Chemicals in the Environment - ACS Publications - American

Chapter 4

Speciation and Chemical Reactions of Phosphonate Chelating Agents in Aqueous Media 1

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Downloaded by MIT on June 3, 2013 | http://pubs.acs.org Publication Date: January 24, 2002 | doi: 10.1021/bk-2002-0806.ch004

A. T. Stone , M. A. Knight , and B. Nowack 1

Department of Geography and Environmental Engineering, 313 Ames Hall, The Johns Hopkins University, Baltimore, MD 21218 Environmental Engineering and Science Program, 138-78 Keck Laboratories, The California Institute of Technology, Pasadena, C A 91125 Institute of Terrestrial Ecology, ΕΤΗ Zürich, Grabenstrasse 11a, CH-8952 Schlieren, Switzerland 2

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ABSTRACT 2-

Organic chemicals possessing phosphonate functional groups (RP0 ) are being used in a growing number of applications. Through a series of examples, the speciation and chemical reactions of phosphonates in aqueous solutions and (hydr)oxide mineral suspensions are systematically explored. Concepts are introduced that are useful for assessing the consequences of intentional or inadvertent release into environmental media. Such concepts should aid the development of more environmentally benign synthetic organic chemicals. 3

INTRODUCTION Chemists mix and match various functional groups and structural moieties in an attempt to develop manufactured chemicals with desirable properties. In recent decades, phosphonate functional groups (RP0 ) have appeared in a 2

3

© 2002 American Chemical Society In Chemicals in the Environment; Lipnick, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.

59


60 rapidly growing number of manufactured chemicals. Phosphonate-based chelating agents are used to prevent the formation of undesirable precipitates, and to protect others from dissolution. Such chelating agents are used to depress free metal ion activity and to increase total dissolved metal ion concentrations. Others can convert unreactive metal ions into reactive ones (and vice versa). In many of these applications, molecular charge, protonation level, and ability to bind metal ions are critical; phosphonate-based chelating agents yield properties and reactivities that are more desirable than those obtained with alternative Lewis Base functional groups. Other classes of phosphonate-containing synthetic organic compounds (or their metal ion complexes) possess pronounced biological properties. Phosphonates are used as pesticides, as growth regulators, as pharmaceuticals, and as sterilants. Biologically-active phosphonates possess some chemical traits that are similar to a naturally-occurring biochemical, and some traits that are different. The proper balance between similar and dissimilar traits is important for achieving efficacy. Desirable (or undesirable) attributes of manufactured chemicals must always be viewed within the context of a particular system and it contents. Phases (and interfacial regions) set the stage for partitioning. Other chemicals present in the system affect speciation, and serve as reactants, catalysts, or inhibitors. Organisms exhibit enormous variations i n their susceptibilities towards different chemical species. This chapter begins with a comparison of the Lewis Base properties of carboxylate, phosphonate, and amine functional groups. The formation of metal ion-chelating agent complexes in solution and the adsorption of phosphonates onto (hydr)oxide mineral surfaces are then discussed. In the remaining portions of this chapter, interconnections between the coordination chemistry and chemical reactivity of phosphonates are explored. The underlying message is that the environmental chemistry of metal ions (whether dissolved or particulate-bound) and phosphonates are closely linked. The more we know about coordination chemistry, the better we can predict the effects of phosphonate-containing synthetic chemicals i n the environment.

Molecular Charge, Basicity, and Equilibrium Speciation in Solution E D T A (ethylenediaminetetraacetic acid), first patented i n Germany i n 1935 and in the U.S. in 1946, has long dominated the market for synthetic chelating agents (1). Four carboxylate ( R C O O ) and two amine (R N) Lewis Base groups are favorably placed within the E D T A structure for occupying up 2

In Chemicals in the Environment; Lipnick, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.


