[advances in chemistry] aquatic humic substances volume 219 (influence on fate and treatment of...

25 Reactions Between Fulvic Acid and Aluminum Effects on the Coagulation Process

Brian A. Dempsey

Department of Civil Engineering, Pennsylvania State University, University Park, PA 16802

The effects of fulvic acids on the speciation of aluminum are measured by timed colorimetric analyses. Precipitates of Al(OH)3(s) (pKsp = 32.8 for pH 4.5-6.5) form in every region where alum has been shown to be successful for the removal of fulvic acids. Stability functions (average log Κ = 3.39) are reported for the formation of soluble aluminum-fulvic acid complexes. Adsorption functions (fulvic acid on freshly precipitated Al(OH)3(s)) are more than 10 times larger than the stability functions for complexation of fulvic acids with dissolved aluminum.

F'ULVIC ACID (FA) is USUALLY REMOVED from raw water by the coagulation process followed by sedimentation and rapid sand filtration. Filter alum (Al 2(SO 4) 3-14.3H 2 0) is the coagulant most commonly used in the United States. However, ferric chloride, organic polyelectrolytes, and other salts of aluminum(III) or iron are also used. I have previously reported on the removal of F A with salts of aluminum (1-3).

The objectives of the work presented here are to determine the effects of F A on the speciation of aluminum (especially when filter alum is used as the coagulant) and to use this information to predict the mechanism of F A removal during water treatment. The experimental time frame in this study (minutes to days) corresponds to the hydraulic residence time of conventional water-treatment and distribution systems. The relatively short time for re-

0065-2393/89/0219-0409$06.00/0 © 1989 American Chemical Society

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In Aquatic Humic Substances; Suffet, I., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1988.

410 AQUATIC HUMIC SUBSTANCES

action and the presence of fulvic acid means that analytically identifiable crystalline aluminum oxides or hydroxides are not expected to be important in these systems.

Removal of Contaminants Using Aluminum Salts Amirtharajah and Mills (4) studied the removal of suspended solids from water with alum as the coagulant. They identified p H values and alum doses that resulted in the formation of a voluminous precipitate of Al(OH) 3(s). Combinations of p H and coagulant dose that result in a heavy floe that settles are usually identified as the "sweep-floc" zone. The use of these conditions for the coagulation process results in good removal of clays, F A , and other contaminants.

Amirtharajah and Mills (4) also identified p H and coagulant dose values at which removal of the contaminants occurs by the charge-neutralization mechanism. In this case, the negative charge of the contaminants is just neutralized by the cationic coagulant species. An equivalent coagulant-to-contaminant dosage must be used to produce charge neutralization. Overdosing results in restabilization. Lower doses of chemical are required for charge neutralization than for the sweep-floc zone. Although using lower doses has potential advantages, contaminant removal rates are often lower and operational control may be more difficult than for the sweep-floc zone.

Amirtharajah (5, 6) explained his experimental results for coagulation with filter alum by assuming that Al(OH) 3(s) forms whenever coagulation is successful. However, he noted that the results are also consistent with de-stabilization by soluble polymeric species of aluminum. Edwards and Amirtharajah (7) reported that the removal zones for humic acid involve p H and alum doses similar to those required for the removal of clays, except that the stability zone shifts slightly toward lower p H values (7).

Sricharoenchaikit (8) also assumed that the precipitation of Al(OH) 3(s) precedes the destabilization of contaminant species. Packham and associates (9-11), on the other hand, suggested that F A and the dissolved aluminum from alum can react directly to form a precipitate without the preliminary formation of Al(OH) 3(s). They stated that F A is precipitated by soluble hydrolyzed species of aluminum when salts of aluminum are used as coagulants (9), whereas they suggested that removal of clays is dependent on the preliminary precipitation of Al(OH) 3(s) (JO, II). Rebhun and Narlds (12) suggested that removal of humic materials by alum near p H 6.7 is due to direct precipitation by the polymeric species A l 8 ( O H ) 2 0

4 + . Matijevic and co-workers (13, 14), Hayden and Rubin (15), and Stumm and associates (16,17) also suggested that polymers of hydrolyzed aluminum are active reagents when filter alum is used as a coagulant.

Dempsey and co-workers (1-3) used various coagulants that contain aluminum to investigate the conditions required for the coagulative removal

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25. DEMPSEY Reactions Between Fulvic Acid and Aluminum 411

of fulvic acids. Some of their results are shown in Figure 1A. The black area designated I in Figure 1A corresponds to conditions in which a sweep-floc can be generated when aluminum is added in the form of filter alum. Removal of fulvic acid apparently occurs by charge neutralization with conditions designated by black area II. Under conditions within the gray dotted region, fulvic acids can be removed by membrane filtration, but not by sedimentation. Dempsey and co-workers (2) suggested that removal using alum at p H above 6 is due to adsorption of F A on Al(OH) 3(s), but that some removal by alum at lower p H values is due to the direct precipitation of F A by polymers or even monomers of aluminum hydroxide.

