Degradability of Iron(III)-aminopolycarboxylate Complexes in Alkaline

Alkaline Media: Statistical Design and X-ray Photoelectron. Spectroscopy Studies .... level design (pH ) 8, 9, 10; T ) 25, 40, 55 °C; I ) 0.025,. 0.1...
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Ind. Eng. Chem. Res. 2005, 44, 5053-5062

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Degradability of Iron(III)-aminopolycarboxylate Complexes in Alkaline Media: Statistical Design and X-ray Photoelectron Spectroscopy Studies Simon Piche´ and Faı1c¸ al Larachi* Department of Chemical Engineering, Laval University, Que´ bec, Canada G1K 7P4

Use of ferric-aminopolycarboxylate complexes for odor control via the oxidative scrubbing of H2S and CH3SH contained in pulp and paper noncondensable gas emissions is evoked as potentially beneficial from the standpoint of iron-sequestration and protection against precipitation in the alkaline environments characteristic of the Kraft mill sulfate-pulp processes. In this study, the degradability of two ferric-aminopolycarboxylate complexes in alkaline solutions was investigated by means of replicated four-factor three-level fixed-effects completely randomized factorial (2 × 34) designs. Expressed in terms of ferric-ethylenediaminotetraacetate (Fe3+EDTA4-) and ferric-trans-1,2-diaminocyclohexanetetraacetate (Fe3+CDTA4-) daily degradation rates, the degradability response was monitored via UV-vis spectrophotometry as a function of temperature (T ) 25, 40, 55 °C), alkalinity (pH ) 8, 9, 10), ionic strength (I ) 0.025, 0.1, 0.5 M) and ferric concentration (CFe ) 175, 280, 450 µM). Analysis-of-variance (ANOVA) of the factorial design suggests that pH and temperature are the main factors increasing Fe3+EDTA4- and Fe3+CDTA4- degradation rates. To a lower extent, ionic strength and ferric chelate concentration also promote degradation. At the most severe factor-level combinations (T ) 55 °C, pH ) 10, and I ) 0.5 M), up to 40% of Fe3+CDTA4- and 54% of Fe3+EDTA4- degraded after 1 day, confirming that CDTA is a superior chelating agent against iron precipitation in alkaline solutions. The brownish fresh-state Fe3+EDTA4- or Fe3+CDTA4- solutions evolved with degradation into turbid solutions whereof the precipitated solid was recovered and its surface probed through X-ray photoelectron spectroscopy (XPS). XPS revealed that the solid degradation product was inorganic and mostly contributed by Fe(OH)3. It was, however, not possible to identify which one of the organometallic complex degradation or the ferric dechelation was responsible for iron(III) hydroxide formation since both routes can contribute to its formation. 1. Introduction Synthetic aminopolycarboxylic acids (e.g., nitrilotriacetic acid (NTA), ethylenediaminotetraacetic acid (EDTA), and trans-1,2-cyclohexanediaminotetraacetic acid (CDTA))1-2 are used to prevent the inherent precipitation in moderately acid to alkaline solutions of some multivalent ions of the first transition metal series (e.g., Fe). Several redox reactions involving iron cations (Fe3+/Fe2+) are indeed accelerated in alkaline media. For instance, the use of iron chelate systems (i.e., Fe3+Ln-/ Fe2+Ln-, where Ln- denotes an anionic organic ligand with n- charges) to remove hydrogen sulfide (H2S) from sour gas and water effluent streams requires 2 mol of hydroxide for every mole of oxidized H2S (eq 1).3-10 The dissolved H2S also dissociates in alkaline solutions to form successively HS- (pKa ) 7.1 at 25 °C) and S2- (pKa ≈ 14 at 20 °C),11 which both promote the absorption rate of hydrogen sulfide during scrubbing.12

H2S(aq) + 2Fe3+Ln- + 2OH- f S0 + 2Fe2+Ln- +2H2O (1) One main problem afflicting industrial application of such processes is the magnitude of iron precipitation. * To whom correspondence should be addressed: Tel.: (418) 656-3566. Fax: (418) 656-5993. E-mail: faical.larachi@ gch.ulaval.ca.

Studies have been conducted to unveil the causes and understand the mechanism of ligand degradation during H2S removal in aqueous streams. For instance, Chen and co-workers13,14 observed that polyaminocarboxylic acids tend to rupture at the weakest locations, for example, ethylene moiety of EDTA.13,14 Cleavage is presumably ascribed to the presence of hydroxyl free radicals produced from the reoxidation of the Fe2+Lnproduct into active Fe3+Ln- via a Fenton mechanism (eqs 2 and 3). Addition of free radical scavengers (i.e., hyposulfite) in slightly alkaline solutions is found to reduce significantly the rate of ligand degradation.