61 to six coordinative positions of a central metal ion. Suppose that a carboxylate group is replaced by a phosphonate group. How does the nature of metal ionchelating agent binding change, and how is the equilibrium speciation of the metal ion of interest altered? These and other questions pertaining to chelating agent design will be addressed i n this section. Acetic acid, methylenephosphonic acid, and methylamine are useful archetypes for more complex molecules (Figure 1). Oxygen atoms serve as sites of proton and metal ion binding to carboxylate and phosphonate groups. Bonds are both ionic and covalent in nature. As far as ionic bonding is concerned, bond strength increases as the charge-to-radius ratio of the Lewis Base group increases. The dianionic nature of the phosphonate group boosts coordination relative to the carboxylate group, but this effect is partially offset by its' larger size (2). It is also worth noting that the carboxylate group is planar, while the phosphonate group is tetrahedral (2). Although the amine group bears no net charge, there is significant electron density directed towards the free p-orbital (3). The amine group is classified as a "hard" Lewis Base, although the mix of covalent versus ionic contribution to bonding is somewhat higher than observed with carboxylate and phosphonate groups (4). When comparing one ligand to another, basicity is of paramount importance. A s the ability of a Lewis Base to coordinate protons is increased, its' ability to coordinate metal ions also increases. Hence, with increasing p K , there is a corresponding increase i n logK values for metal ion coordination. O f the three ligands shown in Figure 1, methylamine is by far the most basic, and yields logK values (e.g. for complexes with +11 metal ions, not shown) that are higher than for the other two ligands. It is important to keep i n mind, however, that metal ions must compete with protons for available ligands. A s solution pH is decreased, a point is reached where protons out-compete metal ions for coordinating available ligands. Similarly, ligands must compete with hydroxide ions (OH") at high p H for available metal ions. A s solution p H is increased, a point is reached where either hydroxo complexes i n solution (e.g. Fe OH*) or the precipitation of (hydr)oxide solids (e.g. Fe (OH) (s)) dominate metal ion speciation. Replacing a carboxylate group with a phosphonate group has a number of important consequences. Near neutral pH, the phosphonate dianion (RP0 ~) is more basic than the carboxylate group, and, all other factors being equal, better able to coordinate metal ions. The methylenephosphonate dianion, for example, exhibits a p K of 7.82, while the corresponding p K for acetate monoanion is 4.67 (Figure 1). Even under acidic conditions, phosphonates are considerably more effective chelating agents than the corresponding carboxylates (5). For example, while methylenephosphonic acid is in a monoanionic form above p H 2.3, acetic acid is i n a protonated, unavailable form until a p H of 4.6 is reached a

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63 (see Figure 1). Greater electron density resides on the oxygen atom i n P=0 than i n C = 0 (6). Thus, phosphonate diesters (RP(0)(OR) ) coordinate metal ions more strongly than carboxylate esters (RC(O)OR). Coordination to the oxonate oxygen may also be important in species such as RP(0)(OH)0" and RP(0)(OH) . Two (or more) Lewis Base groups can simultaneously coordinate a central Lewis Acid, provided that the resulting ringed structure is not unduly strained. Protons assume this arrangement via relatively weak hydrogen bonds. Metal ions much more readily exhibit coordination numbers of two or more. The corresponding structure, termed a "chelate", is often an effective means of ensuring that metal ions successftdly out-compete protons for available Lewis Base Groups. In Figure 2, four new ligands are assembled using amine, phosphonate, and carboxylate Lewis Bases and methylene (-CH -) linkages. For the most part, the methylene linkages electronically isolate each Lewis Base, minimizing inductive and resonance effects on basicity. (A modest electronic effect causes the p K for the amine group in glycine and iminodiacetic acid (Figure 2) to be one p H unit lower than the p K for methylamine.) These structures enable formation of five-membered hydrogen-bonded and chelate rings, as illustrated in Scheme 1. The dramatic lowering of p K values for phosphonate dianions and carboxylate monoanions in all four compounds can be attributed to hydrogen-bonding. I D A (iminodiacetic acid) and glyphosate (N-(phosphonomethyl)glycine), used for the calculations shown in Figures 3-6 are high-volume synthetic chelating agents. (Equilibrium constants were derived from (7)). Despite having three Lewis Bases suitably placed for chelate ring formation, l.OxlO" M I D A fails to capture Fe (l.OxlO" M ) below p H 6.0, as shown i n Figure 3. Metal coordination is driven by bond formation with the two free carboxylate groups. The amine group is protonated below p H 9.6, and requires metal ioninduced deprotonation i n order for Fe L° to form. Glyphosate captures Fe at slightly lower pHs than observed with I D A , a result attributed to the monoprotonated species Fe HL°. The placement of protons within free (8) and metal ion-complexed phosphonate species is currently being investigated i n a number of laboratories, frequently with N M R (9-14). The stoichiometry Fe HL° can either be explained by (i) Fe coordination via carboxylate and phosphonate groups, without participation by R N H / or (ii) a proton shift to one of two anionic oxygen atoms of the phosphonate group, allowing both the amine and the phosphonate to participate in Fe coordination (as evoked by Sawada et al. (10) for the coordination of Fe by N T M P ) . Figure 3 illustrates one additional distinction between the dicarboxylate and the carboxylate/phosphonate ligands. With IDA, an increase i n p H causes 2


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In Chemicals in the Environment; Lipnick, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002. 4

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Figure 3. Speciation of 1. 0x10' MFe in the presence of L 0x10" M iminodiacetic acid and l.OxlO" Uglyphosate (10 mMNaN0 constant ionic strength medium).

