.6 The Cr(VI) concentration in U.S. drinking water was measured to be between 0.057 and 7.51 μg/L based on results from the 2018–2020 U.S. EPA Round 4 Unregulated Contaminant Monitoring Rule Program (UCMR4).7

The sources of Cr(VI) in drinking water include anthropogenic industrial waste discharges and natural geological weathering of Cr-containing minerals.8−11 Cr(VI) can be effectively removed from drinking water by reductive coagulation and filtration, ion exchange, adsorption, or membrane filtration.12,13 However, Cr(0) in its metallic form is commonly used in iron-based corrosion-resistant alloys such as cast iron and steel materials.8 Cast iron pipes (1–4% Cr(0)) and steel pipes (≥11% Cr(0)) together account for approximately 60% of all piping materials in drinking water distribution systems (DWDSs) in the United States.14 Reactions between disinfectants and aging iron pipes can lead to the release of Cr(VI) in drinking water.15 Analysis of the UCMR4 database also showed that more than 40% of the drinking water that left distribution systems had increased in Cr(VI) concentration relative to the entry point.16

Cr(0) solid was recently discovered to exist in iron corrosion scales and significantly contributed to Cr(VI) formation by oxidative conversion to Cr(VI) by the residual free chlorine (denoted HOCl as the dominant species at pH 7 in the following text) in DWDSs.15 The Cr(0) in the cast iron pipe material exhibited approximately 2 orders of magnitude higher reactivity with HOCl than Cr(III) hydroxide-based solids in the iron corrosion scales.15

Little knowledge exists on how to inhibit the Cr(0) redox reactivity to minimize Cr(VI) formation. Traditionally, adjustment of chemical parameters has been employed or proposed as an effective corrosion strategy to minimize metal release from DWDSs. For example, pH and alkalinity adjustment have been used for iron corrosion control, which can be accomplished by aeration and chemical dosing, including lime Ca(OH)2 and soda ash Na2CO3.17,18 Raising the pH to 8 was reported to effectively control iron, lead, and copper corrosion,19−21 as a higher pH can favor the formation of a passivation layer.22 Water pH can significantly impact the speciation of surface hydroxo complexes of Cr(III) solids and the speciation of chlorine that in turn can affect the reaction kinetics of Cr(VI) formation through the oxidation of Cr(III) solids by HOCl.23,24 Alkalinity also can significantly influence the corrosion of metallic pipes and release different metals including iron and lead in DWDSs.19,25,26 High calcium and alkalinity can lead to calcium carbonate precipitation that acts as a passivation layer to inhibit lead and copper corrosion.27,28 However, the effectiveness of alkalinity and calcium on the control of chromium release is unknown.

Furthermore, phosphate and silicate can be used as corrosion inhibitors and sequestrants.29,30 Adding phosphate effectively controls lead corrosion by forming Pb phosphate solids that have an extremely low solubility.31 Silicate has been shown to reduce lead and copper levels and may form an adherent film on the surface of the pipe that acts as a diffusion barrier.32 Silicate has also been used as corrosion inhibitor in iron-based water pipes, as silicate can complex Fe2+ and form a passivation layer of FeSiO3(s).33

However, drinking distribution systems are highly complex; thus, a particular control strategy may be system specific. For example, it was reported that both weight loss and iron concentration decreased on increasing the pH from 8.5 to 9.2, and red water was received when the alkalinity was maintained at higher values.34 Higher levels of silica can cause more iron release to the water and decreased the size of suspended iron particles.35

Currently, very little is known on how significant the adjustment of these water chemistry parameters as corrosion control strategies can affect the surface redox reaction between Cr(0) solid and HOCl. Filling this knowledge gap is critical to understand and control Cr(VI) formation in water from DWDSs. Accordingly, the objectives of this study were (1) to investigate the reaction kinetics and mechanisms of Cr(0)(s) oxidation by chlorine in conjunction with different corrosion control strategies based on adjustments of chemical water parameters (i.e., pH, silicate, phosphate, calcium, and alkalinity) and (2) to explore new control strategies to inhibit Cr(VI) release from DWDSs. A thorough understanding of the influence of these water parameters on Cr(VI) formation is essential to develop Cr(VI) control strategies in different drinking water distribution infrastructures.

Materials and Methods

Chemicals

All chemicals used in this study were reagent grade or higher. All solutions were freshly prepared using deionized water (>18.2 MΩ·cm, Millipore System). Cr(0)(s) with a uniform particle size of 10 μm was purchased from Sigma-Aldrich. The chlorine working solution (HOCl) was prepared freshly using a NaOCl stock solution. Four Cr(III) reference solids, including CrPO4(s), Cr2(SiO3)3(s), Cr2(CO3)3(s), and Cr10Ca(CO3)16(s), were synthesized by a precipitation method. The detailed synthesis procedure and characterization information is given in Text S1 of the Supporting Information (SI).