2Fe2+Ln- + O2 + 2H2O f 2Fe3+Ln- + H2O2 + 2OH(2) Fe2+Ln- + H2O2 f Fe3+Ln- + OH- + OH•

(3)

Another concern regarding iron precipitation is related to the thermodynamic equilibrium constant, KML, of the Fe3+/Fe2+ chelate formation. Also known as the stability constant, KML characterizes the iron chelate capability to bind without splitting into distinctive species (eq 4). For example, knowledge of ligand deprotonation constants (e.g. KD

EDTA 798 EDTA4-

10.1021/ie049636f CCC: $30.25 © 2005 American Chemical Society Published on Web 08/18/2004

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KD1-KD4) and ferric hydroxide dissociation constants (i.e. KS

Fe(OH)3 798 Fe3+ KS1-KS3) is required to establish the proportion of Fe3+EDTA4- as a function of pH. KML

Fe3+ + EDTA4-798Fe3+EDTA4-

(4)

Based on the work by Smith et al.,2,15,16 trivalent ions of the first transition metal series (i.e., Mn3+, Fe3+) combined with aminopolycarboxylates yield highly stable chelates with log(KML) well over 20. For the valence III iron chelates involved in this study (log(KML) ) 25 for Fe3+EDTA4-; log(KML) ) 30 for Fe3+CDTA4-), dechelation should be nonexistent throughout the normal pH range. The above literature review suggests that since ferric chelates of amino carboxylic acid exhibit high stability constants while supposedly the iron chelates may degrade only in the presence of free radicals, one should expect that Fe3+EDTA4- and Fe3+CDTA4- would undergo neither dechelation nor organic degradation in H2S-free alkaline solutions. Yet, preliminary observations appear to contradict this conclusion. Slow degradation (or dechelation) rates were monitored using samples of Fe3+CDTA4- and Fe3+EDTA4- stock solutions mixed into alkaline pH buffers. An investigation using replicated completely randomized factorial designs was initiated to assess the extent of degradability experienced by buffered Fe3+EDTA4- and Fe3+CDTA4solutions through variations of solution pH, temperature, and ionic strength. As a complement, recovery of the degradation solid from the resulting turbid aqueous solutions was analyzed using X-ray photoelectron spectroscopy (XPS) to gain some knowledge on the type of ferric chelate loss, whether it is organic degradation or dechelation. 2. Factorial Design and Analysis A full factorial design featuring four fixed-effects factors (pH, temperature (T), ionic strength (I), and ferric chelate concentration (CFe)) was used with the goal to quantify the level of iron loss from Fe3+CDTA4- and Fe3+EDTA4- aqueous solutions (see Figure 1). A threelevel design (pH ) 8, 9, 10; T ) 25, 40, 55 °C; I ) 0.025, 0.1, 0.5 M; CFe ) 175, 280, 450 µM) was implemented to illustrate the nonlinear behavior of ferric chelate degradation within the interrogated factorial space. Duplication of the 81 (34) treatment combinations for each ferric chelate was used for experimental error assessment. The factorial boundaries were imposed according to the following specifications. Slightly alkaline conditions (8 < pH < 10) promote both H2S absorption and H2S + Fe3+Ln- reaction rates.3 Gas absorption of total reduced sulfur species, and particularly H2S, from typically warm noncondensable gases would normally increase the scrubbing liquid temperature to levels usually between 20 °C and 60 °C. Untreated water streams, used for gas scrubbing, usually contain inorganic ions in amounts that could reach relatively high ionic strengths. Ferric chelate degradation was monitored by UV-vis spectrophotometry. The maximum concentration boundary (450 µM) was dictated by the UV-light molar

Figure 1. Chemical structures of ferric trans-1,2-diaminocyclohexanetetraacetic acid (Fe3+CDTA4-) and ferric ethylenediaminotetraacetic acid (Fe3+EDTA4-).

absorbtivity of ferric chelates. Higher concentrations could not be measured precisely with the instrument. 2.1 Experimental Procedure. Quantitative measurement of ferric chelates was performed with a UVvis spectrophotometer (Varian Cary 300 model). As describedpreviously,17 twoferricCDTAspecies(Fe3+CDTA4and Fe3+OH-CDTA4-) are involved in slightly alkaline solutions (pKa ) 9.7). A similar arrangement was expected for ferric EDTA (pKa ) 7.2).18,19 Each species generates different UV absorbance spectra in the 225325 nm spectral range. On behalf of consistency, calibration measurements were acquired to determine spectrum-to-concentration relationships for individual ferric EDTA species (Fe3+EDTA4- and Fe3+OH-EDTA4-). The calibration measurements for ferric CDTA species were readjusted from our previous work.17 Globally, two sets of 15 [225-325]-nm spectra for different Fe3+L4(or Fe3+OH-L4-) concentrations from 10 to 500 µM were processed for each one of EDTA and CDTA using PLSplus IQ principle component regression (PCR) analysis from Thermo Galactic Grams/32 AI software. Calibration measurements were done at pH boundaries where 100% of the nonhydroxylated species (pH ) 4) and 100% of the hydroxylated species (pH ) 12) were observed. All species’ UV absorbance follows the BeerLambert law perfectly with molar absorbtivities (i.e., at 260 nm) of 8300 L/mol for Fe3+CDTA4-, 7615 L/mol for Fe3+OH-CDTA4-, 8055 L/mol for Fe3+EDTA4-, and 7200 L/mol for Fe3+OH-EDTA4-. Ferric chelate degradation experiments were performed in 150-mL glass bottles in which 125 ( 2 mL of a pH buffer solution (Fisher Scientific) was added (pH ) 8.00 ( 0.05: Potassium phosphate monobasic, sodium hydroxide buffer; pH ) 9.00 ( 0.05 and 10.00 ( 0.05: Potassium chloride, boric acid, sodium hydroxide buffers). Sodium chloride (99%, EMD) was dissolved into the buffer solution to reach the desired ionic strength (mNaCl ) 0.18 ( 0.01 g for I ) 0.025 ( 0.001 M; mNaCl ) 0.73 ( 0.01 g for I ) 0.100 ( 0.002 M; mNaCl ) 3.66 (