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a shift away from the 1:1 complex (Fe L°) towards the 1:2 complex ( F e L ) . With glyphosate, the 1:2 complex is a much less important species. This is a common observation with phosphonate chelating agents (2), and it arises from unfavorable steric and coulombic interactions between two large, negativelycharged functional groups. This phenomenon can also be seen i n speciation calculations where the p H is fixed and the total ligand concentration increased (Figure 4). The 1:1 complex with glyphosate ( F e L ) forms at lower total ligand concentrations than the corresponding complex with IDA, owing to the higher basicity of the phosphonate Lewis Base group, and persists to much higher total ligand concentrations, owing to the relative instability of the 1:2 complex, Fe L ". 2

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Calculations performed with Fe are illustrated i n Figure 5. The +ÏÏI charge and smaller radius of this metal ion yield substantially greater logK values for the formation of hydroxo species and (hydr)oxide solids. Fe(OH) (amorphous), used in the calculations as the solubility-limiting phase, forms at p H 3.9 in the presence of l.OxlO" M IDA, and at p H 5.1 i n the presence of l.OxlO" M glyphosate. The higher basicity and negative charge of the phosphonate-containing ligand clearly improves its performance as a chelating agent. It is interesting to note that the ternary complex with hydroxide ion, Fe (OH)L , is more important i n the presence of I D A than i n the presence of glyphosate. We can speculate that more hydroxide-ligand coulombic repulsion occurs in the glyphosate complex than i n the I D A complex. 3

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With I D A and glyphosate, logK values are available for the complexation of both Fe and F e . Based upon this information, we can explore the effects of chelating agent concentration on the reduction potential for the Fe *(aq) + é = Fe (aq) half-reaction. Suppose a system is comprised of a fixed total amount of reduced and oxidized forms of a metallic element. In our case, l.OxlO' M Fe and l.OxlO" M Fe have been added to solution. Each oxidation state will be distributed over a variety of species with different numbers of coordinated ligands and different protonation levels. With Fe and glyphosate, for example, F e (total dissolved Fe ) will be comprised of Fe (aq), Fe HL°, Fe L", and Fe L ", plus other, less abundant species. F e (total dissolved F e ) will be comprised of F e L ° and F e ( O H ) L , plus other less abundant species; precipitated solids (e.g. Fe(OH) (am.)) must, of course, be appropriately substracted from the mass balance equation. Once the free metal ion concentration is calculated for each oxidation state, the reduction potential is found using the Nernst Equation: 11

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70 (In this case, E° = +0.770 volts.) As shown i n Figure 6, increasing amounts of either I D A or glyphosate cause a decrease i n reduction potential under acidic p H conditions, and an increase i n reduction potential under alkaline conditions. Figures 3 and 5, along with a consideration of Fe(II) and Fe(III) solubility in the presence and absence of chelating agent, readily explain this surprising finding. Under strongly acidic conditions, Fe (OH) (s) does not form, and addition of sufficient amounts of chelating agent brings about the complete dissolution of Fe (OH) (am). F e and F e (total dissolved concentrations of Fe and F e ) are therefore constant. Complex formation constants (logKs) for oxygen-donor and most nitrogen-donor ligands are higher for Fe than for Fe . Hence, as more chelating agent is added, [Fe (aq)] decreases more than [Fe (aq)], causing the reduction potential to decrease. Under alkaline conditions, Fe (OH) (s) requires pHs greater than 9.6 to form, while Fe (OH) (am) forms over a much wider p H range. As long as any Fe (OH) (am) is present in the system, [Fe (aq)] at constant p H is fixed by the dissolution reaction Fe (OH) (am) + = Fe (aq) + 3 H 0 . Under these conditions, increasing total ligand concentrations can still depress [Fe (aq)], but leave [Fe (aq)] unchanged. The net result is an increase i n the reduction potential. N T A (nitrilotriacetic acid), which contains one more carboxylatecontaining arm than IDA, is a much more effective chelating agent (Figure 7). Under the conditions employed (10 μΜ F e and 100 μΜ chelating agent), I D A prevents Fe(OH) (am) formation up to a p H of 3.7, while N T A prevents formation up to p H 7.0. With both I D A and N T A , ternary complexes with hydroxide ions ( O H ) are important. N T M P (nitrilotri(methylenephosphonic acid)) is the phosphonate analog to N T A . As shown in the right panel of Figure 7, N T M P performs substantially better than I D A and glyphosate i n solubilizing Fe , but is slightly less effective than N T A under neutral and alkaline conditions. The advantages of a phosphonate group ν/5· a vis a carboxylate group become less important as the denticity of the chelating agent increases and as the number of phosphonate groups within the molecule increases. There are several reasons for this (2). Unfavorable electrostatic interactions between arms of a multidentate ligand are more important for dianionic functional groups than for monoanionic functional groups. With multidentate ligands, the greater steric requirements of the tetrahedral phosphonate group relative to the planar carboxylate group becomes important. The driving force for protonation becomes quite large, and each monoprotonated phosphonate group is less basic (and hence capable of forming bonds to metal ions) than a deprotonated carboxylate group. n

