Chemistry of riverine and estuarine suspended particles from the Ouse-Trent system, UK

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Colloids and Surfaces A: Physicochemical and Engineering Aspects 120119971183 198

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Chemistry of riverine and estuarine suspended particles from the Ouse-Trent system, UK J.R. Ferreira ~, A.J. Lawlor, J.M. Bates, K.J. Clarke, E. Tipping * Institute o[ Freshwater Ecology, Ambleside. Cumhria. 1,A22 OLP. UK Received 2 October 1995: accepted 24 April 1996

Abstract Three samples of suspended particulate matter (SPMt, two riverine and one estuarine, were characterised by bulk chemical analysis, electron microscopy, chemical extraction, electrophoresis and acid base titration. The river samples contained both mineral and biological material and were rich in organic matter (R. Ouse sample 24%, R. Trent sample 54%1, while the estuarine sample consisted mostly of clay minerals, with only 9% of organic matter. The river particles had higher contents of base-extractable carbon and phosphate and acid-extractable trace metals than the estuarine sample. The electrophoretic mobilities of all three samples were negative, but acid base titrations revealed differences in their proton-dissociating characteristics, the R. Ouse sample having ~ 2 meq g ~ of pH-dependent charge titrating between pH 4 and 9, approximately twice the content of the R. Trent sample, and four times that of the estuarme sample. However, the estuarine sample had a relatively large fixed (negative) charge ( ~ - 0 . 7 meq g t l. Centrifugation depletion experiments showed that for all three samples, phosphate adsorbed most strongly at pH 6 7, the particles releasing phosphate to solution at both low and high pH values. Trace metal adsorption experiments, carried ou! by centrifugation-depletion and, in the case of Cu, with an ion-specitic electrode, showed that adsorption increases v~ith pH, except in some cases at pH > 7, where solution complexation reverses the trend. A simple metal proton exchange model provided an approximate description of the observations, and allowed comparison amongst metals and particle samples. The three samples adsorbed trace amounts of metals with the same selectivity: Cs < Sr < C o - N i - C d < Zn < Cu < Pb < Eu. The strengths of adsorption of a given metal by the three samples varied over an order of magnitude, increasing in the order Ouse estuary < R. Trent ~
1. Introduction S u s p e n d e d p a r t i c u l a t e m a t t e r ( S P M ) in rivers a n d estuaries is c o n s i d e r e d to exert a m a j o r influence on the b e h a v i o u r a n d b i o a v a i l a b i l i t y of solutes, especially metals a n d p h o s p h a t e [ 1 - 4 ] . * Corresponding author. 1 O11 leave from: lnstituto de Pesca and CENA/USP, Sao Paulo, Brazil.

The interactions d e p e n d on the types and a m o u n t s of particle present, pH, c o n c e n t r a t i o n s of the solutes in question, chemical i n t e r a c t i o n s in the solution phase, and ionic strength. Research perf o r m e d d u r i n g the last two decades has p r o v i d e d c o n s i d e r a b l e insight into the processes a n d mechanisms of e n v i r o n m e n t a l surface chemistry (see e.g. Ref. [ 5 ] ) , allowing the qualitative or semiq u a n t i t a t i v e i n t e r p r e t a t i o n of field observations. P e r h a p s the biggest obstacle to m o r e q u a n t i t a t i v e

0927-7757:97/$17.00 Copyright ~" 1997 Elsevier Science B.V. All rights reserved PII S0927-7757( 96)03721- 1

184

J.R. Ferreira et al./Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 183-198

explanation and prediction is uncertainty regarding the definition of the types and amounts of surfaces involved in real environmental systems [6]. M011er and Sigg [7] reported order-ofmagnitude variations in solid solution partition coefficients for river water SPM. Sung [8] analyzed published data for solid-solution partitioning and showed that particulates from different rivers and estuaries displayed substantial variations in the strengths of their interactions with Cu, Cd and Zn. These findings reinforce the view [9] that a "universal" description of adsorption by SPM requires detailed consideration of the contributions of the component adsorptive phases. The present study was undertaken in order to explore variations in particle properties and particle-metal and particle-phosphate interactions within a river-estuary system in the UK. Samples of natural SPM were taken from three sites in the Ouse Trent system of northern England, which is a focus of the Land-Ocean Interaction Study (LOIS) currently being carried out under the auspices of the U K Natural Environment Research Council [10]. The SPM samples were subjected to a series of laboratory characterisation studies, and also to sorption experiments with metals and phosphate.