Ind. Eng. Chem. Res., Vol. 44, No. 14, 2005 5055 Table 1. Averaged (Six-Value) ( Standard Deviation on the Relative Degradation Rate (∆CFe/CFe) Rate in Terms of pH, Temperature (T), and Ionic Strength (I) I ) 0.025 M pH

CDTA (%)

I ) 0.1 M

EDTA (%)

CDTA (%)

I ) 0.5 M EDTA (%)

CDTA (%)

EDTA (%)

1.2 ( 0.2 8.0 ( 1.1 8.7 ( 1.0

1.3 ( 0.2 2.9 ( 0.4 10.5 ( 1.5

1.8 ( 0.5 7.2 ( 0.9 14.6 ( 0.7

8 9 10

0.5 ( 0.3 2.8 ( 0.5 7.4 ( 1.1

1.0 ( 0.6 8.2 ( 0.7 7.6 ( 0.4

T ) 25 °C 0.8 ( 0.3 3.2 ( 0.4 8.7 ( 1.8

8 9 10

3.5 ( 0.4 3.5 ( 0.2 7.3 ( 1.1

1.4 ( 0.6 7.2 ( 0.3 11.5 ( 1.2

T ) 40 °C 3.6 ( 0.4 4.2 ( 0.7 10.6 ( 1.1

1.4 ( 0.5 7.0 ( 0.7 17.3 ( 0.8

3.3 ( 0.6 5.5 ( 0.6 11.6 ( 0.9

1.9 ( 0.7 9.0 ( 1.2 35.4 ( 1.9

8 9 10

2.5 ( 0.6 5.5 ( 0.8 16.9 ( 1.2

3.1 ( 0.9 8.3 ( 0.4 18.4 ( 2.5

T ) 55 °C 2.5 ( 0.4 6.5 ( 0.6 23.6 ( 2.3

2.8 ( 0.9 9.8 ( 0.9 27.8 ( 3.4

3.4 ( 0.8 9.2 ( 1.3 40.1 ( 1.8

3.4 ( 0.7 13.0 ( 1.6 53.7 ( 2.6

0.04 g for I ) 0.500 ( 0.010 M). Then, a specified volume of a ferric chelate (Fe3+CDTA4- or Fe3+EDTA4-) stock solution of 38.5 ( 0.6 mM (pH ≈ 3.5) was added to the bottle (VFe ) 0.57 ( 0.01 mL for CFe ) 175 ( 5 µM; VFe ) 0.91 ( 0.01 mL for CFe ) 280 ( 7 µM; VFe ) 1.46 ( 0.02 mL for CFe ) 450 ( 12 µM). Ferric CDTA and EDTA stock solutions were prepared using the procedure described previously.17 Due to the inherent variation of pH with temperature, correction of the trial solution pH was needed to obtain the required pH at temperatures of 40 °C and 55 °C. The experimental pH was adjusted at 25 °C with a 1 M NaOH solution using the following scheme: pH(T ( 1 °C) ( 0.05 ) pH(25 ( 1 °C) ( 0.05-8.00(40 °C) ) 8.07 (25 °C), 8.00 (55 °C) ) 8.12 (25 °C), 9.00 (40 °C) ) 9.12 (25 °C), 9.00 (55 °C) ) 9.22 (25 °C), 10.00 (40 °C) ) 10.15 (25 °C), 10.00 (55 °C) ) 10.40 (25 °C). A maximum of 4 mL of the NaOH solution was added per trial, which is considered low enough to be neglected in subsequent calculations. After pH adjustment and thorough mixing, a sample of the ferric chelate buffer solution was extracted from the bottle with a peristaltic pump (Varian Routine Sampler Accessory) and conveyed to the spectrophotometer inline cavity for reading. The time was noted. The bottle was then placed into an oven at the preset temperature (25 °C, 40 °C, or 55 °C). After 24 ( 0.5 h of incubation, another sample was analyzed. 2.2 Data Analysis. Spectral analysis was performed with previously established PCR calibration models. AdditionoftherespectiveFe3+CDTA4- andFe3+OH-CDTA4concentrations gives the total ferric CDTA concentration. The same analogy applies for ferric EDTA. For the sake of clarity, this paper will report the concentration difference between the first sample (t ) 0 h) and the last sample (t ) 24 ( 0.5 h), referred to here as the daily degradation rate (∆CFe, µM/day). Tables S1-S6 (in Supporting Information) give the raw experimental measurements in terms of absolute daily degradation rate (µM/day) as well as relative degradation rate (∆CFe/ CFe, day-1). Table 1 summarizes the relative degradation rate average for each treatment combination involving temperature, pH, and ionic strength. Accordingly, the six-value average combines the measurements for each duplicated concentrations (CFe ) 175, 280, and 450 µM). The following observations, consistent for both ferric chelates, are noted. Increasing the solution’s pH greatly increases the amount of Fe3+CDTA4- and Fe3+EDTA4- degradation especially at pH ) 10. Rising temperature increases the degradation rate. It is notably apparent at pH ) 10.