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Figure 6. Effect of increasing iminodiacetic acid and glyphosate concentrations on the reduction potential for the Fe (aq) + e~ -> Fe (aq) half-reaction. Reaction conditions: l.Oxl a M Fe", l.OxlO' M Fe , 10 mM NaN0 constant ionic strength medium. Fe(OH) (amorphous) serves as the solubility-limiting phase for Fe . The dashed line corresponds to the reduction potential in the absence of added ligand.



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Figure 7. Speciation of l.OxlO' M Fe in the presence of l.OxlO" M NTA and L Oxl O" M NTMP (10 mMNaN0 constant ionic strength medium). Fe(OH) (amorphous) is used as the solubility-limiting phase. logK values for Fe -NTMP complexes were obtained from ref (72).

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73


Rates of Metal Ion and Ligand Exchange Surface waters, soils, and sediments are often divided into "parcels" that are modeled as open systems (16). Inputs and outputs of water, water-borne solutes, and water-borne particles change over time. A s a consequence, chemical conditions within the parcel change. If metal ion exchange and ligand exchange reactions take place quickly, equilibrium descriptions of chemical speciation are appropriate. If these reactions take place slowly, however, then an appraisal of speciation requires knowledge of reaction kinetics (17). In dissociative exchange, a coordinated ligand leaves the inner coordination shell before the replacement ligand enters. In associative exchange, the incoming ligand enters before the original ligand exists, yielding an intermediate with a higher coordination number. Interchange mechanisms (I and I ) fill out a continuum between the two mechanisms (18). The incorporation of isotopically-labelled water into the inner coordination sphere of aquated metal ions is one of the simplest exchange reactions observed. Rate constants for this reaction are used by chemists to make generalizations about exchange rates for different metal ions. A few illustrative rate constants for this reaction are presented below (from (19) and (20)): d

a

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2

1

n

Mn (H 0) Fe (H 0) Co (H 0) Ni (H 0) Cu (H 0)

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Differences in speciation have an enormous effect on rates of exchange. C r ^ O H J ^ C O n - i ^ exhibits water exchange rates that are 75-times higher than for C i ^ ( H 0 ) ( 1 8 ) . Under any realistic set of chemical conditions, a chelating agent of interest is likely to exist i n two or more protonation levels, and is likely to be coordinated to common metal cations such as C a and M g . A metal ion of interest is likely to be coordinated by major anions (e.g. CI", S0 ", and C 0 " ) or naturally-occurring organic ligands. As reflected in rate constants for water exchange (above), Cu undergoes ligand exchange far more rapidly than the other +11 metal ions listed above. Despite this fact, the achievement of equilibrium between Cu and strong ligands like E D T A under seawater 3+