Oxidation Experiments of Cr(0) Solid by Chlorine

A 200 mL suspension containing 440 mg/L (8.5 mM) Cr(0)(s) was mixed with 20 mg of Cl2/L (0.28 mM) of HOCl in a 250 mL glass vessel in the dark. HOCl at 20 mg Cl2/L was used instead of the lower concentration in DWDSs (<2 mg/L Cl2/L) because the higher concentration provides valuable insight into the oxidation kinetics, while the kinetics and mechanisms obtained at high-concentration conditions can be extrapolated to predict reactions at low-concentration conditions.15,36,37 The oxidation experiment was conducted for 6 h. The chemical parameters were varied to reflect different drinking water conditions, including pH (6–8), silicate (0–10 mg/L), phosphate (0–10 mg/L), calcium (0–300 mg/L), and alkalinity (0–200 mg CaCO3/L). The pH of the suspension was maintained throughout the reaction with a Eutech Instrument Alpha pH200 controller. The concentrations of silicate, phosphate, calcium, and alkalinity were prepared by adding Na2SiO3, Na3PO4, CaCl2, and NaHCO3 stock solutions, respectively. After HOCl is dosed into the reactor, 5 mL of the suspension was collected every 30 min and filtrated through a 0.22 μm filter (nylon membrane, Agilent). The HOCl in the filtrate was immediately quenched with excess (NH4)2SO4, and the total amounts of Cr and Cr(VI) were measured. This followed an established protocol for sample pretreatment.15 An additional 0.5 mL of suspension was collected every 2 h, filtered through 0.22 μm filters, diluted by 10 times, and quantified for HOCl concentration. In addition to kinetic studies with Cr(0), the oxidation of four Cr(III) reference solids by HOCl was conducted at pH 7 as a control.

Acid Digestion and Solubility Experiments of Cr(III) Solid

To quantify the amount of in situ surface precipitation of Cr(III) solids during Cr(0)(s) oxidation by HOCl, solids were collected and acid digested by EPA Method 3050B after the 6 oxidation reaction.38 The brief principal was that HNO3 and 30% H2O2 can dissolve Cr(III) solids with heating at 95 °C. In brief, after a 6 h reaction, the suspension was quenched with excess (NH4)2SO4 and centrifuged for 10 min. The Cr solids were washed twice with DI water and centrifuged to efficiently remove the adsorbed ions. The final Cr solids were freeze dried and digested. In addition, the Cr(0)(s) from a control experiment without free chlorine was digested to deduct the background of any residual of adsorbed ions. This was proved to be a reliable method to quantify Cr(III) solids without inference from Cr(0)(s) because EPA Method 3050B can completely dissolve Cr(III) solids and cannot dissolve Cr(0) solid.15 The concentrations of Cr, P, and Si in the digested solution were analyzed using ICP-MS to quantify the Cr(III) solid species on the surface of Cr(0)(s).

In addition, control experiments to evaluate the adsorption of Cr(VI) onto Cr(0)(s) were conducted by dosing 0.1–1 mg/L Cr(VI) using potassium dichromate (K2Cr2O7) in a 0.44 g/L Cr(0) suspension. After 6 h, the aqueous Cr(VI) concentration was quantified by Method 3050B. On the basis of the mass balance, the percentage of Cr(VI) adsorbed on Cr(0)(s) was calculated. The adsorption of Cr(VI) on Cr(0)(s) was negligible after 6 h adsorption at pH 6–8 (Figure S1); therefore, Cr(VI) adsorption did not interfere with the quantification of precipitated Cr(III) solids on the surface of Cr(0)(s).

To further understand the Cr(III) solid formation under different water chemistry, the solubility product constant (Ksp) for each synthesized Cr(III) reference solid was quantified based on a 6 h dissolution experiment. The dissolution time was chosen as 6 h because the dissolved Cr(III) concentrations reached plateaus at 6 h as shown in Figure S2. Briefly, the suspension with 0.1 g/L of each Cr(III) reference solid was stirred at 700 rpm for 6 h, and then, the dissolved Cr that can pass through a 0.22 μm filter (nylon membrane, Agilent) was quantified and used to calculate the Ksp (detailed information in Text S2). All experiments were conducted at 25 °C.