Polyaluminum chloride (PAC) is a commercially available coagulant with the formula Al(OH) I (Cl) y (S0 4 ) 2 . Typically, χ is 1.2-2.0, and ζ is usually 0.16 or less. Both direct and indirect evidence indicates that PAC contains ther-modynamically stable polymers of aluminum, especially Al0 4 (Al(OH)2)i 2

Figure 1. Aluminum doses and pH values that permit removal of fulvic acid from water when (A) alum or (B) polyaluminum chloride is the coagulant.

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(subsequently abbreviated A l 1 3 ) ; this radical has a +7 charge. PAC achieves good removal of contaminants even at p H less than 4 and A l , less than 10" 4 5 M), conditions under which neither Al(OH) 3(s) nor A l 1 3 are thermo-dynamically stable (part of the black areas in Figure IB). This coagulative behavior, analogous to that of synthetic organic polyelectrolytes (12, 18), indicates that the polymeric species in PAC are relatively inert.

Some of the investigators cited have measured both the removal of contaminants and the speciation of hydrolyzed aluminum. Evidence for the presence or activity of a certain species of hydrolyzed aluminum is usually based on indirect experimental evidence and logical argument (2, 9, 13-15, 19, 20). Hundt (21) characterized the species that form after the addition of alum, aluminum chloride, or PAC to water in the absence of F A or clays. He showed that when alum was added and the p H was above 4.5, most of the aluminum was retained by a membrane filter. In contrast, filterable species persisted until higher p H values for PAC. Tambo (22), using alum, isolated labile (less than 10-min longevity) species of aluminum with highly positive electrophoretic mobilities. These species have not been determined to be either solid or polymeric.

The speciation of the aluminum coagulants, after addition to the raw waters, is uncertain. Critical compilations of thermodynamic data (23, 24) and evaluations of the literature regarding the hydrolysis of aluminum (24, 25) are available. However, reports disagree regarding the best free energy values, and enthalpy values have not been obtained for many of the hydrolysis reactions of aluminum. Additionally, the short time frame of the coagulation process and the presence of contaminants during coagulation make the prediction of speciation difficult.

Experimental Methods The collection, extraction, cleanup, preservation, and characterization of FAs have been previously described (26). Two FAs, designated FA1 and FA4, were used in these experiments. Both of these materials are derived from Lake Drummond, VA. Solid-sample l 3 C NMR studies indicate that approximately 80% of the organic carbon is aliphatic. FA1 has 11.4 meq and FA4 has 11.6 meq of carboxyl functional groups per g of organic carbon; more than a third of these acidic groups are ionized at pH 3. Aldrich humic acid (HA) was also used in a few experiments. Water was distilled and then treated by a microfilter (Milli-Q) system. All other chemicals were reagent grade or better.

Experiments were run in 40-mL glass sample bottles with poly(tetra-fluoroethylene) (Teflon)-lined caps at room temperature (23-25 °C). FA solution and stock acetate were added to water to give between 0 and 84 mg/L of FA or HA organic carbon and total acetate of 8 X 10̂ M. For comparison, 2.1 to 84 mg/L of FA4 contains 2.4 Χ 10-5 to 9.7 Χ ΙΟ -4 M of carboxylic functional groups. The pH was adjusted to the desired value by using HCI or NaOH. Then the appropriate amount of alum was injected and the solution was mixed. The stock alum solution contained 0.0018 M total Al(III). The final concentration of aluminum varied, but it was 1.05 mg/L (3.9 X 10 5 M) in most cases. Samples were taken at 5 min, 1 h,

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4 h, 1 day, and sometimes 7 days for the analysis of aluminum. The pH was measured for each sampling period. Except for the 5-min samples and when filtration was used as a separation process, the sample bottles were placed in an ultrasonic bath for 10 min prior to sampling in order to break up any particulates that may have formed.

The concentration and speciation of aluminum was determined by the ferron (8-hydroxy-7-iodoquinolinesulfonic acid) method. A 2-mL aliquot of the analyte was placed in a cuvette (path length = 1 cm) and 0.8 mL of the ferron reagent as described by Bersillon (20) was added. The cap was inserted, and macromixing was completed within 5 s of the ferron addition. The pH after addition of ferron must be very consistent (27); addition of acid for digestion of aluminum species should be avoided unless the sample is back-titrated or the molar absorptivity is shown to be unchanged. Absorptivity was 7.78 mM"1 cm 1, with a standard deviation of 0.03 at 370 nm. Monomeric aluminum reacts very rapidly, with a reported pseudo-first-order k = 2.3 min 1 (28).