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74 conditions can require timescales of minutes to hours (21-23). Similar results have been obtained with F e F e speciation during infiltration into groundwater appears to be under kinetic control (24). Twenty days are sometimes required for equilibrium between F e and strong ligands like E D T A to be achieved i n streamwaters (25). N i E D T A " , formed during wastewater treatment, has been shown to persist for significant periods of time following release into estuarine waters (26). Cr exhibits rates of water exchange that are orders-of-magnitude lower than for the other metal ions listed above. It is tempting to assume that all Cr complexes are exchange-inert, and hence immutable. Several points must, however, be kept i n mind. If rates of dissociative exchange are low, exchange via associative mechanisms might become important. (When associative mechanisms are predominant, exchange rates become a function of the identity of the incoming ligand.) With multidentate ligands, the linkage between two Lewis Base groups may introduce steric strain that causes exchange rates to increase (27). Other solute species can catalyze exchange reactions. Bicarbonate ion, for example, facilitates entry of E D T A into the inner coordination sphere of Cr" , and hence catalyzes C r ^ E D T A " formation (28). A preliminary experiment has been conducted which compares the effects of carboxylate and phosphonate functional groups on rates of C r ligand exchange. Laboratory practices outlined i n previous publications (29-30) were followed. A l l chelating agents were purchased from Aldrich Chemical Co. and Fluka Chemical Co. and used without additional purification. The potassium salt of u-fac-C^TDAh' was synthesized according to procedures outlined by Weyh and Hamm (31). This is one of three possible isomers of the 1:2 complex with iminodiacetic acid shown in Scheme 2. A s described i n a previous publication, the u-fac and s-fac isomers can be readily distinguished from one another using capillary electrophoresis (29). The mer- isomer is probably not important under the conditions employed i n our experiments. Track-etched polycarbonate filters and hydrophilic cellulose membrane filters (both 0.2 micron pore diameter, Whatman Filter Co.) were used to recover supernatant solutions. Total chelating agent and metal ion-chelating agent complex concentrations were determined using a Quanta 4000E capillary electrophoresis instrument (Waters Corp.) and bare fused-silica capillaries (75 microns wide, 60 cm long, Polymicro Technol.). The capillary electrolyte consisted of 25 m M phosphate buffer (pH 7.0) and 0.5 m M tetradecyl trimethylammonium bromide (TTAB) electroosmotic flow modifier (see (29)). In the absence of other chelating agents, there is a gradual decline i n the concentration of u-fac-QF(lDPi)i, attributed to interconversion to the s-fac form of the complex. Our exchange experiments employed 100 μΜ u-facC r ^ I D A V , 5.0 m M acetate buffer, and 1.0 m M concentrations of the m

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76 following chelating agents: N T A , E D T A , N T M P , B P M G (bi(phosphonomethyl)glycine), P M I D A (phosphonomethyliminodiacetic acid), and E D T M P (ethylenediaminetetra(methylenephosphonic acid)). N T A (with no phosphonate groups), P M I D A (with one phosphonate group), and B P M G (with two phosphonate groups) share the same three-arm structure. As shown i n Figure 8, the two chelating agents possessing phosphonate groups capture Cr via ligand exchange more rapidly than N T A . Similarly, E D T M P (with four phosphonate groups) captures C r ^ considerably more rapidly than E D T A (with no phosphonate groups). We can conclude that phosphonate groups facilitate the kinetics of Cr capture at this pH. A second set of experiments has been performed i n solutions buffered to p H 7.2 using 5.0 m M M O P S ^ ο φ ΐ ι ο ϋ η ε ρ π ^ η ε β ι ι ΐ ί ο η ΐ ϋ acid). During 21 days of reaction, no decline i n w/bc-Cr (EDA)2" concentration was observed, regardless of whether strong chelating agents were added. 111


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in

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These experiments demonstrate that rates of exchange reactions involving C r and phosphonate chelating agents are sensitive to p H and medium composition, and slow enough to control speciation i n a number of important aqueous systems. A tempting response is to develop computer-based models for predicting toxic metal ion speciation under kinetically-controlled conditions. It should be kept i n mind, however, that the data requirements of kinetic models (e.g. forward and reverse rate constants, rate constants for catalytic processes, pH and major ion concentrations as a function of time) are much greater than those of equilibrium models. Quantitative models serve a didactic purpose, and encourage us to think realistically about factors affecting speciation. O n a practical level, however, it is time for attention to be focused on analytical methods for directly determining speciation in environmental samples. Direct speciation measurements provide a strong basis for making environmental management decisions, and provide a way of testing computer-based speciation models. A n integration of speciation measurements with modeling efforts provides the best prospect for predicting speciation. 01

Adsorption and Precipitation (Hydr)oxide and aluminosilicate products of rock weathering comprise much of the available surface area in soils and sediments. Such inorganic solids can be either amorphous or (micro)crystalline. In the presence of vapor or liquid water, surfaces are hydrated. In adsorption experiments, it is common to compare adsorption onto two inorganic solids that are chemically distinct yet possesses the same physical characteristics (e.g., surface area and surface charge). Similarly, the adsorption behavior of two chemically distinct solutes