Analytical Methods

Cr(VI) concentration was measured using the standard diphenylcarbazide (DPC) method with a spectrophotometer.39 The brief principal was that DPC reacted with Cr(VI) to form the colored Cr(VI)–DPC complex, and the absorbance of the solution was then measured at 540 nm. HOCl concentrations were measured using the standard DPD method.39 The concentrations of dissolved Cr, Si, P, and Ca were analyzed using inductively coupled plasma mass spectrometry (ICP-MS, Agilent 7700). Ca was measured under the hydrogen mode to reduce interference.40 The Brunauer–Emmett–Teller (BET) surface areas of Cr(0) and Cr(III) solids were measured using a Micromeritics ASAP 2020 surface area analyzer. All experiments were conducted in duplicate, and the error bars were calculated accordingly. Two-way ANOVA statistical analysis was conducted using Software R (v4.0.3) to test the significant difference for the effects of pH and silicate on Cr(VI) formation.

Modeling of Reaction Kinetics

The reaction rate constant of Cr(VI) formation (kCr(VI)) via Cr(0)(s) oxidation by chlorine was obtained by fitting experimental data of Cr(VI) release using the software OpenModel. A second-order reaction kinetics model was introduced to quantify the reaction rate constant of Cr(VI) formation (kCr(VI)) based on eq 1

graphic file with name es2c05324_m001.jpg 1

where kCr(VI) is the surface-area normalized second-order rate constant for Cr(VI) formation from Cr(0) powder (L·m–2.min–1), [HOCl] is the HOCl concentration (mol/L) at any reaction time t (min), [Cr(s)] is the Cr solids concentration (mol/L), SCr(0) is the Cr(0)(s) BET surface area (m2/g), and AWCr is Cr atomic weight of 52 g/mol. [Cr(s)] is the initial Cr(0)(s) concentration minus the dissolved Cr(VI) concentration at time t.

All of the data for the Cr(VI) concentrations and free chlorine concentrations at different time points were input into the model. kCr(VI) was an experimentally observed rate constant combining both the one-step oxidation of Cr(0)(s) to Cr(VI) and the oxidation of Cr(III)(s) to Cr(VI). All modeled fitting curves matched the observed Cr(VI) concentrations well with R2 ranging between 0.972 and 0.996.

To evaluate the control strategies for Cr(VI) release in DWDSs dominated by iron pipe, a kinetics model was established using the reaction rate constants (kCr(VI)) obtained from this study (with details of the model found in Text S3). The relative Cr(VI) formation was predicted based on typical drinking water conditions, and it was calculated by dividing the Cr(VI) formation under one certain water condition by Cr(VI) formation at pH 6 with only Cr solid and free chlorine.

Results and Discussion

Effect of pH on Cr(VI) Formation

The redox potential of Cr(VI)/Cr(0)(s) (−1.49 V) is lower than that of HOCl/Cl– (1.25 V) in typical drinking water chemical conditions;15 thus, the overall oxidation of Cr(0)(s) to Cr(VI) by HOCl is thermodynamically feasible as

graphic file with name es2c05324_m002.jpg R1

Cr(VI) formation decreased with increasing pH from 6 to 8 (Figure 1A). Statistical analysis showed that compared to the Cr(VI) formation at pH 7, all pH values had significant effects on Cr(VI) formation except for pH 6.5. Correspondingly, surface-area-normalized rate constants of Cr(VI) formation (kCr(VI)) decreased by 31% from 1.9 × 10–3 to 1.3 × 10–3 L·m–2.min–1 (Figure 1B). During the oxidation reaction, the Cr(III) solid concentration on the surface of Cr(0)(s) increased from 0.17 to 0.47 mg/g when the solution pH increased from 6 to 8 (Figure S3); thus, a higher pH resulted in an enhanced Cr(III) solid formation on the surface of Cr(0)(s). CO2 in air has a negligible effect on the Cr(III)(s) formation. At pH 7, the carbonate in equilibrium with headspace air in the reactor resulted in an negligible amount of Cr(III)(s) formation during the oxidation reaction. In addition, the 250 mL glass vessels were covered by parafilm to avoid contact to air. Because OH– is the only ligand to complex dissolved Cr(III), the precipitated Cr(III) solid is expected to exist as Cr(OH)3(s) with a maximum dissolved Cr(III) concentration of 0.1 μg/L at pH 7 based on the solubility product constant of Cr(OH)3(s) (Ksp = 6.3 × 10–31).41 Furthermore, control experiments with pure Cr(OH)3(s) solid exhibited a much lower rate constant for Cr(VI) formation via HOCl oxidation (i.e., 0.2 × 10–3 vs 1.67 × 10–3 L·m–2.min–1; Figure 2) than Cr(0)(s). Therefore, the enhanced formation of Cr(OH)3(s) on the surface of Cr(0)(s) at higher pH inhibited Cr(VI) formation.