This value was confirmed in our experiments for times greater than 15 s, but the reaction was even faster before 15 s, so that 85% of the total aluminum in dilute alum solutions had reacted within 30 s. Reaction rates for most other species of aluminum are considerably slower (19, 27, 28) and the blank-corrected absorbance at 30 s is claimed to represent inorganic monomeric aluminum. Although the rate of nonmonomeric color development increased with increasing concentrations of either FA or total aluminum, the extrapolated contribution of these nonmonomeric species to the 30-s reading was typically less than 10% of the 30-s reading.

This contribution was determined on the basis of the slopes (absorbance versus time) at 2 min, when the monomeric aluminum was 99% reacted. Reagent blanks were analyzed by measuring the absorbance of samples (minus aluminum) against distilled-deionized (DDI) water. Ferron reagent was added to both cuvettes. Sample blanks were analyzed by comparing the absorbance of the whole samples (including aluminum and FA) against D D I water; ferron reagent was not added.

Solutions of polyaluminum chloride that contain the polymer Al04(Al(OH)2)i27+

(abbreviated Ali3) have an initial rate of color development of 0.071 min 1 (our data) to 0.075 min 1 (28); thus, only 3.5% of such aluminum is reacted with ferron at 30 s. The rate of color development from the PAC that was used in these experiments slows substantially, however, so that only 57% of the Al i 3 is reacted after 22 h. On the other hand, the aluminum in suspensions that are predominantly Al(OH)3(s) is sometimes totally reacted in less than 1 h, a result indicating greater lability with respect to the ferron reagent than for Ali3 or other polymeric materials. As a result, definitions of the aluminum that reacts with ferron in 2 h as monomeric plus polymeric (Ala plus Al6) and the remaining aluminum (Alc) as Al(OH)3(s) cannot be justified for the situations that we have studied.

Some plots of absorbance versus time are shown in Figure 2 for diluted alum (very rapid reaction), Al i 3 (very slow and monotonie reaction), amorphous Al(OH)3(s) (S-shaped curve), and aluminum that is complexed by FA (very rapid reaction for inorganic monomeric aluminum and very slow, monotonie reaction for organically bound aluminum). The S-shaped curve occurred at some sample time (typically at 5 min, 1 h, and 4 h after coagulation) in every case in which Al(OH)3(s) was visually observed. The inflection point in the S-shaped curve always occurred within 2 h after the addition of ferron reagent, and the incremental absorbance that occurred after the inflection point could often be removed by membrane filtration. In this work the ferron test is used to determine monomeric (30-s) aluminum and as evidence for the presence of Al(OH)3(s). These two uses have been corroborative in every case.

Some data for filterable aluminum are presented in this chapter. Filterable aluminum is defined as the fraction that passes membrane filters with 0.2-μ m pores.

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0.7

0 6 + ^' ' υ* β <^ ° ' u m

£ 0.5 f Ο

CD U C Ο

Ο ω

Χ)

< 0.1 +

0.0 0 10 20 30 40

m m

Figure 2. Absorbance versus time after addition of the ferron reagent for diluted alum (no base added), polyaluminum chloride (C0 = 1.3 M), amorphous

Al(OH)3(s), and aluminum complexée with FA.

Aliquots of 10 mL of filtrate per cm 2 of exposed filter were discarded before keeping filtrate for analysis. The filtrate was analyzed by electrothermal atomic absorption spectroscopy (29) or by ferron colorimetric technique.

The monomeric speciation of aluminum was calculated by using data from the 30-s ferron reactions and equilibrium constants taken from Smith and Martell (23) or Baes and Mesmer (24). Thus the overall stability constants (log β* at ionic strength 7 = 0) were -4.99, -9.3, -15.0, and -23.0. Activity coefficients were calculated with the extended Debye-Huckel equation (30) and / = 0.001. The logarithmic values of the apparent overall constants for the sequential hydrolysis of A l 3 + were therefore -5.07, -9.42, -15.13, and -23.12.

Precipitation of AliOH)3(s) One objective of this work is to determine the effects of F A on the speciation of aluminum during coagulation processes. Experiments have been run to determine whether Al(OH) 3(s) forms during the coagulation process. A test for the presence of this phase can be performed by observing whether consistent values exist for the ion activity product ({Al 3 +}-{OH~} 3); a consistent value occurs when the slope is -3.0 for the plot log [ A l 3 + ] versus p H . Additional tests for the presence of Al(OH) 3(s) are an S-shaped ferron trace and the removal of nonmonomeric aluminum by membrane filtration.

Data are presented first for experiments run for p H values between 3.5 and 7. The calculated concentrations of A l 3 + (based on the 3-s ferron test

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and the apparent hydrolysis constants) are shown in Figure 3. These data are for analyses run at 4 h and 1 day. The total aluminum concentration was either 3.9 X Ι Ο ^ ο ^ . β X 10~ 5 M, and the concentration of humic substances (FAI , FA4, or Aldrich humic acid) was either 0 or 4.2 m g / L of organic carbon (i.e., 4.8 x 10"5 e q / L of total carboxylic acid groups).