Figure 2. Evolution of ferric chelates daily degradation rate (∆CFe) with pH (0, 9 represent Fe3+CDTA4-; O, b represent Fe3+EDTA4-). Plotted values correspond to the average of duplicated treatment combinations.

Ionic strength slightly increases the degradation rate. Ingeneral,Fe3+CDTA4- ismorestablethanFe3+EDTA4-. The difference is more apparent at pH ) 9 (T ) 25-55 °C) and pH ) 10 (T ) 40-55 °C). Figure 2 clearly shows that pronounced alkalinity afflicts stability whether EDTA or CDTA is used as the chelating agent. The degradation rate intensifies after pH ) 9. As a result, Fe3+CDTA4- and Fe3+EDTA4degradation can be expected to achieve high rates (∆CFe/ CFe > 50%) after only 1 day for pH g 10. Experimental values overlapping for different temperatures (i.e., T ) 40 °C, 55 °C) suggest a possible link between pH and temperature effects. Based on Figure 2, the influence of pH on the degradation rate accelerates (i.e., the slope increases) as temperature is increased. In agreement, chelate degradation rates accelerate within the temperature range despite some apparent experimental inaccuracy (Figure 3). Although not apparent for Fe3+EDTA4-, it seems that an increase of the solution’s ionic strength intensifies the temperature effect on Fe3+CDTA4- degradation. As alleged previously, ionic strength boosts the degradation rate especially for Fe3+EDTA4- solutions (Figure 4). It becomes apparent when operating conditions involve high pH and temperature. The ferric chelate concentration factor is also expected to influence ∆CFe linearly since ∆CFe/CFe remained fairly constant for the same T, pH, and I treatment combination, see

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Figure 3. Evolution of ferric chelates daily degradation rate (∆CFe) with temperature (T) (0, 9 represent Fe3+CDTA4-; O, b represent Fe3+EDTA4-).

Figure 5. Evolution of ferric chelates daily degradation rate (∆CFe) with concentration (CFe) (0, 9 represent Fe3+CDTA4-; O, b represent Fe3+EDTA4-). Table 2. Sum of Squares (SS) and F Values Based on the Full Factorial Analysisa effects

DF

average 1 T 2 pH 2 I 2 CFe 2 T × pH 4 T×I 4 T × CFe 4 pH × I 4 pH × CFe 4 I × CFe 4 T × pH × I 8 T × pH × CFe 8 T × I × CFe 8 pH × I × CFe 8 pH × T × I × CFe 16 error 81 total 161

SS CDTA 3865.8 6.030 17.804 0.966 3.802 2.108 0.122 0.084 0.046 0.068 0.036 0.535 0.286 0.038 0.100 0.111 0.234 3898.2

EDTA 3689.2 3.299 31.791 1.592 3.418 0.682 0.028 0.188 0.503 0.123 0.003 0.246 0.115 0.055 0.017 0.080 0.647 3732.0

F value CDTA 1.34 × 106 1043.7 3081.5 167.2 658.0 182.4 10.6 7.3 3.9 i 5.9 i 3.1 i 23.1 12.4 1.6 i 4.3 i 2.4 i

EDTA 4.62 × 105 206.5 1990.0 99.7 214.0 21.3 0.9 i 5.9 i 15.7 3.8 i 0.1 i 3.8 i 1.8 i 0.9 i 0.3 i 0.6 i

F1% (DF, 81) 6.90 4.88 4.88 4.88 4.88 3.56 3.56 3.56 3.56 3.56 3.56 2.74 2.74 2.74 2.74 2.25

Figure 4. Evolution of ferric chelates daily degradation rate (∆CFe) with ionic strength (I) (0, 9 represent Fe3+CDTA4-; O, b represent Fe3+EDTA4-).

a Double barred values, F < F 1% (DF, 81); barred values, F < 2 × F1% (DF, 81).

