Grantee Research Project Results
2004 Progress Report: Toxic Metal Ion-Synthetic Chelating Agent Interactions in Aqueous Media
EPA Grant Number: R829356Title: Toxic Metal Ion-Synthetic Chelating Agent Interactions in Aqueous Media
Investigators: Stone, Alan T. , Ball, William P.
Institution: The Johns Hopkins University
EPA Project Officer: Hahn, Intaek
Project Period: October 1, 2001 through September 30, 2005 (Extended to September 30, 2006)
Project Period Covered by this Report: October 1, 2003 through September 30, 2004
Project Amount: $333,057
RFA: Complex Chemical Mixtures (2000) RFA Text | Recipients Lists
Research Category: Hazardous Waste/Remediation , Land and Waste Management , Safer Chemicals
Objective:
This work focuses on metal ion-chelating agent mixtures that yield adverse environmental impacts that cannot be anticipated by studying either the metal ion or the chelating agent alone. The objectives of this research project are to:
- identify ways in which contaminant metal ions are rendered more toxic or otherwise problematic through reaction with naturally occurring or contaminant-derived chelating agents;
- and identify synthetic chelating agents, which are innocuous as “free” species but yield toxic or unusually reactive species when they react with naturally occurring metal ions.
Progress Summary:
Dissolution of Chromium(III) (Hydr)Oxides
Chromium-contaminated sediments (e.g., Baltimore Harbor) and soils (e.g., Hudson County, New Jersey) occur at several U.S. locations. Chelating agents are found in detergents, household and industrial cleansers, shampoos, and a wide range of other products. Incomplete removal during waste treatment may yield appreciable concentrations in wastewater effluent. Glyphosate and a handful of other herbicides possess chelating agent properties. Residential and agricultural uses of such herbicides represent potential nonpoint sources for entry into surface waters. Herbicides used for weed control at brownfield sites may also come into direct contact with chromium-contaminated soils.
Cr(OH)3(s, amorphous) synthesized in our laboratory serves as a surrogate for adsorbed and precipitated chromium(III) in soils. The naturally occurring chelating agent citric acid and the synthetic chelating agents iminodiacetic acid (IDA), nitrilotriacetic acid (NTA), ethylenediaminetetraacetic acid (EDTA), trans-1,2-cyclohexaneethylenediaminetetraacetic acid (CDTA), and trimethylenediaminetetraacetic acid (TMDTA) all bring about Cr(OH)3(s) dissolution within some portion of the pH domain of natural waters and soils. Chelating agent adsorption is a prerequisite for dissolution. To learn about mechanisms of dissolution, it is useful to explore dissolution rates normalized per mole of adsorbed chelating agent, RN. Under neutral and acidic pH conditions, RN is largely independent of pH for EDTA and TMDTA. With IDA, NTA, citric acid, and CDTA, however, RN increases substantially with decreasing pH, indicating either that:
- protons and chelating agents act in a synergistic manner to bring about dissolution;
- or adsorbed chelating agents shifts from a less reactive deprotonated form to a more reactive protonated form as the pH is decreased.
Interpreting changes in RN within the alkaline pH range are more complex. A dramatic spike in values of RN is observed in the pH range of 8.0 - 9.0 with the chelating agents EDTA and TMDTA and to a lesser extent with CDTA. The most plausible explanation is that hydroxide ions and chelating agents act in a synergistic manner to bring about dissolution. Citric acid, IDA, and NTA are considerably weaker chelating agents; relatively little dissolution is observed under alkaline pH conditions.
Chromium(III) Oxidation by MnO2(s,birnessite)
Manganese(III,IV) (hydr)oxides accumulate at oxic-anoxic interfaces within soils and sediments. It is important to establish how rapidly manganese(III,IV) (hydr)oxides oxidize chromium(III) to chromium(VI), because the two oxidation states have quite distinct toxicological properties. (Oxidation by O2 is known to be too slow to be significant in most environmental settings.)