The concentration of the A l 3 + species and the shape of the ferron trace are constant until a p H of about 4.75. At all higher p H values and for 0 or 4.2 m g / L of D O C , the concentration of monomeric aluminum is diminished and the ferron trace assumes the S-shape that has been identified with amorphous Al(OH) 3(s). Nonmonomeric aluminum can be removed by membrane filtration. A linear least-squares analysis of the data for p H 4.75-6.5 reveals that

log [ A l 3 + ] = 9.06 - 2.96 p H (r 2 = 0.987) (1)

where [ A l 3 + ] refers to concentration. The theoretical slope is -3.00 if aluminum speciation is controlled by Al(OH) 3(s). The p K s p (corrected to zero ionic strength) was calculated for each datum within the given range of p H values; the average and standard deviation for p K s p was 32.82 ± 0.17. These data are consistent with the formation of microcrystalline gibbsite. Nordstrom and Ball (31) and Driscoll and co-workers (32) discovered similar relationships between [ A l 3 + ] and p H for stream waters affected by acid-mine drainage and by acid precipitation, respectively.

- 2

- 4 -

- 6 -

o ο Ê ο Avg p K s p = 32.8

for pH > 5.0

-10-

-12-

° Alum only A w / A l d HA ο w/FA#1 ο w /FA#4

•14-

pH

Figure 3. log [Al3+] versus pH under several experimental conditions. The [Al3+] is based on the 30-s ferron result (this research) and stability constants

(23, 24).

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Moderate concentrations of humic materials (up to at least 5 m g / L of DOC) enhance the precipitation of Al(OH) 3(s) at the slightly acidic p H values (4.75-5). This effect of humic materials on the precipitation of Al(OH) 3(s) is illustrated in Figure 4, where the inorganic monomeric aluminum (30-s ferron results) or A l f l is plotted as a function of increasing concentration of humic material.

Changes in p H account for a minor part of the decreases in inorganic monomeric aluminum; the range of p H values was 0.01 for the Aldrich H A (pH 5.0) tests and 0.09 for the FA1 (pH 5.0) tests. The sample blanks for these tests were identical to the reagent blanks, a result indicating that the changes in absorbance that are shown in Figure 4 are not due to analytical anomalies, such as unaccounted light scattering. The minimum value for inorganic monomeric aluminum occurs when the F A normality (in total carboxylic groups) is equivalent to the molar concentration of aluminum (1.05 mg/L). The charge density of FA1 is greater than that for Aldrich H A . These conditions have been designated zone II in Figure 1A.

At much higher concentrations of humic materials (e.g., the curve in Figure 4 for FA1 and p H 5.0 for D O C greater than 20 mg/L), the ferron trace assumes completely monotonie response at all reaction times. The concentration of inorganic monomeric aluminum decreases with increasing D O C . No filterable precipitate is formed. These observations are consistent with the formation of a soluble complex between aluminum and the humic materials.

0.8

0.2 +

0 . 0 - 1 — · — ι — · — ι — » — ι — « — ι — . — ι — « — ι — « — ι — « — I 0 10 20 30 40 50 60 70 80

[FA] in m g / L of DOC

Figure 4. Monomeric aluminum as determined by the 30-s ferron test versus the concentration (DOC) offulvic acid. See text for explanation of the inflection

points.

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The concentration of inorganic monomeric aluminum increases between D O C of 4.2 and 20 m g / L of FA1 and p H 5.0 (Figure 4). These conditions are analogous to the gray-hatched area below zone II in Figure 1A (i.e., a region in which the ratio of coagulant to humic material is inadequate for charge neutralization). The increase in inorganic monomeric aluminum may be due to an increasing disorder (and specific surface area) for Al(OH) 3(s) with excessive and increasing concentration of the humic materials. The rate of the increase in absorbance for the nonmonomeric fraction of aluminum is directly proportional to the concentration of humic material. In addition, the ferron trace is S-shaped, and some aluminum can be removed by filtration (but not by sedimentation) in this zone.

The part of this study that deals with the speciation of aluminum at slightly alkaline p H values is an extension of prior work (2), which showed a linear relation at p H 7.2 between the concentration of humic materials and the required dose of alum for the removal of the F A by coagulation and filtration. Thus, removal of 3.5 m g / L of D O C required at least 1.3 mg /L of aluminum, 10 m g / L of D O C required at least 3.6 mg/L of aluminum, and 35 mg /L of D O C required at least 13 mg /L of aluminum before any removal of F A occurred. Some of these results are displayed by the curves in Figure 5. These doses of aluminum are orders of magnitude in excess of the solubility of Al(OH) 3(s) at p H 7.2 and in the absence of humic materials.

