EDDS
Ethylenediamine-N,N'-disuccinic acid (EDDS) is an aminopolycarboxylic acid that functions as a strong, biodegradable chelating agent for transition metals, offering a sustainable alternative to the persistent synthetic chelator ethylenediaminetetraacetic acid (EDTA).[1] With the molecular formula C10H16N2O8 and a molecular weight of 292.24 g/mol, EDDS exists as a colorless to white solid with a melting point of 220–222 °C, low water solubility for the acid form though its salts exhibit high solubility in water, and forming stable complexes with metal ions such as iron, copper, and lead.[2][3] The (S,S)-isomer, known as [S,S]-EDDS, is the most commonly used stereoisomer due to its superior chelating efficiency and rapid biodegradability, achieving over 80% degradation within 28 days under aerobic conditions, unlike EDTA which persists in the environment.[1] EDDS is synthesized through the reaction of ethylenediamine with fumaric acid or maleic anhydride derivatives, often enzymatically or chemically to yield the desired stereochemistry, and is typically employed in its trisodium salt form (CAS 178949-82-1) for practical applications due to enhanced water solubility and pH stability above 9.0.[4][5][6] Its coordination chemistry involves forming hexadentate complexes with divalent and trivalent metals, providing effective sequestration in hard water and alkaline environments where EDTA performance diminishes.[3] Key applications include detergents and cleaning products, where it prevents metal-catalyzed oxidation and scale formation; environmental remediation, such as enhancing phytoextraction of heavy metals like lead and nickel from contaminated soils by increasing their bioavailability for plant uptake; and industrial processes like metal recovery from spent catalysts, achieving up to 84% extraction efficiency for nickel. As of 2024, advancements such as BASF's improved production process have enhanced its commercial viability.[7][8][9][10] Environmentally, EDDS demonstrates low toxicity to aquatic organisms, with EC50 values exceeding 1000 mg/L for fish and Daphnia, though algal growth inhibition (EC50 = 0.290 mg/L) arises from metal chelation rather than direct toxicity, resulting in predicted no-effect concentrations (NOEC) of 0.125–0.500 mg/L in field scenarios.[1] Its low sorption to sludge (Kp = 40 L/kg) and high removal rates (>96%) in sewage treatment plants ensure minimal accumulation in water bodies, with predicted environmental concentrations below 5 µg/L at typical usage levels, yielding a risk assessment ratio (PEC/PNEC) less than 1.[1] Approved for eco-labels and phosphorus-free formulations, EDDS supports green chemistry initiatives by reducing the ecological footprint of chelating agents in household and industrial products.[7]Structure and Properties
Molecular Structure
EDDS, or ethylenediamine-N,N'-disuccinic acid, is an aminopolycarboxylic acid with the chemical formula C₁₀H₁₆N₂O₈ and a molecular weight of 292.24 g/mol.[11] Its IUPAC name is 2-[2-(1,2-dicarboxyethylamino)ethylamino]butanedioic acid.[11] The molecular structure features an ethylenediamine backbone (H₂N-CH₂-CH₂-NH₂) where each nitrogen atom is substituted with a succinic acid group (-CH(COOH)-CH₂-COOH), resulting in four carboxylic acid functional groups.[12] This arrangement forms a hexadentate ligand, with the two tertiary amine nitrogens and four carboxylate oxygens serving as donor atoms for coordination.00082-9) The carbon chain connectivity can be represented as:where the nitrogens link the ethylenediamine to the alpha carbons of the succinic moieties.[2] EDDS possesses two chiral centers located at the alpha carbons of the succinic acid groups adjacent to the nitrogens.[11] These chiral centers give rise to three stereoisomers: the (S,S) enantiomer, the (R,R) enantiomer, and the meso (R,S) form.00082-9) The (S,S)-EDDS isomer, derived from L-aspartic acid configurations, is the primary biologically relevant form due to its enhanced biodegradability compared to the other isomers.00082-9)HOOC-CH₂-CH(N-CH₂-CH₂-N-CH-CH₂-COOH)-COOH | | COOH COOHHOOC-CH₂-CH(N-CH₂-CH₂-N-CH-CH₂-COOH)-COOH | | COOH COOH
Physical and Chemical Properties
EDDS is a white granular solid with no characteristic odor. It melts at approximately 220–222 °C, decomposing at this temperature. The free acid form exhibits low solubility in water (0.015 g/100 g at 20 °C) but its trisodium salt is highly soluble (>1000 g/L at 20 °C); overall, EDDS shows low solubility in organic solvents such as ethanol. As a polyprotic acid, EDDS possesses four ionizable protons with pKa values of 2.4, 3.9, 6.8, and 9.8, corresponding to the stepwise deprotonation of its carboxylic acid groups and the protonated amine group. At high pH values, the fully deprotonated EDDS⁴⁻ form predominates. EDDS demonstrates hydrolytic stability under neutral conditions but is subject to biodegradation.| Property | Value | Conditions |
|---|---|---|
| Appearance | White granular solid | - |
| Melting point | 220–222 °C (decomposes) | - |
| Water solubility (free acid) | 0.015 g/100 g | 20 °C |