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φ

SO

80 "Ν" represents a nitrogen-donor atom (e.g., from an amine group) while each " 0 " represents an oxygen-donor atom (e.g., from either a carboxylate or phosphonate group). Depending upon the accessibility of one, two, or three coordinative positions around a central F e ("mononuclear") atom, monodentate, bidentate, and tridentate surface complexes are possible. If the "arms" of the chelating agent connecting the Lewis Base groups are long enough, bridging between two surface-bound F e atoms is possible (e.g. the binuclear bidentate complex). As mentioned above, Lewis Base groups not engaged i n bond formation "dangle" at some distance away from the surface. Details regarding the adsorption of monophosphonates (34-37) and polyphosphonates (41-42) are provided in the references just cited. Here, we focus upon the adsorption properties of polyphosphonates, which are particularly intriguing. A comprehensive study of the adsorption of eight phosphonate-containing compounds onto FeOOH(goethite) has recently been published (41). A t p H 10, the extent of adsorption of 40 μΜ N T M P onto 0.42 g/L FeOOH(goethite) is approximately 30 percent. A thousand-fold increase i n the background electrolyte concentration (from 1.0 millimolar to 1.0 molar N a N 0 ) has no effect on the extent of adsorption. This lack of an ionic strength effect is surprising, given that FeOOH(goethite) bears a net negative charge at this pH, and that the predominant solution phase species for N T M P is HL ". Long-range electrostatic interactions which repulse like-charged ions away from the surface are much stronger at low ionic strength than at high ionic strength (33). Three different phenomena may be responsible for the lack of an ionic strength effect, (i) Each Fe -phosphonate bond lowers the charge of the adsorbed complex by one unit. If all three phosphonate groups of N T M P are bonded in this way, the surface complex would have a stoichiometry and charge of (>Fe ) HL ". (ii) Phosphonate-containing "arms" that are not bonded to the surface may position themselves away from the surface, where electrostatic interactions with the charged surface are diminished, (iii) Phosphonatecontaining "arms" that are not bonded to the surface may protonate or form ion pairs with electrolyte cations (e.g. Na ) to a greater extent than phosphonate groups i n bulk solution. When added phosphonate concentrations are increased at fixed surface loading and pH, the extent of adsorption eventually levels out. Maximum extents of adsorption provide an estimate of the "footprint" of each molecule on the surface. Footprint size may reflect the numbers of surface-bound ¥q atoms engaged i n bond formation with the adsorbed molecule, or may indicate crowding by neighboring molecules. Electrostatic repulsion between likecharged adsorbate molecules on the surface may also be important. m


m

3

5

ffl

m

2

3

4

m



81 Maximum extents of adsorption at p H 7.2 for eight phosphonatecontaining compounds are shown i n Figure 10. Going from left to right, the first three compounds, H M P (hydroxymethylphosphonic acid), M P (methylphosphonic acid), and H E D P ( 1 -hydroxyethane-1,1 -diphosphonic acid) do not contain amine groups. Maximum surface coverage by the diphosphonate H E D P is half that of the monophosphonate H M P and 40 % less than the monophosphonate M P . The next three compounds possess a primary, a secondary, and a tertiary amine group, and one, two, and three phosphonate groups, respectively. The effect of amine groups on footprint size does not result i n clear trends; maximum surface coverage by A M P is 18 % less than by M P , but maxium surface coverage by IDMP is 24 % greater than by HEDP. (Hydrogen-bonding, chelating ring formation, and differences in adsorbed species charge may all be contributing factors.) Maximum surface coverage by the triphosphonate N T M P is 35 % lower than the coverage than by the diphosphonate IDMP. E D T M P , with four phosphonate groups, and DTPMP (diethylenetrinitrilopentakis(methylenephosphoriic acid), with five phosphonate groups, adsorb to approximately the same extent as N T M P . The fact that the fourth and fifth phosphonate groups do not increase the molecular footprint indicates that the additional phosphonate groups are not associated with the surface. Steric considerations may prevent them from coordinating surface-bound F e atoms, and electrostatic repulsion may force them away from the anchoring phosphonate groups (41). Engineered systems and environmental media contain other solutes that can influence the extent of phosphonate adsorption. Other ligands (e.g. carbonate, phosphate, natural organic matter) can compete with phosphonate molecules for available surface sites, thereby lowering the extent of phosphonate adsorption. Metal ions can either raise or lower the extent of phosphonate adsorption, depending upon the balance of a number of conflicting phenomena. A s discussed in previous sections, the formation of dissolved metal ion-phosphonate complexes occurs to the greatest extent i n the mid-pH range. By providing another "compartment" for phosphonates in solution, complex formation in solution can lower adsorption. Metal ions that adsorb as separate entities cause the surface charge to shift in the positive direction. A t high pH, this shift i n surface charge lessens the extent of long-range electrostatic repulsion between negatively-charged surfaces and phosphonate polyanions, thereby raising the extent of adsorption. Multidentate ligands can bridge between surface-bound metal atoms and dissolved metal ions, forming "ligand-like" ternary complexes. Alternatively, metal ions can bridge between surface-bound oxo or hydroxo groups and dissolved ligands, forming "metal ion-like" ternary complexes. Cooperative adsorption of this kind, illustrated in Figure 11, is the subject of a comprehensive review (43). m



82



Figure 10. Maximum extents of adsorption for eight phosphonate-containing compounds onto 0.42 g/L FeOOH(goethite). Reaction conditions: 1 mM MOPS buffer (pH 7.2), 10.0 mMNaN0 (data from ref (41).