Figure 1.

Figure 1

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Six hour oxidation of Cr(0)(s) by free chlorine at different pH levels. [HOCl]TOT = 20 mg Cl2/L (0.282 mM); [Cr(0)(s)] = 440 mg/L (8.46 mM); molar ratio of Cr(0)/HOCl = 30:1; ionic strength = 10 mM, [Ca] = 0 mg/L. (A) Cr(VI) formation; (B) kCr(VI). p value for Cr(VI) formation compared to that at pH 7: pH 6 (0.0105), pH 6.5 (0.9890), pH 7.5 (0.0066), and pH 8 (<0.0001).

Figure 2.

Figure 2

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kCr(VI) from the oxidation of different Cr solids by free chlorine at pH = 7 with 10 mM ionic strength. [HOCl]TOT = 20 mg Cl2/L (0.282 mM); Cr solids = 440 mg/L.

Effect of Phosphate on Cr(VI) Formation

The impact of phosphate on Cr(VI) formation depended on the applied concentration (Figure 3A). When the phosphate dosage was increased from 0 to 0.1 mg/L as PO4, the rate of Cr(VI) formation increased by 53% (1.7 × 10–3 vs 2.6 × 10–3 L·m–2.min–1; Figure 3B). The presence of phosphate was previously reported to block the active sites for malachite Cu2CO3 (OH)2(s) precipitation during the copper corrosion process, resulting in an increase in copper release.42 Similarly, in the system with Cr solids, a low concentration of phosphate suppressed Cr(OH)3(S) formation on the surface of Cr(0)(s) because phosphate can compete with OH– for adsorption sites on the surface Cr(0)(s), consequently decrease the surface concentration of OH–, and suppress Cr(OH)3(s) formation. This surface reaction mechanism was supported by the observed Cr(OH)3(s) formation during the oxidation reaction. In the absence of phosphate, 0.38 mg/g Cr(OH)3(s) was formed on the surface of Cr(0)(s) after 6 h oxidation (Figure S4). When 0.1 mg/L phosphate was added to the system, Cr(OH)3(s) formation decreased by 53% (0.18 vs 0.38 mg/g). Therefore, 0.1 mg/L phosphate enhanced Cr(VI) formation by decreasing Cr(OH)3(s) formation on the surface of Cr(0)(s). It was found that 0.5 mg/L phosphate had lower Cr(VI) formation than 0.1 mg/L phosphate, which was possibly due to Cr(PO4)(s) beginning to form on the surface of Cr(0)(s). In addition, approximately 70% of phosphate was transferred from the aqueous phase to the solid phase at the end of the Cr(0)(s) oxidation reaction (Figure S5), which indicates the adsorption of phosphate during the reaction. There is little in the literature on how other aqueous species affected the OH– adsorption. Previous studies showed that phosphate can decrease the adsorption of negatively charged species such as carbonate species, silicate, and citrate through competing for the adsorption sites.43−45 Therefore, it is possible that the dominant anions in drinking water such as HCO3–, Cl–, SO42–, and NO3– may reduce OH– adsorption due to the competitive adsorption.

Figure 3.

Figure 3

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Six hour oxidation of Cr(0)(s) by free chlorine at different phosphate levels. [HOCl]TOT = 20 mg Cl2/L (0.282 mM); [Cr(0)(s)] = 440 mg/L (8.46 mM); molar ratio of Cr(0)/HOCl = 30:1; pH = 7; ionic strength = 10 mM, [Ca] = 0 mg/L. (A) Cr(VI) formation; (B) kCr(VI).

In contrast, when the phosphate concentration was ≥0.75 mg/L, Cr(VI) formation was inhibited. kCr(VI) was between 0.8 and 1.1 × 10–3 L·m–2.min–1 with 0.75–10 mg/L phosphate (Figure 3B), 35–53% lower than that in the absence of phosphate. Accordingly, the concentration of Cr(III) solids on the surface of Cr(0)(s) was 28–38% higher than that without phosphate (Figure S6). Phosphate can react with dissolved Cr(III) to form insoluble CrPO4(s) with a Ksp of 7.2 × 10–18 (Table S1) (reaction 2)

graphic file with name es2c05324_m003.jpg R2

The Ksp was calculated based on the 6 h of dissolution experiments to represent newly formed Cr(III) solids during the 6 h oxidation experiment. The Ksp of CrPO4(s) (7.2 × 10–18) was much higher than that in the previous study (2.4 × 10–23),46 possibly due to the formation of a more soluble form of CrPO4(s) with a shorter duration of dissolution during the synthesis in this study.