2. Study sites and sampling Information on chemistry, flows and suspended solids concentrations of the two rivers is given in Table 1. In the low-salinity region of the estuary, the concentration of orthophosphate was 5-10 ~tM during the years 1977-1981 [11], while concentrations of dissolved organic carbon (DOC) are in the range 4-20 mg 1 1 (M.R. Williams and G.E. Millward, personal communication, 1995). Suspended sediments in the Humber estuary originate mainly from the eroding Holderness coast, to the north of the mouth of the estuary, and from the North Sea [12,13]. In this work, sahaples were taken from Naburn Lock ( N G R SE 591455) and Cromwell Lock ( N G R SK 810613), which are just upstream of the tidal limits of the Rivers Ouse and Trent respectively. The sampling dates were 26 April 1994

Table 1 General information on the Rivers Ouse and Trent. The values for chemical variables are means for the first 8 months of 1994. Ranges of flow and suspended solids concentrations for the same period are also presented. These data were obtained as part of the LOIS project (G.J. Leeks, P.A. Cranwell, D.V. Leach, J.P. Lishman, A.F.H. Marker, C. Neal, A.C. Pinder, E. Rigg, G. Ryland, C.R. Smith and P. Wass, unpublished observations, 1994) Parameter

River Ouse

River Trent

pH Alkalinity (meq 1 1) Mg (meq 1 1) Ca (meq l 1/ DOC (rag I 1) Phosphate (/aM) Temperature ('C) Flow(m3 s ~) SPM (rag I 1)

8.0 2.8 1.0 3.4 4.2 2.1 2 22 10 300 2 200

7.8 3.5 1.9 5.3 5.3 30.0 2 23 20 400 5-200

(R. Ouse) and 9 May 1995 (R. Trent). The rivers were both at relatively low flow, the discharge of the R. Ouse being 4 7 m 3 s 1 and that of the R. Trent being 41 m 3 s 1 (see Table 1). The concentrations of SPM were 7.9 mg 1 1 (R. Ouse) and 15.4 mg l 1 (R. Trent). The estuarine sample was taken at Goole Bridge (NGR SE 733266), at the turbidity maximum of the Ouse estuary (salinity 1.0%), on 16 October 1994; the concentration of SPM was 1400mg 1 1 To obtain gram quantities of SPM from the two river waters, several hundred litres of water were taken from each site. For the estuarine site, a sample of 10 1 proved sufficient. The samples were collected in polyethylene containers that had been cleaned with acid and thoroughly rinsed with distilled water. Before sampling, the bottles were rinsed twice with the natural water in question. The samples were returned to the laboratory within 8 h of sampling, and stored at 4°C before retrieval of the SPM. Solids were collected by centrifugation (Beckman J2-21), using a fixed-angle rotor with 350 cm 3 bottles for the estuarine sample (45 rain at 12000g) and a JCF-Z flowthrough rotor for the river samples (201 h 1 32000g). Recoveries of SPM were ~60% for the two river samples and greater than 90% for the estuarine sample. The lower recoveries for the river materials are proba-

J.R. Ferreiraet al./Colloids Surlhces A: Physicochem. Eng. Aspects 120 (1997) 183 19,~'

bly due to losses of slowly-sedimenting (lowdensity and/or small-diameter) particles that escaped sedimentation during their passage through the flow centrifuge. After concentration, the samples were washed by repeated (four times) suspension in distilled water and centrifugation. In the case of the River Ouse material, the SPM was suspended finally in 0.001 M NaC1, while for the other two samples distilled water was used. Final concentrations of SPM were 4.7 g 1 1 (River Ouse), 13.0 g 1 ~ (River Trent) and 70 g 1 ~ (Ouse estuary). These suspensions were each divided into 30 cm 3 aliquots, and then stored frozen until required.