84

li/te

Ligand-Like

Metal Ion-Like

Ternary Complex Formation Figure 11. Ligand-like and metal ion-like ternary complex formation.


85 We have completed an adsorption study i n systems containing FeOOH(goethite), representative phosphonate ligands, and metal ions likely to be encountered in wastewater effluents (42). Transition metal ions are typically encountered at concentrations comparable to those of phosphonate ligands. 10 μΜ Fe exerts a very slight diminishing effect on the adsorption of 10 μΜ E D T M P and D T P M P , but no discernable effect on the adsorption of A M P , IDMP, HEDP, and N T M P . Cu and Ζη (also at 10 μΜ) exerted no discernable effect. C a is a major ion in most waters; concentrations from 50 μΜ to 1.0 m M significantly increased N T M P adsorption (42). Examining how phosphonates affect the adsorption of metal ions is, of course, important too. A t p H 4.0, 10 μΜ N T M P , E D T M P , and D T P M P substantially raise the extent of 10 μΜ Cu adsorption. "Ligand-like" ternary complex formation and the lowering of surface charge are believed responsible (42). A t p H values greater than 7, 10 μΜ N T M P , E D T M P , and D T P M P substantially lower the extent of 10 μΜ Cu adsorption. Here, the formation of Cu -phosphonate complexes in solution draws Cu ion away from the FeOOH(goethite) surface. Weaker chelating agents such as A M P , H E D P , and IDMP exhibited only slight effects on Cu adsorption (42). Poor environmental practices in past decades have left a legacy of toxic metal-contaminated sediments i n many rivers, lakes, and estuaries. It would be desirable to determine whether inputs of synthetic chelating agents can solubilize toxic metal ions from such sediments. Working with sediments downstream from a copper electrorefinery plant, Bordas and Bourg (44) determined that greater than 100 μΜ concentrations of E D T A and N T M P were able to solubilize 10 % or more of particle-bound Cd, Pb, and C u within a 30hour time period. Wastewater effluents typically contain chelating agent concentrations that are tens- to hundreds-of-times lower (45). m

11

π

n


11

11

n

11

11

Chemical Transformations Preceding sections have pointed out ways in which the identity, number, and arrangement of Lewis Base functional groups within chelating agent molecules affect the complex formation and adsorption reactions of phosphonates. Here, we provide a few illustrative examples of how these properties affect the chemical reactivity of phosphonates. F e - E D T A and other F e complexes with (amino)carboxylate chelating agents readily photolyze, whereas free (amino)carboxylates and their complexes with common +11 metal ions do not (46). Indeed, the kinetics of ¥q capture and release control ambient concentrations of Fe -(amino)carboxylate complexes, and hence control degradation rates in surface waters (25). m

m

m

ffl


86 ffl

(Amino)phosphonate chelating agents behave i n the same way. Fe -EDTMP, for example, readily photolyzes, while free E D T M P does not (47). In the laboratory, photolysis of Fe -EDTMP yields orthophosphate ion and N-methyl H M P . Intermediates that chelate F e (e.g. N-methyl IDMP) are themselves rapidly photolyzed (47). By analogy, N T M P and most of the phosphonatecontaining chelating agents discussed i n this chapter should also be subject to photolysis. It is also important to identify chemical breakdown processes i n environments where photolysis cannot take place, e.g. deep portions of lakes and rivers where sunlight does not penetrate, as well as within soils and aquifers. In carefully prepared transition metal ion-free solutions, chemical breakdown of a wide range of (amino)phosphonates (e.g. by hydrolysis) was found to be very slow, requiring timescales of 100 days or more (48). Using hydrogen peroxide concentrations in the 0.1 m M range and comparable concentrations of Fe , appreciable degradation was observed (48). The exceedingly low hydrogen peroxide concentrations found i n actual groundwater samples (e.g. 20 n M reported by (49)) are probably not enough to induce degradation. The complete degradation of N T M P into IDMP, H M P , and A M P within approximately 48 hours in non-illuminated surface and ground waters has been reported (50). Comparable degradation was observed in a laboratory test medium that contained a number of inorganic anions ( N 0 \ CI", S0 ", B(OH) " , M o 0 " , and E D T A ) and metal ions (Na , Mg , Ca , Mn , Co , Cu , Z n , and Fe ). N T M P undergoes degradation in non-illuminated Baltimore City tap water. A 70 % loss was observed over a 7-day period, corresponding to a halflife of approximately 4 days (51). A series of single metal ion experiments indicated that ambient concentrations of M g and Ca and 10 μΜ concentrations of Cu , Zn , and Fe yielded no discernable degradation of N T M P . Rapid abiotic degradation was observed i n 0 -saturated solutions containing 10 μΜ Μη ; no degradation was observed i n 0 -firee solutions. It can be concluded that Mn -catalyzed autooxidation is responsible for N T M P degradation. The reaction scheme is presented in Figure 12. Interested readers should refer to Nowack and Stone (51) for details regarding reaction product identification, the effects of pH and medium composition on reaction rates, and comparisons with other phosphonate-containing ligands. The mechanism of Mn -catalyzed N T M P autooxidation is unique i n several respects. Although the oxidation of Mn (aq) by 0 is thermodynamically favorable at p H values greater than 4.7 (52), reaction ffl