Using capillary electrophoresis, we were able to discern specific chromium(III)-chelating agent complexes generated in the dissolution experiments described in the preceding section. As shown below, different complexes brought into contact with MnO2(s,birnessite) at pH 7.0 yield more than five orders of magnitude differences in rates of chromium(III) oxidation:
CrIII(EDTA)- greater than 2 years CrIII(NTA)° 34 minutes CrIII(HEDTA)- 1.6 years CrIII(IDA)+ 8.2 minutes CrIII(IDA)2- 12 days Note: Hydroxyethylenediaminetriaacetic acid (HEDTA) is structurally similar to EDTA. One of the carboxylate Lewis Base groups has been replaced by an ethanol group.
As noted in previous progress reports, chromium(III) oxidation to chromium(VI) liberates the chelating agent. Free chelating agent molecules are relatively resistant to oxidative degradation under strongly alkaline conditions, but degrade more rapidly as the pH is decreased. Oxidative degradation of free chelating agent molecules is important because
- degradation products are much less effective than parent chelating agents in solubilizing chromium(III);
- and the oxidative degradation reaction consumes MnO2(s,birnessite), and thereby lessens rates at which chromium(III)-chelating agent complexes are oxidized.
Reductive Dissolution of MnO2 by Organic Ligands With Keto-Enol Structures
As noted in previous progress reports, naturally occurring and synthetic chelating agents react with manganese(III,IV) (hydr)oxides solely via ligand-assisted dissolution (yielding soluble manganese(III)-chelating agent complexes), solely via reductive dissolution (yielding oxidized organic products and manganese(II)), or via both processes simultaneously.
Previous progress reports noted that malonic acid reacts with MnO2(s,birnessite) solely via reductive dissolution. We hypothesized that malonic acid is activated toward oxidation through the generation of the enol tautomer. Once the tautomer is formed, it is consumed by two parallel pathways:
- oxidation to tartronic acid, glyoxylic acid, and finally formic acid;
- and oxidation to dihydroxymalonic acid followed by hydrolysis to oxalate, and then oxidation to carbon dioxide.
If our hypothesis that reaction takes place via the enol tautomer is correct, then other organic substrates undergoing keto-enol tautomerization should react via a similar mechanism. Malonic acid, acetylacetic acid, and acetylacetone provide for a particularly informative comparison. As noted in Figure 1, keto-enol equilibria, complete with protonation level determinations, have been worked out for acetylacetic acid and for acetylacetone. With acetylacetic acid, the enolate species “D” accounts for approximately 0.3 percent of the total dissolved concentration. With acetylacetone, the deprotonated enolate species E- dominates above pH 9, and the protonated species HE- accounts for approximately 43 percent of the total dissolved concentration within the neutral and acidic range.
It would be valuable to know the extent of adsorption onto MnO2, but the subsequent redox reaction is too fast for reliable measurements to be made. For comparison purposes, we have measured the extent of adsorption onto a non-redox active surface, TiO2. The greatest extent of adsorption for malonic acid is approximately three-times higher than that for acetylacetic acid and for acetylacetone. The maximum extent of adsorption takes place at pH 3.0 for acetylacetic acid, at pH 5.0 for malonic acid, and pH 8.2 for acetylacetone (Figure 2).
Figure 3 shows rates of MnO2 reductive dissolution as a function of pH. Acetylacetic acid shows the steepest decline in rate as the pH is increased, the expected finding based upon the low pH of maximum adsorption for this compound. The higher pH of maximum adsorption for malonic acid makes the decline in rate with increasing pH more gradual. Acetylacetone, with a maximum in absorption at slightly alkaline pH, exhibits the most gradual decline in rate with increasing pH. Near pH 7, rates of reductive dissolution by acetylacetone are nearly independent of pH. Efforts are currently underway to change the substituents within the acetylacetone structure to learn more about how keto-enol and protonation equilibria affect rates of MnO2 reductive dissolution (see Figure 4).