3. Methods 3.1. Electron microscopy

Suspensions were shaken for l min, then subsamples were taken, diluted with distilled water and sonicated for 5 rain. Droplets of the dispersion were placed onto ionised Formvar-coated EM support grids, air-dried slowly and shadow-cast with chromium. The preparations were examined with a J E O L 100CX T E M S C A N instrument.

3.2. Elemental analysis

The C, H and N contents of the particles were determined with a Carlo-Erba Elemental Analyzer, Model 1106.

3.3. Inductively-Coupled Plasma-Mass Spectrometry ( I C P - M S )

A VG-elemental PQe instrument operating under standard conditions was used. maRh was used as an internal standard to compensate for matrix effects and instrumental drift. D a t a were analyzed with PQ VISION software, metal concentrations being determined by fully quantitative multi-element calibration.

185

3.4. Extraction~s

Suspensions of SPM in 0.01 M N a O H were prepared, mixed for 2 h, then centrifuged (30 min, 20000g) and the supernatants were analyzed for D O C with a Phase-Sep TOCSin II analyzer and for soluble reactive phosphorus by the method of Murphy and Riley [14]. Extraction with 0.1% HNO3 was performed, followed by centrifugation and the determination of dissolved metal concentrations by ICP-MS. To determine base cation contents of the SPM, samples were extracted with 0.1 M HCI, and dissolved metal concentrations were measured by atomic absorption spectrometry: interferences were suppressed with a 2% lanthanum solution. Extractions to determine the contents of Mn and Fe oxides were done with 0.1 M hydroxylamine0.01 M HNO3 and with 0.2M a m m o n i u m oxalate/0.2 M oxalic acid respectively, according to the methods of Shuman [ 15]. 3.5. A c i d base titrations

Suspensions of SPM were prepared at concentrations of 100-200 mg 1 1 in 0.001 M or 0.1 M NaCI in a 100 c m 3 glass bottle. Sufficient acid was added to achieve a pH between 3 and 4, the suspension was equilibrated by shaking overnight at 10 C , and then purged for 30min with CO2free wet air in order to expel CO 2. Titrations were carried out by adding known volumes of standardised CO,-free N a O H (0.05 0.1 M) from a Radiometer ABU-80 automatic burette, and measurement of pH with a radiometer glass electrode (GK2401Ct. Throughout the titrations, the suspension was blanketed with CO2-free wet air. After each addition of base, equilibration was allowed to take place and was judged to be complete when the millivoltmeter reading changed by less than 1 mV over a 2t1 min period. Because of the lengthy periods required for this to happen (up to 3 h in some cases), the number of titration points practically achievable in a working day was limited to between 5 and 10. 3.6. Electrophoresis

Measurements were made with a Rank Brothers Mark II instrument, using a flat cell at 20 C .

186

J.R. Ferreira et al./Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 183 198

Particles were suspended at ~ 10 mg 1 1 in 0.001 M NaC1, different amounts of NaHCO3 were added, equilibration with air was allowed to occur by shaking overnight, and then pH values were determined. Usually 10 particles were timed in each direction at both stationary levels, but about half this number of timings was made for samples with low mobilities.

3.7. Phosphorus sorption Suspensions of the particulate samples were prepared in a background electrolyte of 0.1 M NaC1, to which had been added different amounts of HC1 or NaHCO3. The suspensions were shaken overnight at 10°C, then centrifuged (18000g, 30 rain), and the supernatants were analyzed for orthophosphate [ 14].