m

11

2

3

2

1

11

11

4

11

11

4

11

n

4

111

11

11

11

11

111

2

11

2

n

n

2+

2



11

Figure 12. Mechanism of Μη -catalyzed

NTMP autoxidation (adapted ref.

(51)).


88 kinetics in the absence of added ligands is extremely slow below p H 8.5 (e.g. ref. (53)). The N T M P autoxidation reaction occurs at an appreciable rate across a wide p H range (4 < p H < 8.5), indicating that complexation by N T M P (i) broadens the p H range where M n reaction with 0 is thermodynamically favorable, and (ii) activates coordinated M n towards reaction with 0 . A s far as the first point is concerned, it is important to note that complexation by N T M P very likely lowers the reduction potential for the M n / M n halfreaction, since the logK for complexation of M n is larger than the logK for complexation of Mn . As far as the second point is concerned, relatively little is known about how the identity of coordinated ligands and composition of the aqueous medium affect rates of M n oxidation by 0 (54). In one study, E D T A and pyrophosphate were found to greatly inhibit the reaction at p H 10 (55). Whether or not this same result would be obtained under neutral or acidic conditions was not investigated. 11

2

11

2

ffl

n

m

11

11


2

Mn^chelating agent complexes can be synthesized in a variety of ways, and the intramolecular oxidation of the coordinated (amino)carboxylate chelating agents has been extensively studied (e.g. ref. (56-59)). Thus, once M n N T M P is formed by the autooxidation reaction, N T M P breakdown is quite reasonable. We can speculate that phosphonate-containing ligands might exist that promote the reaction of M n with 0 , but resist intramolecular oxidation by Mn . Soluble M n complexes generated i n this way would be strong oxidants in a thermodynamic sense, which might react quickly with other naturallyoccurring or contaminant-derived chemical species. D T P M P , E D T M P , and N T M P are all subject to Mn -catalyzed autooxidation, but H E D P and the breakdown products I D M P and F I D M P are not (51). Equilibrium speciation models indicate that the fraction of available M n complexed by H E D P and IDMP is small, which may limit the rate of the forward reaction. Properties of the alpha-amino group may also be important; we are currently investigating whether phosphonates containing secondary amine versus tertiary amine groups exhibit different susceptibilities towards oxidation. Of course, many more potential pathways for the chemical degradation of phosphonates need to be explored. The C-P bond in alkylphosphonates, for example, is weakened by electron-withdrawing substituents. Alpha-carbonyl phosphonates such as the antiviral pharmaceutical phosphonoformate are especially susceptible towards hydrolysis (60). Metal ion-catalyzed conversion of phosphonates into phosphonamides has been observed; nucleophilic attack of amine groups at the electrophilic phosphorus atom is believed responsible (61). m

11

2

111

111

n

11


89

Conclusions Synthetic organic compounds possessing a phosphonate group exist i n multiple protonation levels, coordinate dissolved metal ions, and adsorb onto inorganic solid surfaces. During and subsequent to use, knowledge of the speciation of phosphonate-containing compounds is crucial for understanding chemical transformations and environmental effects. Given the widespread use of a variety of phosphonate-containing chemicals, it is surprising how little is known about their thermodynamic properties (pK and logK values), and about their reactions with the chemical constituents of natural waters. When a study has been performed on one chemical within a class, it is commonly assumed that the entire class with exhibit comparable behavior. This review indicates, however, that very small changes i n structure (e.g., replacing one functional group with another) can markedly alter speciation, reaction pathways and reaction rates. A s i n most areas of environmental chemistry, phosphonate-containing synthetic organic compounds represent a rich new subject waiting to be explored.


a

Acknowledgements U.S. Environmental Protection Agency, National Center of Environmental Research and Quality Assurance (Office of Exploratory Research) support for our work on the environmental chemistry of phosphonates under Grant R826376 is gratefully acknowledged.

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