Reaction of Citric Acid with MnO2
Previous progress reports have presented our findings regarding the reaction of citric acid and structurally related compounds with MnO2 and MnOOH. Our objective in this work is to provide the basis for predicting rates of dissolved MnII and MnIII appearance in solution once contact with chelating agents has taken place. Dissolved MnIII species are especially relevant because of high inherent toxicities.
During the past year, a major portion of our effort has gone into the quantitative modeling of the citric acid plus MnO2 reaction. The reaction is strongly autocatalytic in nature, as described by the following two elementary reactions:
A + S C + P k1* (1) A + S + C 2C + P k2* (2)
In this scheme, S represents MnIII,IV surface sites, A represents unreacted citric acid, C represents Mn II(aq), and P represents the citric acid oxidation products 2-ketoglutaric acid and acetylacetic acid. The differential equation for appearance of P can be solved analytically. Values for k1 and k2 were obtained by applying this model to an experiment employing 50 μM citric acid and 200 μM MnO2 at pH 7.1. The resulting numerical simulation is presented as Case B in Figure 5.
Reactions 1 and 2 represent two reactions in parallel. Reaction 1 is a simple reductive dissolution reaction, initiated by the adsorption of citric acid onto the MnO2 surface. Reaction 2 yields the same products but is autocatalytic in nature. MnII product forms a complex with citric acid, such as Mn II(citrate)-. MnIII,IV surface sites oxidize Mn II(citrate)- to MnIII(citrate)°(aq). Finally, intramolecular electron transfer converts MnIII(citrate)°(aq) into Mn II, 3-ketoglutaric acid, and acetylacetic acid.
As indicated by Cases A and C in Figure 5, raising k1* at constant k2* causes the induction period characteristic of autocatalysis to shorten and eventually disappear. Similarly, raising k2* at constant k1* (Cases D and E) shortens the induction period. This quantitative model has been crucial to exploring how pH, citrate concentration, and the presence of additional chemicals (e.g., ZnII, pyrophosphate, orthophosphate) affect time course plots.
Figure 1. Equilibrium Speciation of 5 mM Acetylacetone and 5 mM Acetylacetic Acid as a Function of pH. Different protonation levels and keto-enol tautomers are noted.
Figure 2. Adsorption of Malonic Acid, Acetylacetic Acid, and Acetylacetone Onto TiO2 as a Function of pH. Reaction conditions: 100 μM organic substrate, 5.0 g/L TiO2, 1.0 mM NaCl (pH is fixed using additions of HCl and NaOH).
Figure 3. Rates of MnO2(s,birnessite) Reduction by Malonic Acid, Acetylacetic Acid, and Acetylacetone as a Function of pH. Reaction conditions: 200 μM MnO2(s), 5.0 mM organic reactant, and 10 mM buffer (4-sulfobenzoic acid buffer for pH 3.0 and 3.5, butyric acid for pH 4.0, 4.5, and 5.0, MES for pH 6.6, and MOPS for pH 7.0).
Figure 4. Changing the Substituents Within the Acetylacetone Structure Affect Keto-Enol and Protonation Equilibria. These changes, in turn, affect rates of MnO2 reductive dissolution (ro, in micromoles .liter-1hour-1).
Figure 5. The Reaction Between Citric Acid and MnO2 Exhibits Autocatalysis. Simulated time course plots (upper) and rates as a function of time (lower, corresponding to the first derivative) are based upon Reactions 1 and 2, which are presented in the text. Case B represents the best fit to an experiment employing 50 μM citric acid and 200 μM MnO2 at a pH of 7.1.
Future Activities:
In addition to completing manuscripts in the areas mentioned previously, we also would like to pursue the contaminant mixture issues described below.