| Water solubility (trisodium salt) | >1000 g/L | 20 °C |
| pKa values | 2.4, 3.9, 6.8, 9.8 | - |
Synthesis
Chemical Synthesis from Aspartic Acid
The chemical synthesis of (S,S)-ethylenediamine-N,N'-disuccinic acid (EDDS) from L-aspartic acid involves the alkylation of the amino group of L-aspartic acid with a 1,2-dihaloethane reagent, such as 1,2-dibromoethane, to form the ethylenediamine bridge, followed by cyclization and acidification to yield the disuccinic acid structure. This stereospecific process utilizes the (S)-configuration of L-aspartic acid to produce exclusively the (S,S)-isomer of EDDS, avoiding racemization under the reaction conditions.[4] The reaction proceeds without the need for amino group protection, as the basic medium deprotonates the amino functionality, enabling nucleophilic attack on the alkylating agent.[13] The process begins with the dissolution of L-aspartic acid in an aqueous medium, followed by addition of a strong base such as sodium hydroxide to achieve a pH of 9–13, typically 10–11, which facilitates the deprotonation.[4] A stoichiometric deficiency of 1,2-dibromoethane (molar ratio 0.1–0.45:1 relative to L-aspartic acid) is then introduced, and the mixture is heated to 80–120°C under ambient or slightly elevated pressure (up to 50 psig) for 4–9 hours, promoting the formation of the ethylene bridge and subsequent cyclization to the disuccinate.[4] The resultant basic solution is co-fed with a mineral acid, such as hydrochloric acid (2–40 wt%), into water while maintaining a pH of 2–6.5 (preferably 2.6–5.0) to protonate and precipitate the EDDS free acid.[4] Alternative alkylating agents, like 1,2-bissulfooxyethane, have been employed in refined variants to enhance selectivity, operating at pH 8–11 and temperatures up to 100°C.[13] This synthetic route was first reported in 1968 by Neal and Rose, who described the reaction of L-aspartic acid with 1,2-dibromoethane in the presence of base, achieving a modest yield of approximately 25% due to significant side products such as oligomers and hydroxyethylamine derivatives. Subsequent developments in the 1990s addressed scalability challenges by optimizing reagent ratios and pH control during acidification, reducing by-product formation (e.g., 2-bromoethylamine N-succinic acid at 5–20 wt%) and improving overall conversion to 30–50% with selectivity exceeding 83%.[4] Advanced protocols using cyclic sulfates as alkylating agents have further boosted yields to 79% for the free acid form.[13] Purification typically involves filtration of the precipitated EDDS, followed by washing with water and crystallization from aqueous media to achieve purity levels of 98–99 wt%.[4] While effective, the chemical route faces industrial limitations from the formation of halogenated waste and moderate yields (typically 50–70% after purification), prompting exploration of biotechnological alternatives for more sustainable production.[13]Biotechnological Synthesis
Biotechnological synthesis of (S,S)-ethylenediamine-N,N'-disuccinic acid (EDDS) relies on enzymatic catalysis using EDDS lyase, an enzyme originally isolated from the actinomycete Amycolatopsis japonicum. This lyase facilitates the stereospecific condensation of ethylenediamine and fumaric acid, yielding the biodegradable (S,S)-EDDS isomer with stereoselectivity greater than 99%. The reaction proceeds under mild aqueous conditions, typically at neutral pH and ambient temperatures, avoiding the harsh acids and high pressures required in traditional chemical routes.[14] To enhance efficiency and scalability, the EDDS lyase gene has been cloned and expressed in heterologous hosts such as Escherichia coli, enabling production of the enzyme for in vitro catalysis. Recent innovations include immobilization of the fumarase-free EDDS lyase on glutaraldehyde-activated amino carrier, which allows for enzyme reuse across multiple reaction cycles—up to 11 batches with 94% conversion—and achieves space-time yields of 1.55 g/(L·h). In optimized enzymatic systems, concentrations as high as 209 g/L have been reported, demonstrating the potential for high-titer production.[15][16] Fermentation-based approaches utilize engineered microorganisms expressing the EDDS biosynthetic gene cluster (aesA-H), with A. japonicum strains serving as the primary host due to their native pathway. These processes employ fed-batch cultivation using glucose as the carbon source and supplemental precursors like ethylenediamine, yielding titers up to 20 g/L while being sensitive to trace zinc inhibition. Heterologous expression in E. coli has also been explored for whole-cell biocatalysis, further broadening platform options for scalable production.[17][18] These methods offer key advantages over chemical synthesis, including superior stereoselectivity for the environmentally degradable (S,S)-form, reduced energy consumption, and minimized waste from byproducts. Post-2010 advancements in enzyme immobilization and metabolic engineering have improved reusability and process economics. Industrial-scale biotechnological production has been developed since the 1990s, supporting commercial applications in eco-friendly chelating agents. As of June 2025, further engineering of EDDS lyase through site-saturation mutagenesis has enhanced its activity for stereoselective synthesis of related compounds.[16][19]Coordination Chemistry