These experiments were repeated to determine the speciation of the aluminum. Experiments were performed for 0, 3.5, and 10.0 mg/L of D O C (FA1) and for five concentrations of total aluminum. The D O C that passes through a membrane filter is shown as a function of the coagulant dose in Figure 5. In addition, the ion activity product for Al(OH) 3(s) has been determined (on the basis of the 30-s ferron results and the hydrolysis constants for monomeric aluminum), and these values are shown in Figure 5. These values were affected neither by the total concentration of aluminum nor by the extent of removal of F A .

The consistency of the calculated ion activity products at a p H of 7.2 and the S-shaped ferron traces are strong indications of the presence of Al(OH) 3(s) and of the control of aluminum speciation by this precipitate. This evidence indicates that Al(OH) 3(s) has formed, although neither aluminum nor F A can be removed by membrane filtration. Prior research has demonstrated that F A adsorbs strongly on Al(OH) 3(s) at this p H , with 3 mol of F A carbon (e.g., 0.4 eq of carboxylic functional groups) adsorbed on every mole of Al(OH) 3(s) when the residual F A D O C is only 1 mg/L. The resultant negative charge density of these aggregates causes electrostatic repulsion between particles (33, 34) and, for these experiments, has resulted in stabilization of the submicrometer-sized particles of F A and Al(OH) 3(s).

The zones in which F A can be removed by coagulation with alum and filtration by membrane filters are compared with the conditions required for the precipitation of Al(OH) 3(s) in Figure 6. The zone of precipitation of Al(OH) 3(s) is indicated by the solid line. At slightly acidic values of p H the

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Ο ο Q

CP Ε

ο 3

25

20 +

15 +

10

pK Sp before, at, and after breakpoint

no FA (5 doses of Al) p K c n = 32.41

3.5 mg/L FA p « s p = 32.57, 32.53, 32.51 10 mg/L FA pK„ n = 32.37, 32.32, 32.24

10

Aluminum dosage (mg/L) Figure 5. Removal offulvic acid at pH 7.2 as a function of the initial fulvic acid DOC and the dose of alum. The ion activity products for Al(OH)3(s) are

indicated for various experimental conditions.

-3.5

2 -4 .0

CD Ο

-4.5+

-5.0

7 mg/L FA4

3 . 5 mg/L FA1

PH

Figure 6. Comparison of the boundaries for removal (after coagulation and filtration) of various fulvic acids and the boundaries for the formation of

Al(OH)3(s). Alum is the coagulant in all of these experiments.

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removal of F A occurs at p H values and alum doses that are appropriate for the precipitation of Al(OH) 3(s). At higher p H values the formation of Al(OH) 3(s) is not sufficient for the removal of F A . According to the hypothesis that has been proposed here, the particles are dispersed until higher ratios of positively charged Al(OH) 3(s) to negatively charged F A are obtained. The coherence of removal zones and solubility zones for alum is an indication that Al(OH) 3(s) is present when F A is removed. It is not necessarily a negation of the simultaneous presence of reactive polymers of aluminum.

Complexation Between Al(HI) and FA Stability functions for the complexation of aluminum by F A have been calculated for zones where Al(OH) 3(s) is not present (according to criteria that have been discussed). The data are shown in Table I. A l l stability functions are in the form

= [complex] [free Al] x [free FA] K )

where [complex] is the molarity of complexed aluminum (i.e., the total aluminum minus [free Al]); [free Al] is the molarity of aluminum that reacts with ferron in 30 s ; and [free FA] is the total normality of carboxylic acid groups minus [complex]. Thus an unsupported assumption is made that each complexed aluminum reacts with only one carboxylate group. The stability functions that correspond to data where the ion activity product for Al(OH) 3(s) exceeds 10" 3 2 8 are marked in Table I with an asterisk. The presence of very high concentrations of F A may inhibit the precipitation of Al(OH) 3(s); nevertheless, the marked data should be viewed with skepticism. Some stability functions for the complexation of calcium with F A are also shown in Table I. These functions are based on charge-balance calculations (26). The numerators for the Ca data have been changed from normality (as reported in ref. 26) to molarity to obtain greater compatibility with other data in Table I.

The values for complexation of aluminum with F A that are shown in Table I are close to those that have been reported by Backes and Tipping (35) (log Κ = 2.90), who used a dialysis technique and concentrations for total aluminum and humic material similar to those reported here. Their p H values ranged from 3 to 4.8. Backes used a high-molecular-weight fraction of humic acid to minimize losses through the dialysis membrane.