Figure 5. The coupling effect of concentration and temperature is also demonstrated in the case of Fe3+EDTA4-, whereas it is almost nonexistent for Fe3+CDTA4- (parallel lines). 2.3 Factorial Analysis. All studied factors (pH, T, I, and CFe) exhibit more or less important main effects on Fe3+CDTA4- and Fe3+EDTA4- degradability. Even interaction effects (e.g., the binary interactions T × pH and T × I) appear to increase ferric chelate instability. The importance of main and interaction effects was assessed using the F values from the univariate modeling/analysis-of-variance (ANOVA) technique (Table 2). Both factorial designs (2 × 34 ) 162 measurements) were processed with the SPSS 10.0.0 software. Note that a logarithmic response, -log(∆CFe), and a type III sum of squares (SS) model were used in the ANOVA analysis. As expected, all four main effects contribute largely to the degradation rate, pH being the most influential of all. Temperature and ferric chelate concentration, although significant factors, induce much less degradation as such when compared to pH, while the ionic strength confirms its relatively low incidence on the

degradation rate. With the exception of the T × pH interaction for Fe3+CDTA4-, all joint effects may be considered irrelevant compared to the main factors. Despite a severe significance level (1%) on the F values, only few interactions could be discarded (F < F1% (DF, 81); 2 out of 15 for CDTA, 6 for EDTA, see Table 2). However, it should be noted that most interaction F values are close to F1% (DF, 81) and even if they remain statistically relevant, their contribution to the variance remains very marginal compared to the main factor effects. As an arbitrary choice, further simplification was therefore instigated with elimination of interactions exhibiting F values < 2 × F1% (DF, 81), thus allowing significant reduction in the complexity of the ANOVA models. Figure 6 exposes the potential of these simplifications.Withninemain/combinedfactorsforFe3+CDTA4(pH, T, I, CFe, T × pH, T × I, T × CFe, T × pH × I, and T × pH × CFe), the simplified ANOVA model assigns -log(CFe) values within an average absolute relative error (µ) of 1%. The residuals dispersion is satisfactory while properly following the normal distribution with the exception of 4-6 points that likely exhibit some

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Figure 6. Quality-of-fit description of the optimized ANOVA models (1, parity diagram; 2, residual plot; 3, normal quantile-quantile plot for the residuals).

experimental inconsistencies. Likewise, the simplified (six factors; pH, T, I, CFe, T × pH, and pH × I) ANOVA model for Fe3+EDTA4- expresses reasonably good predictability with µ ) 1.4%. Here, 80-90% of residuals follow a normal distribution. Model oversimplification

does not explain alone why the remaining 10-20% of measurements deviate from normality since the full factorial ANOVA model presented the same behavior. Indeed, most residual values lying outside the normal distribution (Figure 6) are specific to low degradation

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rate experiments. In such cases, the standard deviation remains close to the already low six-value average (i.e., µ ) 0.5 ( 0.3%, CDTA degradation at T ) 25 °C, pH ) 8.00, and I ) 0.025 M; see Table 1), which inflicts much more disparity between duplicated logarithmic responses, -log(∆CFe), and the ANOVA model prediction. Such divergence decreases for higher degradation rate experiments even if the absolute difference between duplications is larger. 3. Characterization of Degradation Product X-ray photoelectron spectra on three samples (S1: CDTA; S2: Fe3+CDTA4-; and S3: Fe3+CDTA4- degradation product) was acquired at room temperature on an AXIS ULTRA X-ray photoelectron spectrometer from Kratos Analytical using a monochromatic Al KR (hv ) 1486.6 eV) X-ray source operated at 300 W. The photoelectron kinetic energies were measured using a hemispherical electrostatic analyzer working in a constant pass energy mode. Survey scans (0-1150 eV) and high-resolution spectra on C 1s, N 1s, O 1s, and Fe 2p were acquired pass energies of 160 and 20 eV, respectively which correspond to nominal resolutions of 1000 and 50 meV. The studied samples were obtained as follows. First, pure CDTA (C14N2O8H22) powder labeled S1 was analyzed to give precise information on the organic molecule surface structure in terms of carbon (C 1s), nitrogen (N 1s), and oxygen (O 1s) core-level XP spectra. Assignment of feature spectral signatures (i.e., aliphatic carbon, β-C) will assist in determining whether, for instance, the full organic structure of CDTA is detected in the degradation product or not. Meanwhile, a 100-mL sample of the initial Fe3+CDTA4- stock solution (38.5 mM; pH ≈ 3.5) was slowly dehydrated at 50 °C for a couple of days to obtain a solid sample of ferric CDTA (S2). Due to the usually small quantities of precipitate formation during experiments, a series of 10 bottles providing highly degrading conditions (pH ) 10, T ) 55 °C, I ) 0.5 M, and CFe ) 450 µM) were incubated for 1 day. The precipitate was then filtered and carefully washed with a NaOH solution, thus avoiding the possible dissolution of Fe(OH)3. The degradation product (S3) was allowed to dry before analysis. XPS data reduction of the measured C 1s, N 1s, and O 1s high-resolution spectral envelopes for the CDTA (S1) and Fe3+CDTA4- (S2) samples was performed by curve fitting synthetic peak components using the software package CasaXPS (computer aided surface analysis for X-ray photoelectron spectroscopy, Casa Software Ltd., United Kingdom). Symmetric GaussianLorentzian product functions were used to approximate the line shapes of the fitting components. Peak synthesis was performed after Shirley background subtraction. The Fe 2p high-resolution spectra were also analyzed for Fe3+CDTA4- (S2) and the degradation product (S3) samples. Accounting for iron multiplet splitting, a tailing function with an asymmetry index of 0.2081 was added to the Gaussian-Lorentzian product functions. 3.1 Survey Scans. The relative atomic surface concentration based on survey scans is given in Table 3. Primarily, it shows a C:O:N atom ratio of 14:6:2. This is uncharacteristic when compared to the expected 14: 8:2 stoichiometric ratio for a single CDTA molecule. It was speculated that the lack of two oxygen atoms could be ascribed to possible formation of a CDTA dimer by removing two water molecules (Figure 7), which are