Because CrPO4(s) solid exhibited a much lower reactivity than Cr(0)(s) to form Cr(VI) (0.03 × 10–3 vs 1.67 × 10–3 L·m–2.min–1) (Figure 2), formation of CrPO4(s) can inhibit Cr(VI) formation. The acid digestion result confirmed the formation of CrPO4(s) on the surface of Cr(0)(s). When 1 mg/L phosphate was added to the system, both Cr(OH)3(s) (0.31 mg/g as Cr(III)/Cr(0)) and CrPO4(s) (0.17 mg/g as Cr(III)/Cr(0)) were formed on the surface of Cr(0)(s) (Figure S4). In addition, although 10 mg/L phosphate could have more CrPO4(s) precipitation than 5 mg/L phosphate, some of the CrPO4(s) may detach from the Cr(0)(s), resulting in a comparable Cr(0)(s) surface area that is exposed to free chlorine. Thus, a similar inhibitory effect was observed at 5 and 10 mg/L phosphate.

Many studies have shown that the adsorption capacity of phosphate decreased with increasing pH. Three different studies showed that the adsorption of phosphate on different metal oxides decreased by approximate 25% when increasing two units of pH: goethite with pH from 6 to 8,47 hydrous ferric oxide with pH from 6.5 to 8.5 (40 vs 32 mg/g),48 and zirconium-modified coal gasification coarse slag with pH from 6 to 8 (16 vs 12 mg/g).49 Assuming that the phosphate adsorption on Cr(0)(s) was affected to the same extent, it is estimated that 0.44 mg/L phosphate addition at pH 6 and 0.57 mg/L phosphate addition at pH 8 would enhance Cr(VI) formation. Thus, the critical phosphate level was approximately 0.5 mg/L across the typical pH range (6–8) in drinking water. However, the critical value should be further experimentally investigated due to the higher complexity of the drinking water distribution systems.

Effect of Silicate on Cr(VI) Formation

Statistical analysis showed that silicate (SiO32–) had no significant effect on Cr(VI) formation at a concentration of ≤2.5 mg/L but exhibited an inhibitive effect on Cr(VI) formation at a concentration of ≥5 mg/L (Figure 4A). The rate constant with 10 mg/L silicate was 24% lower than without (1.3 × 10–3 vs. 1.7 × 10–3 L·m–2.min–1) (Figure 4B). Correspondingly, when the silicate concentration was ≥2.5 mg/L, the solid Cr(III) concentration on the surface of Cr(0)(s) was 66–97% higher than that without (0.63–0.75 vs 0.38 mg/g) (Figure S7). Silicate can react with dissolved Cr(III) to form insoluble Cr2(SiO3)3(s) with a Ksp of 9.8 × 10–36 (Table S1) (reaction 3)

graphic file with name es2c05324_m004.jpg R3

Figure 4.

Figure 4

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Six hour oxidation of Cr(0)(s) by free chlorine at different silicate levels. [HOCl]TOT = 20 mg Cl2/L (0.282 mM); [Cr(0)(s)] = 440 mg/L (8.46 mM); molar ratio of Cr(0)/HOCl = 30:1; pH = 7; ionic strength = 10 mM, [Ca] = 0 mg/L. (A) Cr(VI) formation and (B) kCr(VI). p value for Cr(VI) formation compared to that without silicate: 1 (0.9779), 2.5 (1.000), 5 (0.0100), and 10 mg/L (<0.0001).

When 0.5 mg/L silicate was added to the system, only Cr(OH)3(s) (0.35 mg/g as Cr(III)/Cr(0)) was formed on the surface of the Cr(0)(s), which was comparable to the conditions without silicate (0.38 mg/g as Cr(III)/Cr(0)) (Figure S4). In the presence of 10 mg/L silicate, both Cr2(SiO3)3(s) (0.39 mg/g) and Cr(OH)3(s) (0.35 mg/g) were formed on the surface of Cr(0)(s) (Figure S4). Because Cr2(SiO3)3(s) had a much lower reactivity than the Cr(0)(s) reaction to form Cr(VI) (0.018 × 10–3 vs 1.67 × 10–3 L·m–2.min–1) (Figure 2), silicate can inhibit Cr(VI) formation through formation of Cr2(SiO3)3(s). The stoichiometry of the solid was confirmed. The Cr and Si concentrations in the digested sample were measured by ICP-MS, and the molar ratio of Cr to Si was close to 2:3.

Effects of Alkalinity and Calcium on Cr(VI) Formation

Increasing alkalinity to 50 mg CaCO3/L can suppress Cr(VI) formation (Figure 5A and Figure S8). kCr(VI) with 50 mg CaCO3/L alkalinity was 24% lower than that without alkalinity (Figure 5B). Alkalinity promotes the formation of insoluble Cr2(CO3)3(s) with a Ksp of 1.1 × 10–45 (Table S1) (reaction R4)

graphic file with name es2c05324_m005.jpg R4

Figure 5.