Schnitzer and Hansen (36) used two techniques (method of continuous variations and the ion-exchange equilibrium method) to determine stability functions between aluminum and soil F A at p H 2. Both techniques gave log Κ values of 3.7. The Κ value was expressed in terms of molarity of the F A , and values would be somewhat lower if expressed in normality of carboxylic

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Table I. Complexation Functions for FA and Stability Constants for Analogues of FA with Aluminum or Calcium

Log of Function Ligand Metal or Constant Conditions FA1 Al 3.59 pH 4.30, DOC 17.1 mg/L

Al 3.62e pH 5.07, DOC 17.1 mg/L Al 3.32° pH 9.07, DOC 17.1 mg/L Al 3.49 pH 4.15, DOC 20.9 mg/L Al 3.52e pH 5.07, DOC 20.9 mg/L Al 3.27* pH 8.90, DOC 20.9 mg/L AJ 3.23 pH 4.12, DOC 40.9 mg/L Al 3.32fl pH 4.98, DOC 40.9 mg/L Al 3.15e pH 8.72, DOC 40.9 mg/L Al 3.36 pH 4.03, DOC 78.2 mg/L Al 3.27 pH 4.77, DOC 78.2 mg/L Al 3.58û pH 8.08, DOC 78.2 mg/L

F A P Ca 1.72 pH 5, pCa 3, TOC 666 mg/L Ca 2.09 pH 7, pCa 3, DOC 666 mg/L Ca 2.25 pH 9, pCa 3, DOC 666 mg/L Ca 1.64 pH 5, pCa 2.4, DOC 666 mg/L Ca 1.91 pH 7, pCa 2.4, DOC 666 mg/L Ca 2.10 pH 9, pCa 2.4, DOC 666 mg/L

Acetate Al 1.51 /-> 0 Ca 1.18

Oxalate Al 6.1 I = 1.0 Ca 1.66 / = 1.0

Tartrate Al 5.32 ί = 1.0 Ca 1.80 / = 0.1

Salicylate Al 12.9 I = 0.1 Ca 0.15 ί = 0.16

Phthalate Al 3.18 / = 0.5 Ca 1.07 I = 0.1

"Ion activity products ({Al3+}{OH}3) greater than or equal to 10-328, but ferron traces without indication of Al(OH)3(s). bSee ref. 28 for more data regarding complexation of Ca with FA.

functional groups on the F A . Schnitzer used higher concentrations of F A and aluminum than for data reported in Table I, but used FA-to-aluminum ratios similar to those used here. Pott and co-workers (37) used a cation-exchange technique to measure complexation functions between aluminum and humic materials. They discovered functions 3 orders of magnitude stronger than those reported in Table I.

Because humic materials are mixtures, the metal-to-ligand ratios are important. Pott and co-workers (37) used lower aluminum concentrations than were used for the experiments summarized in Table I, but the ratios of humic material to aluminum fall within the same range. Young and Bache (38) also modeled the complexation reaction, but only used a limited set of data for validation of the model. Both Backes and Pott found that complexation of aluminum with humic materials was insignificant compared to the

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hydrolysis of aluminum at p H 6 or higher. This behavior, which agrees with the results reported here, has substantial impact on the mechanism of removal of F A during coagulation processes.

Some stability constants for the complexation of aluminum with specific organic ligands are also shown in Table I. Monocarboxylic organic acids (e.g., acetic acid), strong chelating agents (e.g., salicylic acid), and weaker chelating agents (e.g., phthalic acid) are often proposed as analogues for the reactive complexing groups on molecules of FA. When we compare F A and specific organic ligands as complexing agents, we must emphasize that F A is a mixture of materials and that the stability function cannot have the same meaning as the stability constants for complexation of aluminum by well-defined l i gands (39, 40).

The following observations are made with respect to the stability functions and constants listed in Table I. First, the formation functions for F A and the selectivity of the ligand for A l versus Ca are both greater for F A than for acetic acid. Second, the stability functions and the selectivity for A l versus Ca are less for F A than for salicylate (and other chelators that form five- or six-membered chelate structures). Finally, the selectivity for A l versus Ca and the stability constant for complexation of aluminum with phthalate are quite close to the values for F A .

The best analogue for these experimental conditions is phthalate, which forms a seven-membered chelate ring with metals. Fettes (41) showed that chelate rings of greater than 30 members are sometimes stable for complexation of metals with macromolecules. Dempsey and O'Melia (26) suggested that such chelation may be important when metals are complexed with F A . The comparisons described support this hypothesis.

Comparison of Complexation and Adsorption Dempsey, Ganho, and O'Melia (2) reported isotherms for the adsorption of F A on Al(OH) 3(s) that is formed during the coagulation process. The isotherms were performed at p H 7 and for total concentrations of F A and aluminum similar to those used in the experiments reported here. For a residual soluble F A concentration of 1 mg/L, over 3 mol of F A carbon was adsorbed per mole of aluminum. If there were 11.24 meq of carboxylic functional groups per g of F A D O C , then the adsorption function would be 3.9 x 10 4 (log Κ = 4.6).