Figure 7. Speculated CDTA dimerization (--- carbonyl condensation, s amine condensation). Overall reaction: nC14O8N2H22 + nH+ T -(C14O6N2H19)n- + 2nH2O. Table 3. Atomic Concentration (at. %) for Each Studied Sample atom

CDTA

Fe3+CDTA4-

degraded product

carbon (C 1s) oxygen (O 1s) nitrogen (N 1s) iron (Fe 2p) sodium (Na 1s) xhlorine (Cl 2s) C/O/N/Fe

63.7 27.0 9.0 0.3 14:5.9:2:0

55.8 25.9 7.2 3.6 4.7 2.8 14:6.5:1.8:0.9

12.4 31.9 1.6 10.9 19.7 23.5 14:36:1.8:12.3

regained anew during CDTA dissolution. Ferric CDTA atom ratios are close to those for pure CDTA, with the exception that the ferric chelate also contains iron with an expected C-to-Fe ratio of 14 to 1. Sample S2 also contains sodium and chloride issued from NaOH and FeCl2 precursorsusedinthepreparationoftheFe3+CDTA4stock solution. The atom fractions of oxygen and iron clearly increased in the degradation product sample S3, suggesting that a large proportion of CDTA carbon and amine (i.e., the diaminocyclohexane ring) was purged during filtration while most of the iron was recuperated (alkaline pH). The presence of carbon is presumed here to be a remnant of Fe3+CDTA4- that had been left absorbed on the precipitate. Assuming that the residual Fe3+CDTA4- exhibits a stoichiometry given by C:O:N: Fe ) 14:8:2:1, a new ratio characterizing the remaining fraction in S3 is calculated to support the Table 3 results. The new corrected ratios (C:O:N:Fe ≈ 0:28:0: 11) fit the stoichiometry of Fe(OH)3 reasonably well. Therefore, preliminary assumptions that most of the degraded product is inorganic (ca. 92%) can be advanced based strictly on the XPS survey scans. This point will be discussed next in detail through the fitting of the different C, N, O, and Fe assignments. 3.2 XPS Characterization of CDTA. A summary of the chemical structure and assessment of the quality of fit for the C 1s, N 1s, and O 1s core electrons’ highresolution spectra is given in Table 4 in terms of binding energy (BE) shift, peak full width at half-maximum, and chemical environment assignment. The spectral correction for charging effects was referred to the binding energy of aliphatic carbon (β-C) at 285.1 eV.20 It resulted

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Figure 9. High-resolution XP N 1s spectra (S1: CDTA and S2: Fe3+CDTA4-) along with the fitted elemental spectra for pure CDTA. Elemental spectra assignments (a, b) are given in Table 4.

Figure 8. High-resolution XP C 1s spectra (S1: CDTA; S2: Fe3+CDTA4-; and S3: CDTA degradation product) along with the fitted elemental spectra for pure CDTA. Elemental spectra assignments (a-f) are given in Table 4. Table 4. XP Spectra Fit Assessment for C 1s, N 1s, and O 1s Core Electrons with Peak Assignment for the CDTA Sample (S1) peak position fwhm (eV) (eV) C 1s a b c d e f N 1s a b O 1s a b c

285.1 286.1 286.7 287.4 288.7 289.4 399.7 402.0 531.4 532.2 533.7

1.15 1.12 1.00 1.40 1.00 1.24 1.33 1.60 1.30 1.98 1.54

assignment (at oms/molecule) C-C-C N-C-C N-C(-C)-C N+-C()O)-C C-C(OOH) C-C(OO-C) C-N(-C)-C C-(C-)N+(-C)-C C(dO)-O-(dO) C(dO)OH C(dO)OH

(4/14) (4/14) (2/14) (1/14) (1/14) (2/14) (1/2) (1/2) (1/6) (4/6) (1/6)

fit assessment ∆BE ) 283-291 eV µ ) 10%, σµ ) 25%

∆BE ) 398-404 eV µ ) 15%, σµ ) 20% ∆BE ) 529-536 eV µ ) 7%, σµ ) 12%

Figure 10. High-resolution XP O 1s spectra (S1: CDTA; S2: Fe3+CDTA4-; and S3: CDTA degradation product) along with the fitted elemental spectra for pure CDTA. Elemental spectra assignments (a-c) are given in Table 4.