Neutral Metal Ion-Chelating Agent Complexes
Transport of intact cationic or anionic metal ion complexes into cells requires receptor sites on the cell membrane surface. Neutral complexes, on the other hand, can pass directly through the cell membrane if they are hydrophobic enough. Octanol-water partition coefficients (Kow values) are a good measure of hydrophobicity, but are not known for most neutral complexes.
Capillary electrophoresis allows us to quantify concentrations of cationic and anionic complexes. From total metal ion concentrations and mass balance, we can calculate concentrations of neutral complexes. (In some instances, neutral complexes can be quantitatively determined using capillary electrophoresis.) The protonation level and the ratio of metal ion to the chelating agent also need to be determined.
MnIII(citrate)°(aq), generated by reaction of citric acid with manganese(IIII,IV) (hydr)oxides, serves as an impetus for new research in this area. The neutral form should be predominant under slightly acidic conditions (e.g., pH 4.0). As the pH is raised, however, a point should be reached where the alcoholate group becomes deprotonated. The stoichiometry for this deprotonation reaction is typically written as:
MnIII(citrate) o(aq) + H2O = MnIII(OH)(citrate)-(aq) + H + pKa (3)
We would like to determine the pKa for Reaction 3. We would also like to explore the kinetics of MnIII(citrate)°(aq) production arising from the oxidation of Mn II(citrate)-(aq) by molecular oxygen (O2). Reaction kinetics are likely to be acutely sensitive to the citric acid concentration and to the ratio Mn T/CITT (total dissolved manganese divided by total dissolved citrate).
The work we envision would help in evaluating possible impacts of a variety of synthetic chelating agents in soils. Our findings also should be relevant to investigations into the human toxicity of inhaled or ingested manganese.
Cr VI Adducts, Inert CrIII Species, and Other Products of Chromate Ion Reaction With Organic Reductants
Previous sections have described our extensive research into dissolved chromium-containing species derived from particulate CrIII. We would also like to investigate dissolved chromium-containing species derived from chromate ion (HCrO4-/CrO42-).
Chromate is used synthetically as an oxidant of organic compounds under strongly acidic conditions. In many instances, reaction is believed to begin via formation of chromate-organic adducts, which subsequently undergo intramolecular electron transfer. Subtle changes in UV/visible spectra have been the primary evidence for adduct formation. Capillary electrophoresis would enable us to physically separate adducts from parent species. Because we employ a photodiode array detector, it should be possible to obtain spectra for specific adducts.
Over time, intermediate oxidation states of chromium (i.e., CrV and CrIV) should be superceded by CrIII. As our prior work has demonstrated, CrIII complexes are substitution-inert and, therefore, readily amenable to analysis by capillary electrophoresis. Determining the stoichiometry, protonation level, and bound ligand identity of these product complexes would tremendously improve our ability to evaluate the adverse environmental effects of chromate ion-chelating agent mixtures.
Journal Articles:
No journal articles submitted with this report: View all 15 publications for this projectSupplemental Keywords:
chelating agents, toxic metal ions, complexation, oxidation-reduction reactions, structure-reactivity relationships, metal ion-chelating agent mixtures, contaminant-derived chelating agents, synthetic chelating agents, wastewater effluent,, RFA, Scientific Discipline, Waste, Ecosystem Protection/Environmental Exposure & Risk, Environmental Chemistry, chemical mixtures, Fate & Transport, Hazardous Waste, Ecology and Ecosystems, Hazardous, hazardous waste treatment, complex mixtures, contaminated sediments, fate and transport, fate and transport , biodegradation, contaminant biodegradation rates, hazardous organic substances, environmental transport and fate, chemical kinetics, hazardous chemicals, capillary electrophoresis, contaminated soils, analytical modelsProgress and Final Reports:
Original AbstractThe perspectives, information and conclusions conveyed in research project abstracts, progress reports, final reports, journal abstracts and journal publications convey the viewpoints of the principal investigator and may not represent the views and policies of ORD and EPA. Conclusions drawn by the principal investigators have not been reviewed by the Agency.