Chelation Mechanism
EDDS acts as a hexadentate chelating agent, coordinating metal ions through two secondary amine nitrogen donors and four oxygen atoms from its carboxylate groups.[20] This multidentate binding enables the formation of stable octahedral complexes with transition metals such as Fe³⁺, Cu²⁺, and Zn²⁺, where the ligand wraps around the metal center to maximize coordination.[20] The ethylenediamine backbone and succinyl side chains position the donor atoms optimally for chelation, similar to related aminopolycarboxylic acids.[21] The binding geometry involves the creation of five- and six-membered chelate rings, with the ethylenediamine moiety forming a five-membered ring via the two nitrogen atoms and the metal, while the succinate arms contribute six-membered rings through carboxylate oxygens.[22] These ring sizes facilitate efficient orbital overlap and steric accommodation in the octahedral arrangement, enhancing the overall chelation process.[23] In contrast to EDTA, which shares a comparable ethylenediamine core but uses acetate arms, EDDS's succinate extensions provide a more flexible structure that supports biodegradability without compromising the fundamental binding motif.[21] Chelation by EDDS exhibits pH dependence, with effective metal binding occurring above pH 6, where partial deprotonation of the carboxylic acid groups exposes the negatively charged oxygens for coordination.[20] At lower pH values, protonation competes with metal ions for the donor sites, reducing complex formation.[20] Spectroscopic techniques confirm the chelation mechanism, particularly through UV-Vis absorption shifts that arise from ligand-to-metal charge transfer upon complexation.[20] For instance, Fe³⁺-EDDS complexes display characteristic bands in the UV region, indicating electronic transitions associated with the coordinated structure.[24]Stability Constants and Complex Formation
EDDS forms stable 1:1 metal-ligand (ML) complexes with various metal ions, primarily through hexadentate coordination involving its two nitrogen atoms and four carboxylate groups. The formation constant K_f for these complexes is defined by the equilibrium \mathrm{M^{2+} + EDDS^{4-} \rightleftharpoons M(EDDS)^{2-}} with K_f = \frac{[M(EDDS)^{2-}]}{[\mathrm{M^{2+}}][EDDS^{4-}]}, where the charges adjust based on the metal ion valence.[20] Stability constants, expressed as log K values, indicate the strength of these complexes and vary by metal ion. For the [S,S]-stereoisomer of EDDS under standard conditions (25°C, ionic strength 0.1 M NaCl or equivalent), representative values include log K = 20.04 for Fe(III), 18.3 for Cu(II), 13.15 for Zn(II), and 8.69 for Mn(II). These constants were determined for aqueous solutions and reflect the ligand's affinity for transition and heavy metals. Similar values from independent studies confirm log K \approx 22.0 for Fe(III), 18.4 for Cu(II), and 13.4 for Zn(II) at 25°C and ionic strength 0.1–0.15 M.[25][11][11] The effective or conditional stability constants are pH-dependent due to protonation of EDDS and potential hydrolysis of metal ions. For Cu(II)-EDDS, the conditional log K peaks at pH 7–9, where the complex achieves maximum stability (over 90% formation), while at lower pH (<4) protonated species like CuHEDDS dominate, reducing effective binding. Fe(III)-EDDS maintains high stability across pH 2–10, but hydrolysis risks increase above pH 10. These pH effects are critical for applications requiring neutral conditions.[25][26] EDDS exhibits selectivity for transition metals over alkaline earth metals, with log K values for Ca(II) and Mg(II) around 6–8, significantly lower than for Cu(II) or Zn(II). This preference arises from the ligand's ability to form more stable chelate rings with softer Lewis acids. At high pH (>10), hydrolysis of trivalent metals like Fe(III) can compete with complexation, lowering effective stability.[11][25] Experimental determination of these constants relies on potentiometric titration, where pH is monitored during addition of base to solutions containing EDDS and metal ions. Titrations are conducted at controlled temperatures (e.g., 25°C) and ionic strengths (0.1 M), with data analyzed using programs like SUPERQUAD or ESTA to fit equilibrium models. Metal-to-ligand ratios (1:1 to 1:4) ensure identification of predominant 1:1 species, with glass electrodes calibrated for accurate H⁺ activity.[25][20]| Metal Ion | log K (ML) | Conditions | Source |
|---|---|---|---|
| Fe(III) | 20.04 | 25°C, I = 0.1 M NaCl | Hyvönen (2008)[25] |
| Cu(II) | 18.3 | 25°C, I = 0.1 M NaCl | Hyvönen (2008)[25] |
| Zn(II) | 13.15 | 25°C, I = 0.1 M NaCl | Hyvönen (2008)[25] |
| Mn(II) | 8.69 | 25°C, I = 0.1 M NaCl | Hyvönen (2008)[25] |