in an overall correction of 2.2 eV for the CDTA sample. Based on the polymeric structure of CDTA speculated in Figure 7, six chemical environments for carbon, two for nitrogen, and three for oxygen are recognized (Table 4). Acceptable fit over XP spectra was obtained with constitutive peaks ranging from 285.1 to 289.4 eV for C 1s (Figure 8). Peak assignment was resolved through knowledge of the molecule stoichiometry (i.e., four aliphatic carbon atoms, C-CH2-C, present in the cyclohexane ring) combined with information previously given by Hamoudi et al.20 Figure 9 confirms beyond a doubt that two structural combinations exist for nitrogen in CDTA. One peak at 399.7 eV clearly represents the tertiary amine (C-N(-C)-C) while the second peak bursting at the higher BE (402.0 eV) characterizes the ammonium configuration (C-(C-)N+(-C)-C) with increased electropositivity.21 This peak clearly disappears during dissolution and formation of Fe3+CDTA4-. Figure

10 provides the XP O 1s spectra for the three samples. Carbonyl and hydroxyl oxygen were assigned based on previously obtained results20 while the carboxylic anhydride function, C(dO)-O-C(dO), formed during polymerization was assigned by default. 3.3 XPS Characterization of Fe3+CDTA4-. A summary of the fit for C 1s, N 1s, and O 1s core electrons’ high-resolution spectra is given in Table 5. Spectral correction for charging effects was 2.2 eV for the Fe3+CDTA4-. As expected, the organic characters exhibited by CDTA and Fe3+CDTA4- yield equivalent XP spectra with some exceptions (Figures 8-10). After dissolution and formation of Fe3+CDTA4-, some elemental structures allocated to the carbonyl function (C 1s, d-f, Table 4) disappear and form a new coordinated iron-carboxylate function, (C-C(dO)-O-Fe). The other carbon atoms involved in the cyclohexane and ethylene functions remain intact with more or less the same peak position, fwhm (full-width at half-maximum), and stoi-

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Table 5. XP Spectra Fit Assessment for C 1s, N 1s, and O 1s Core Electrons with Peak Assignment for the Fe3+CDTA4- Sample (S2)a peak position fwhm (eV) (eV) C 1s

N 1s

O 1s

a

285.1 285.8 286.1 286.6 287.2 288.7 398.5 400.0 401.3 531.7 532.0 536.5

1.10 1.00 1.17 1.01 1.00 1.43 1.50 1.39 1.50 1.35 2.00 1.73

assignment (at. %)

fit assessment

23.0 14.4 23.0 N-C-C N-C(-C)-C 11.5 5.1 23.0 C-C(dO)-O-Fe 4.6 C-(C-)N+(-Fe)-C 87.7 7.7 47.3 C(dO)-O-Fe 47.3 C(dO)-O-Fe 5.4

∆BE ) 283-291 eV µ ) 6%, σµ ) 18%

C-C-C

∆BE ) 398-403 eV µ ) 9%, σµ ) 13% ∆BE ) 529-538 eV µ ) 11%, σµ ) 25%

Blanks refer to unknown assignments.

Table 6. XP Spectra Fit Assessment for Fe 2p3/2 and Fe 2p1/2 Core Electrons with Peak Assignment for the Fe3+CDTA4- (S2) and Degradation Product Sample (S3) peak position shake-up fwhm (eV) (eV) (eV) S2 Fe 2p3/2 Fe 2p1/2 S3 Fe 2p3/2 Fe 2p1/2 a

709.7a 723.2a 709.9a 710.9b 723.4a 724.4b

714.1 727.6 714.3 718.8 727.8 732.3

2.49 2.49 1.98 3.74 1.98 3.74

fit assessment ∆BE ) 705-735 eV µ ) 21.9%, σµ ) 22.9% ∆BE ) 705-747 eV µ ) 10.5%, σµ ) 25.2%

Fe3+CDTA4-. b Fe(OH)3.

chiometric proportion. For the N 1s spectra, one major peak is acknowledged representing the coordinated tertiary amine with iron, C-(C-)N+(-Fe)-C. As for oxygen, two elemental structures of equivalent proportion, namely the carbonyl (532.0 eV) and coordinated hydroxyl (531.7 eV) oxygen were found. Good fit over Fe3+CDTA4- C 1s, N 1s, and O 1s spectral curves (µ ) 6-11%) could be achieved only if some unidentified spectral components were added to the models. They account for 19.5% of the total concentration for C 1s, 12.3% for N 1s, and 5.4% for O 1s. These compounds may be associated with the presence of carbonates (i.e., Na2CO3) in the Fe3+CDTA4- stock solution. However, it is most likely that limited degradation of Fe3+CDTA4with subsequent formation of organic salts (i.e., NaCH2COOH) explains the need for new foreign peaks in the models. 3.4 XPS Characterization of Iron. Simultaneous fits of Fe 2p3/2 and Fe 2p1/2 electrons in the highresolution XP spectra (705-750 eV) were achieved. The qualities of fit for Fe3+CDTA4- (S2) and degradation product (S3) samples are given in Table 6. Spectral correction for charging effects was calculated at 1.6 eV for sample S3. Normally, Fe3+CDTA4- Fe 2p spectra should represent one chemical arrangement, the hexacoordinated ferric ion within the ligand. Owing to iron paramagnetic behavior, a shake-up photopeak was also added to the fitting. This shape-up photopeak normally shows up within 10 eV of the characteristic peak.22-24 As shown in Figure 11, the shake-up peaks emerge +4.4 eV upfield for both the Fe 2p3/2 and Fe 2p1/2 characteristic peaks, which are themselves separated by 13.5 eV. It was estimated in Section 3.1 that 92% of the total iron in the degradation product sample stands for Fe(OH)3 while the remaining 8% is probably due to the presence of unwashed Fe3+CDTA4-. As a result, two sets