Figure 5

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Six hour oxidation of Cr(0)(s) by free chlorine at different alkalinity levels (0–200 mg/L as CaCO3) with and without 200 mg/L Ca. [HOCl]TOT = 20 mg Cl2/L (0.282 mM); [Cr(0)(s)] = 440 mg/L (8.46 mM); molar ratio of Cr(0)/HOCl = 30:1; pH = 7; ionic strength = 10 mM: (A) Cr(VI) formation and (B) kCr(VI).

The equilibrium dissolved Cr(III) factor (α) was introduced to confirm the formation of Cr2(CO3)3(s). The value of α was calculated as the ratio of the dissolved Cr(III) concentration in equilibrium with a Cr(III) solid to the dissolved Cr(III) concentration in equilibrium with Cr(OH)3(s) at pH 7 as shown in eq 2

graphic file with name es2c05324_m006.jpg 2

where Cr(III)dis1 is the dissolved Cr(III) concentration in equilibrium with different Cr(III) solids at pH 7 and Cr(III)dis0 is the dissolved Cr(III) concentration in equilibrium with Cr(OH)3(s) at pH 7. The value of α indicated the sequence of the Cr(III) precipitation. When α > 1, Cr(OH)3(s) will precipitate as the dominant species. When α < 1, other Cr(III) solids will precipitate as dominant species.

The value of alpha predicts the dominant Cr(III) solids based on the equilibrated Cr(III) concentrations under different water chemistry from a thermodynamic point of view with the assumption that kinetics did not play an important role in the precipitation of different Cr(III)(s). The data showed that Cr2(CO3)3(s) has an α value significantly lower than 1 (Figure S9); therefore, Cr2(CO3)3(s) was the dominant Cr(III) solid on the surface of Cr(0) when an alkalinity of 10–200 mg CaCO3/L was added into the system. Because Cr2(CO3)3(s) has a much lower reactivity than Cr(0) to form Cr(VI) (0.12 × 10–3 vs 1.67 × 10–3 L·m–2.min–1; Figure 2), formation of Cr2(CO3)3(s) inhibited Cr(VI) formation.

Under the same alkalinity condition, adding 200 mg/L calcium decreased Cr(VI) formation further by 7–19%, exhibited by lower kCr(VI) values (Figure 5B). Thus, calcium and alkalinity together had a synergistic inhibitive effect on Cr(VI) formation by forming Cr10Ca(CO3)16(s) with a Ksp of 1 × 10–2 50 (Table S1) (Reaction R5)

graphic file with name es2c05324_m007.jpg R5

The value of α for Cr10Ca(CO3)16(s) in the presence of 200 mg/L calcium was significantly lower than that without (3.5 × 10–5–2.9 × 10–7 vs 1.2 × 10–3–1.4 × 10–5) (Figure S9); thus, Cr10Ca(CO3)16(s) would precipitate prior to Cr2(CO3)3(s). The acid digestion confirmed that Cr10Ca(CO3)16(s) was formed on the surface of Cr(0)(s) with 200 mg/L Ca and 200 mg/L alkalinity (Figure S4). Because Cr10Ca(CO3)16(s) has a much lower reactivity than the Cr(0)(s) reaction to form Cr(VI) (3.2 × 10–5 vs 1.67 × 10–3 L·m–2.min–1) (Figure 2), formation of Cr10Ca(CO3)16(s) further inhibited Cr(VI) formation. In addition, without alkalinity, calcium alone (0–300 mg/L Ca2+) had no effect on Cr(VI) formation through oxidation of Cr(0)(s) by chlorine (Figure S10) because calcium alone does not induce the formation of additional surface solid (Figure S4). The dissolution time was chosen as 6 h because the dissolved Cr(III) concentrations reached plateaus at 6 h as shown in Figure S2.

In all of the experiments, the Cr(VI) formation rates were constant throughout the reactions (Figures 1A, 3A, 4A, and 5A). When Cr(0)(s) reacts with free chlorine, Cr(III)(s) can form on the surface of Cr(0)(s), and the Cr(III)(s) can further react with free chlorine to form Cr(VI). The production and consumption of Cr(III) solids was a dynamic process reaching steady state in a short time, resulting in a constant Cr(III) steady-state concentration on the surface of Cr(0)(s). Thus, the inhibitive effect of Cr(III)(s) was constant during a majority of the reaction time, leading to a constant Cr(VI) formation rate.