3.8. Metal adsorption by centrifugation-depletion Experiments were carried out using 35 cm 3 polyethylene centrifuge tubes as reaction vessels. The tubes, and all other plastic and glassware used in the experiments, had been cleaned by soaking in 1% H N O 3 overnight, followed by an overnight soak in deionized water, and thorough rinsing in deionized water. Suspensions of SPM were prepared at a concentration of 100 mg 1-1 in a background electrolyte of 0.1 M NaC1, and a small volume of a solution of metals (Co, Ni, Cu, Zn, Sr, Cd, Cs, Eu, Pb) in 0.17 M HNO3 was added to each centrifuge tube, to give a final concentration of each metal of 1 gM. Differing amounts of NaHCO3 were then added, in order to obtain final pH values in the range 3 9. A small Tefloncoated magnetic stirrer bar was placed in each tube. Adsorption reactions were allowed to occur overnight at 10°C in a shaking incubator. After reaction, each tube was placed on a magnetic stirrer, and a 5 cm 3 sample of the suspension was taken, placed in a clean centrifuge tube, and acidified with 5 cm 3 of 0.28 M HNO3. The acidified sample was shaken overnight, centrifuged (18 000g, 30 rain), and the supernatant taken for the determination of total metal concentrations by ICP-MS. The suspension remaining in the original centrifuge tube was centrifuged, 5 cm 3 of the supernatant was

removed, acidified with 5 cm 3 of HNO3, and taken for the determination of metal concentrations in solution. A sub-sample of the supernatant was taken for pH measurement. By this procedure, account could be taken of any losses of metal onto the walls of the centrifuge tube during equilibration with the SPM; such losses were generally less than 10%, but were up to 40% for Cu, Eu and Pb in experiments with the estuarine SPM. Thus, the amounts of adsorption were obtained by the difference between the total metal concentrations determined on acidified samples and dissolved metal concentrations determined on unacidified samples, not by the difference between the original total added metal and the dissolved metal. The procedure also takes into account any metal associated with the particles, i.e. that present before the metal spike was added. The optical absorbances of the supernatants were determined, in order to estimate concentrations of dissolved organic matter released from the SPM.

3.9. Copper adsorption by potentiometry An Orion Cu electrode (No. 94-29) was used. It was calibrated with solutions of C u ( N O 3 ) 2 at concentrations in the range 1-100~tM. The response was linear, with a mV/loglo[Cu 2+ ] slope of 29-31, close to the theoretical value of 28.0. Experiments were conducted on solutions and suspensions of volume 900 cm 3, kept in a 1 1 beaker thermostatted at 20°C. The beaker was kept in a laminar flow cabinet, and was covered with a thick black cloth in order to exclude light, which affects the electrode performance. A lid with apertures for Cu and glass electrodes and a gas inlet was placed over the beaker. The solutions under study were bubbled continuously with wet air in order to maintain equilibrium with atmospheric CO2. To test the performance of the electrode, 1 0 - 2 M solutions of oxalic and citric acids were prepared in 0.1 M NaNO3 and acidified to p H i 3 . Then Cu(NO3)2 was added to give a total copper concentration of 0.1 or 1 gM. After allowing equilibration (judged to have occurred when the potential changed by less than 1 mV over 5 min), the potentials given by the glass and copper electrodes were recorded, and a small volume of N a O H was added

J.R. l-),rreira et al./Colloids Surjaces A: Physicochem. Eng. Aspect,~ 120 (1997) 183 19,~

187

to increase the pH. After re-equilibration the new millivoltmeter readings were taken, and the process repeated until the pH reached 8. The concentrations of [Cu 2+ ] measured during these experiments were compared with calculated values, using equilibrium constants given by Martell and Smith [16]. Acceptable agreement was found down to [Cu 2+ ] values of 1 nM. Experiments with suspensions of SPM from the Ouse estuary and the River Trent were carried out in a similar manner to those with the organic acids. Sediment concentrations were 100 mg 1 ~ and the total added Cu concentration was 1 btM. The lowest Cu 2+ concentrations recorded were greater than 1 nM, and therefore within the range of reliable performance of the Cu electrode.