Figure 11. High-resolution XP Fe 2p spectra for the Fe3+CDTA4sample (S2) with the fitted elemental spectra. (a) Fe3+CDTA4- Fe 2p3/2 peak, (b) Fe 2p3/2 shake-up(c) Fe 2p1/2 peak(d) Fe 2p1/2 shakeup.

of peak combination (i.e., Fe 2p3/2 peak + shake-up and Fe 2p1/2 peak + shake-up) were processed with CasaXPS while applying the following constraints: (a) the distance between corresponding Fe 2p3/2 and Fe 2p1/2 peaks was set at 13.5 eV, (b) the distance between Fe3+CDTA4characteristic and shake-up peaks was set at +4.4 eV, (c) the fwhm for corresponding Fe 2p3/2 and Fe 2p1/2 peaks were set equal to Avert spectral overfitting, and (d) the area ratio between Fe3+CDTA4- characteristic and shake-up peaks was set at 0.62 as determined in the previous fitting. Take note that cesium (Cs 3d3/2, Cs 3d5/2), probably introduced via NaCl impurities, was detected at low concentration (Figure 12). ThedegradationproductspectralfitgivestheFe3+CDTA4Fe 2p3/2 peak position at 709.9 eV, which is comparable to what was obtained for sample S2 (709.7 eV). The Fe(OH)3 Fe 2p3/2 peak position was determined at 710.9 eV, well within the acceptable range for ferric (Fe3+) complexes (710-712 eV).22-27 In addition, Fe(OH)3 represents 80.3% of total iron concentration in the degradation product, which corresponds roughly to the abundance inferred from the survey scans. The presence of Fe(OH)3 can also be monitored via the O 1s highresolution spectra, clearly revealing distinctive characteristics when compared to the Fe3+CDTA4- O 1s scan (Figure 10). A broad peak at ca. BE ) 530.5 eV, which is 1.2 eV downfield to the lowest assigned peak for Fe3+CDTA4-, establishes the presence, in substantial proportion, of a solid complex involving oxygen functions not present in CDTA. This peak likely represents Fe(OH)3 even if hydroxide compounds are usually assigned to BE shifts within 531-531.5 eV.27-29 The multimodal broad O 1s spectrum for the degradation product sample is therefore ascribed to BE shifts relative to Fe3+CDTA4-, Fe(OH)3, and perhaps other byproducts such as carbonates. 4. Conclusion The degradation of two common Fe3+-bearing aminopolycarboxylic acids, EDTA and CDTA, in alkaline solutions was studied using full three-level fixed-effects factorial designs. The factor-level combinations swept

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gratefully acknowledged. We would also like to thank Maude Gagnon for her assistance in the laboratory experiments as well as Dr. Alain Adnot for his assistance in the XPS study. Supporting Information Available: Raw experimental measurements in terms of daily absolute degradation rate (∆CFe, µM/day) and relative degradation rate (∆CFe/CFe, day-1) are given in Tables S1-S6. This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited

Figure 12. High-resolution XP Fe 2p spectra for the degradation product sample (S3) with the fitted elemental spectra. (a) Peak combination representing Fe3+CDTA4-. (b) Peak combination representing Fe(OH)3.

the following ranges: pH (8 < pH < 10), temperature (25 < T < 55 °C), ionic strength (0.025 < I < 0.5 M), and ferric chelate concentration (175 < CFe < 450 µM). Analysis-of-variance over replicated (2 × 34) designs indicated that pH and temperature are the main factors increasing Fe3+EDTA4- and Fe3+CDTA4- degradation rates, followed, to a lower extent, by ionic strength and ferric chelate concentration. It was determined that pH × T interaction is the most significant combination among all others. CDTA was determined to be a superior chelating agent for operation in alkaline solutions. It was determined through X-ray photoelectron spectroscopy that inorganic Fe(OH)3 was the primary component of the degradation product. In the Fe3+CDTA4degradation product sample, 80-90% of the total iron (Fe3+) was evaluated as Fe(OH)3. The remaining fraction was determined possibly as adsorbed Fe3+CDTA4onto the solid. It was not possible to identify which one of the organometallic complex degradation or the ferric dechelation was responsible for iron(III) hydroxide formation, since both routes can contribute to its formation. While Fe(OH)3 would probably compose the degradation product through dechelation, it is unclear whether Fe3+ is released after organic degradation. Further analysis of the solution chemistry would be necessary. Acknowledgment Financial support from the Natural Sciences and Engineering Research Council of Canada (NSERC) is

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Received for review May 3, 2004 Revised manuscript received July 1, 2004 Accepted July 2, 2004 IE049636F