The terms and derivation of the adsorption function are identical to the complexation functions described here and shown in Table I. Both are expressed in terms of complexed over free F A functional groups divided by the free concentration of A l ions (in monomeric form for complexation and in solid form for adsorption). In both cases a 1:1 stoichiometry is assumed between carboxylic functional groups on the F A and aluminum atoms (a simplification that does not take possible chelation into account).

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The very interesting result is that the adsorption reaction is 16 times stronger than the complexation reaction. This result seems reasonable. The weight-averaged molecular weight of F A is at least 600, and there are over 11 meq of carboxylic functional groups per g of F A carbon, yielding more than six carboxylic groups on the average molecule of F A . Many investigators (e.g., 42) have shown that the strength of adsorption is a function of the molecular weight and the number of adsorbing groups on the adsorbate.

This quantitative comparison corroborates the evidence reported in Figure 5, which indicates that Al(OH) 3(s) has formed at p H 7.2, even when the precipitate cannot be removed by membrane filters with 0.2-μπι pores. There should be similar strength of adsorption at p H 5; the F A is less negatively charged, but the Al(OH) 3(s) is more positively charged. This qualitative argument strengthens allegations that F A is removed by the adsorption mechanism rather than by direct precipitation at p H values close to 5.

Conclusions A n important conclusion from this work is that Al(OH) 3(s) forms in the presence of F A over the entire range of p H and total aluminum conditions that lead to removal of F A during coagulation using alum. The solubility of Al(OH) 3(s) for times ranging from minutes to days and for p H values less than 6.5 is defined by p K s p = 32.8. Solubility is greater than this when the p H exceeds 6.5.

This conclusion does not necessarily indicate that polymeric species of aluminum are unimportant in coagulation with alum, even after several minutes have elapsed, because a mass balance on the aluminum is not reported here. Most likely, polymeric species form when alum is added to water, because polymers may be intermediates between the initial coagulant (monomeric aluminum) and one of the final products (Al(OH)3(s)) (22, 43). It is not certain whether such intermediate species are necessary for effective coagulation. Effective coagulation can be accomplished by using relatively inert polymeric species of aluminum (e.g., polyaluminum chloride).

F A affects the nature of the Al(OH) 3(s). The effects seem to be caused by the relatively strong adsorption of F A and by the moderately high negative charge densities for the F A . As a result, F A enhances the formation of Al(OH) 3(s) at p H ^ 5 . This behavior is similar to that observed for sulfate. At p H 7.2, the F A causes the stabilization of submicrometer-sized particles of Al(OH) 3(s). The high and variable charge density and strong adsorption of F A affect the shape of the typical removal diagram for F A . This effect causes removals at lower p H than for clay for a given coagulant dose and requires a greater coagulant dose at higher p H values than would be needed for clay.

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The complexation of soluble aluminum by these FAs is moderate in strength—less than for model chelators with five- or six-membered chelate rings, much greater than for monoprotic organic acids, similar in strength (and selectivity for aluminum versus calcium) to phthalic acid. This provocative similarity is a clue (but not a confirmation) to the structure of the reactive functional groups on F A . F A is a mixture, and the reported stability functions represent the dominant complexing groups at the degree of experimentally observed complexation.

Adsorption of F A on Al(OH) 3(s) is a much stronger reaction than the complexation of monomeric aluminum with F A . This fact is very useful for a mechanistic understanding of the removal of F A during coagulation with salts of aluminum. The conclusion is reasonable, but not necessary, that the most important mechanism for removal of F A with alum is by adsorption of F A on Al(OH) 3(s). This prediction is made for the entire p H range for which removal occurs and for the low-to-moderate concentrations of F A that have been tested.

References 1. O'Melia, C. R.; Dempsey, B. A. Proc. Annu. Public Water Supply Eng. Conf.;

University of Illinois at Urbana-Champaign, 1982; pp 5-14. 2. Dempsey, Β. Α.; Ganho, R. M.; O'Melia, C. R. J. Am. Water Works Assoc. 1984,

76, 141-150. 3. Dempsey, Β. Α.; Sheu, H.; Ahmed, T. M. T.; Mentink, J. J. Am. Water Works

Assoc. 1985, 77, 74-80. 4. Amirtharajah, Α.; Mills, Κ. M. J. Am. Water Works Assoc. 1982, 74, 210-216. 5. Amirtharajah, A. Coagulation and Filtration: Back to the Basics; from Am. Water

Works Assoc. Semin.; 1981; pp 1-22. 6. Amirtharajah, A. J. Environ. Eng. (N.Y.) 1986, 112, 1085-1108. 7. Edwards, G. Α.; Amirtharajah, A. J. Am. Water Works Assoc. 1985, 77, 50-57. 8. Sricharoenchaikit, P. Ph.D. Thesis, Syracuse University, 1984. 9. Hall, E. S.; Packham, R. F. J. Am. Water Works Assoc. 1965, 57, 1149-1166.