Consumption of Free Chlorine

Free chlorine had a pK of 7.6; thus, HOCl was the dominant species when pH was lower than 7.6 and OCl– was the dominant species when pH was higher than 7.6. HOCl can change to Cl2O with an equilibrium constant (K) of 8.74 × 10–3 M–1 at 25 °C.50 In equilibrium at pH 6, although the Cl2O concentration was approximately 6 orders of magnitude lower than the HOCl concentration (6.5 × 10–10 vs 2.73 × 10–4 M), Cl2O may still play a significant role for the oxidative formation of Cr(VI) due to its high reaction activity.51 In this study, Cr(VI) formation was due to oxidation of all of the chlorine species, and their relative contributions were not further investigated.

For all reactions, less than 20% of the free chlorine were consumed after the 6 h of reaction. For example, 1.7–2.5 mg/L (8.5–12.5%) free chlorine was consumed after the 6 h reaction at different pHs (Figure S11). For the consumed free chlorine, 54–63% was due to oxidation of Cr(0)(s) to form Cr(VI), and the rest was mainly due to oxidation of Cr(0)(s) to form Cr(III)(s) on the surface of Cr(0)(s) and the possible autocatalytic decay of chlorine. The oxidation of Cr(III)(s) by chlorine can generate Cr(V) intermediate species in the form of Cr(V) oxide, and the Cr(V) oxide can generate dissolved oxygen by quickly decomposing to Cr(III)(s).36

Mechanisms of the Inhibition on Cr(VI) Formation

An increase in pH was previously reported to modestly enhance Cr(VI) formation during oxidation of Cr(III)(s) solids by free chlorine due to changes of reactive Cr(III)(s) surface hydroxo species.36 In contrast, Cr(VI) formation was observed to decrease with increasing pH during Cr(0)(s) oxidation by free chlorine. The opposite pH trend strongly suggests that oxidation of Cr(III)(s) was not the rate-limiting reaction for Cr(VI) formation in the system with Cr(0)(s) and free chlorine. In addition, both Cr(0)(s) and Cr(III)(s) coexisted in the naturally formed iron corrosion scales, and Cr(0)(s) was found to mainly contribute to the fast Cr(VI) formation when the iron corrosion scales reacted with free chlorine.15 As a result, the one-step reaction of direct Cr(0)(s) oxidation to Cr(VI) by free chlorine is likely the dominant pathway for Cr(VI) formation.

In this study, Cr(0)(s) exhibited a Cr(VI) formation rate constant that was 1–2 orders of magnitude higher than that for Cr(III) solids (Figure 2); thus, it is possible that a fast one-step Cr(0)(s) oxidation to Cr(VI) was the dominant reaction to form Cr(VI) (Mechanism 1). In addition, the surface speciation measurements of Cr(III) solid (Figures S3, S6, and S7) confirmed that different less reactive Cr(III) solids were formed during oxidation of Cr(0)(s) by free chlorine when pH, silicate, phosphate, calcium, and alkalinity were adjusted to desirable levels. This observation supported the second Cr(VI) formation mechanism (Mechanism 2) that Cr(VI) formation via Cr(0)(s) oxidation involved two steps: Cr(0)(s) oxidation to Cr(III)(s) and then Cr(III)(s) oxidation to Cr(VI) (kIII).

To quantitatively evaluate the relative importance of Mechanisms 1 and 2 on Cr(VI) formation, the branching ratio of Mechanism 1 was calculated (eq 3)

graphic file with name es2c05324_m008.jpg 3

where SCr(0) and SCr(III) are the percentages of Cr(0)(s) surface area covered by Cr(0)(s) and Cr(III)(s) and k0 and kIII are the surface-area-normalized second-order rate constants for Cr(VI) formation from Cr(0) and Cr(III) solids. The details of the calculation of SCr(0) and SCr(III) and the branching ratio are shown in Text S4, and the results are shown in Figures S12 and S13, respectively. The branching ratio of Mechanism 1 (one-step reaction) was overwhelmingly higher (i.e., 0.95–0.98 in most conditions; Figure S13). Overall, these results confirmed that the dominant mechanism for Cr(VI) formation was the one-step oxidation of Cr(0)(s) to Cr(VI). In addition, previous studies showed that Cr(V), as a Cr intermediate species, could be formed during the redox reaction.52,53 However, the Cr intermediate species was not investigated in this study. The proposed one-step oxidation mechanism does not rule out the possibility that a Cr intermediate species could be formed and was quickly oxidized to Cr(VI), which may contribute a significant amount of Cr(VI).