10. Packham, R. F. J. Appl. Chem. 1962, 12, 564-568. 11. Packham, R. F. Proc. Soc. Water Treat. Exam. 1963, 12, 15. 12. Narkis, N . ; Rebhun, M. J. Am. Water Works Assoc. 1977, 69, 325-328. 13. Matijevic, E.; Mathai, K. G.; Ottewill, R. H.; Kerker, M. J. Phys. Chem. 1961,

65, 826-830. 14. Matijevic, E.; Janauer, G. E.; Kerker, M. J. Colloid Sci. 1964, 19, 333-346. 15. Hayden, P. L.; Rubin, A. J. In Aqueous-Environmental Chemistry of Metals;

Rubin, A. J., Ed.; Ann Arbor Science: Ann Arbor, 1974; p 317. 16. Hahn, H. H.; Stumm, W. J. Colloid Interface Sci. 1968, 28, 134-144. 17. Stumm, W.; Morgan, J. J. J. Am. Water Works Assoc. 1962, 54, 971-992. 18. Edzwald, J. K.; Becker, W. C.; Tambini, S. J. J. Environ. Eng. (N.Y.) 1987,

113, 167-185. 19. Smith, R. W. In Non-Equilibrium Systems in Natural Water Chemistry; Gould,

R. F., Ed.; American Chemical Society: Washington, DC, 1971; p 250. 20. Bersillon, J. L. ; Hsu, P . H . ; Fiessinger, F. Soil Sci. Soc. Am. J. 1980, 44, 630-634. 21. Hundt, T. R. Ph.D. Thesis, The Johns Hopkins University, 1985. 22. Tambo, N.; Kamei, T. Chapter 28 in this volume.

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23. Smith, R. M . ; Marteil, A. E. Critical Stability Constants; Plenum: New York, 1974.

24. Baes, C. F., Jr.; Mesmer, R. E. The Hydrolysis of Cations; Wiley: New York, 1976; p 112.

25. Davis, J. Α.; Hem, J. D. In The Environmental Chemistry of Aluminum; Sposito, G., Ed.; CRC Press: Boca Raton, FL, 1988.

26. Dempsey, Β. Α.; O'Melia, C. R. In Aquatic and Terrestrial Humic Materials; Christman, R. F.; Gjessing, E. T., Eds.; Ann Arbor Science: Ann Arbor, 1983; p 239.

27. Schonherr, S.; Gorz, H . ; Gessner, W. Z. Chem. 1980, 20, 422. 28. Gessner, W.; Winzer, M . Z. Anorg. Allg. Chem. 1979, 452, 151-156. 29. Pagoni, P. M.S. Thesis, University of Missouri-Rolla, 1986. 30. Stumm, W.; Morgan, J. J. Aquatic Chemistry, 2nd ed.; Wiley-Interscience:

New York, 1981. 31. Nordstrom, D. K.; Ball, J. W. Science (Washington, D.C.) 1986, 232, 54-56. 32. Driscoll, C. T.; Baker, J. P.; Bisogni, J. J.; Schofield, C. L. In Geological Aspects

of Acid Precipitation; Bricker, O. P., Ed.; Butterworths: London, 1984; p 55. 33. Gibbs, R. J. Environ. Sci. Technol. 1983, 17, 237-240. 34. Neihof, R. Α.; Loeb, G. I. Limnol. Oceanogr. 1972, 17, 7-16. 35. Backes, C. Α.; Tipping, E. Water Res. 1987, 21, 211-216. 36. Schnitzer, M . ; Hansen, Ε. H . Soil Sci. 1970, 109, 333-340. 37. Pott, D. B.; Alberts, J. J.; Elzerman, A. W. Chem. Geol. 1985, 48, 293-304. 38. Young, S. D.; Bache, B. W. J. Soil Sci. 1985, 36, 261-269. 39. MacCarthy, P.; Smith, G. C. In Chemical Modeling in Aqueous Systems; Jenne,

Ε. Α., Ed.; American Chemical Society: Washington, DC, 1979; p 201-222. 40. Perdue, E. M.; Lytle, C. R. In Aquatic and Terrestrial Humic Materials; Christ

man, R. F.; Gjessing, Ε. T., Eds.; Ann Arbor Science: Ann Arbor, 1983; pp 295-313.

41. Fettes, Ε. M . Chemical Reactions of Polymers; Interscience: New York, 1964; p 14.

42. Howard, D. B.; Woods, S. J. In Adsorption from Solution at the Solid-Liquid Interface; Parfitt, G. D.; Rochester, C. H . , Eds.; Academic: London, 1983.

43. Teagarden, D. L . ; White, J. L . ; Hem, S. L. J. Pharm. Sci. 1981, 70, 808-810.

RECEIVED for review July 24, 1987. ACCEPTED for publication March 2, 1988.

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[advances in chemistry] aquatic humic substances volume 219 (influence on fate and treatment of...

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