Although the Cr(III) solids had different reactivities with HOCl, kCr(VI) apparently has an inverse relationship to the Cr(III) solid concentration on the surface of Cr(0)(s) (Figure 6). On the basis of their low branching ratios, Cr(III) solids were not the dominant source of Cr(VI) formation (Figure S13), but they can significantly inhibit Cr(VI) formation through reducing the contact area between Cr(0)(s) and chlorine, decreasing Cr(VI) formation. For example, SCr(0) decreased from 75% to 35% when 200 mg/L Ca2+ was added, and the alkalinity was increased from 0 to 200 mg CaCO3/L (Figure S12). Overall, Cr(III) solids on the surface of Cr(0)(s) are critical for Cr(VI) formation through oxidation of Cr(0)(s) by chlorine; thus, the control strategy for Cr(VI) formation should focus on the Cr(III) solids formation on the surface of Cr(0)(s).

Figure 6.

Figure 6

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Relationship between the kCr(VI) and the Cr(III) solids using all available data of Cr(III) solid formation on the surface of Cr(0)(s).

Environmental Implications

Cr(VI) is a carcinogen, and its presence in drinking water even at a trace level poses a great public health risk. It is possible that more rigorous regulations for Cr(VI) will be enforced in the future. In addition, to solve the issue of water scarcity, more desalinated water would be used as drinking water sources, and the desalinated water with higher bromide as a catalyst would enhance Cr(VI) formation by several times.15,36 Thus, Cr(VI) control is important for the drinking water distribution systems.

On the basis of realistic chemical conditions in the drinking water distribution system and using kinetics data obtained from this study, a simple adjustment of pH from 6 to 8 is predicted to decrease the formation of Cr(VI) by 31% (Figure 7). Combined with increasing the pH from 6 to 7, dosing 1 mg/L phosphate or 10 mg/L silicate can reduce Cr(VI) formation by 41% or 28%, respectively (Figure 7). Phosphate is reportedly dosed at 0.2–6 mg/L as PO4 in drinking water to control corrosion and lead release in England and the United States.54,55 In contrast, 5–25 mg/L as SiO2 of silicate is used primarily for iron/manganese sequestration in drinking water.33 Therefore, phosphate or silicate addition can be applied to inhibit Cr(VI) formation. Furthermore, since increasing the water alkalinity from 10 to 100 mg/L in the presence of 200 mg/L of background calcium reduced Cr(VI) formation by 26% (Figure 7), Cr(VI) formation could potentially be lowered further if a higher pH combined with phosphate dosing in combination with calcium and alkalinity were implemented.

Figure 7.

Figure 7

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Relative Cr(VI) concentration under different control strategies to uncontrolled conditions in drinking water distribution systems with 0.68 mg/L bromide in U.S. source waters. Control was at pH = 6, and other conditions were at pH = 7 unless specified.

Increasing pH/alkalinity in drinking water treatment plants can be a viable control strategy to reduce Cr(VI) formation. Unlike silicate, phosphate has both an enhanced and an inhibitive effect on Cr(VI) formation; thus, more attention should be paid to the use of phosphate in real DWDSs. In the future, it may be necessary to reduce Cr(VI) in drinking water further to protect public health. This study provides knowledge and understanding of Cr(VI) formation kinetics and related mechanisms in drinking waters with different chemistry.

In addition, Fe(II) and organic matter may exist in the iron corrosion scales to compete with Cr solids for the consumption of free chlorine; thus, they could reduce Cr(VI) formation.15 Fe(III) in the iron corrosion scales could precipitate out Cr(III) to form Cr(III)–Fe(III) hydroxides. MnO2 as a catalyst in the corrosion scales could accelerate the oxidation of Cr solids to form Cr(VI).56 Thus, when applying these control strategies in a real drinking water distribution system, more research effort should be taken to understand the complexity of the redox chemistry in the drinking water distribution systems in the future.

Acknowledgments

This research was supported by the U.S. National Science Foundation CAREER Program (CBET-1653931). We thank Soham Shah for BET analysis, group member Liang Wu for assistance in this project at UC Riverside, and Ken Ishida at Orange County Water District for his assistance in the preparation of the manuscript.

Supporting Information Available

The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/acs.est.2c05324.

  • Synthesis of different Cr(III) solids; Cr(III) solids concentration on the Cr(0) after 6 h oxidation reaction at different pH, phosphate, calcium, and alkalinity; calculation of maximum dissolved Cr(VI) factor and Ksp of different Cr(III) solids; kinetic model for the prediction of Cr(VI) concentrations in drinking water distribution systems (PDF)

The authors declare no competing financial interest.

Supplementary Material

es2c05324_si_001.pdf (921.7KB, pdf)

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Supplementary Materials

es2c05324_si_001.pdf (921.7KB, pdf)

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