4. Results Electron micrographs of the three SPM samples are shown in Fig. 1. The specimens from the two rivers are seen to include considerable amounts of material of biological origin, including diatom frustules and the remains of algal and bacterial cells, whereas the estuarine sample is dominated by platy clay mineral particles. A wide range of sizes of primary particles is evident. When suspended in the original waters, the particles are likely to be found in aggregates, ranging in size from 1 to

Ib)

100pm [17]. Table 2 shows compositional data for the three samples. They differ considerably in their organic matter contents: if it is assumed that the organic matter is 50% carbon, then the organic matter contents of the River Ouse and Trent particles are 24 and 54% respectively, whereas that of the Ouse estuary material is only 9%. The C : N ratio of the R. Trent material is 6.3, a low value characteristic of phytoplankton [18], and the alga Melosira was present in large amounts (J.G. Day, personal communication, 1995). In contrast, the value for the estuarine sample is 21.5, typical of well-humified organic matter [19]. The R. Ouse C : N ratio of 12 is consistent with a significant contribution from cultivated soils [19]. The base-extractable carbon is presumed to consist mainly of humictype material The amounts correspond to 10 15%

Icl Fig. 1. Electron micrographs of SPM frmn the Rivers ()use (a) and Trent (b), and the Ouse estuary lc). Tile scale bar represents I bm~.

of the total carbon for the two river samples, but the lack of detectable base-extractable C in the estuarine SPM indicates that any hnmic matter present is very strongly held by the clay minerals that make up most of this sample. All three samples contain appreciable amounts of Mn and Fe oxides.

J.R. Ferreira et al./Colloids SurJhces A: Physicochem. Eng. Aspects 120 (1997) 183 198

188

Table 2 Bulk compositional data for samples of SPM. The C and N values are total contents, other data refer to extractions with N a O H (Cextr), hydroxylamine (MnO2), oxalate (Fe(OH)3), HC1 (base cations), N a O H (orthophosphatet and H N O 3 (trace metals); see the text for details. The value of the Na content for the R. Ouse S P M is given in parentheses because the sample had been stored in a solution of 0.001 M NaCI; the quoted Na content was used with Eq. (1) in order to estimate particle charge, but does not reflect natural conditions Component gg

R. Ouse

R. Trent

Ouse estuary

1

C Cextr N Mn as M n O 2 Fe as Fe(OH)3

0.120 0.013 0.010 0.004 0.029

0.272 0.042 0.043 0.005 0.015

0.043 <0.001 0.002 0.002 0.048

meq g 1 Na Mg K Ca

(1.15) 0.48 0.13 1.63

0.10 0.14 0.03 0.52

0.01 0.40 0.01 2.14

gmol g P Ti V Cr Co Ni Cu Zn Sr Cd Pb

90 ND ND 0.5 0.4 0.6 2.0 19 1.6 <0.01 5.3

141 0.3 0.2 0.5 0.5 1.0 1.1 11 1.1 0.02 0.9

22 2.0 1.0 0.5 0.2 0.4 0.5 3.6 1.1 <0.01 0.6

Both the Ouse and Trent rivers are high in dissolved phosphate, as shown by the data in Table 1. Concentrations in the estuary are also high [ 11 ]. Therefore, the high phosphate contents of the particles are to be expected. The trace metal contents of the particles were determined by extraction with 0.1% HNOB, this relatively mild reagent being chosen to provide an estimate of the metal that might be released from particles subjected to different environmental conditions. The results show that the two riverine SPM samples have higher trace metal contents than do the estuarine particles; this indicates that metals delivered by the rivers, or discharged directly into the estuary, are

substantially diluted with relatively large amounts of marine particles of low metal content. Electrophoresis results (Fig. 2) show that each of the SPM samples has a net negative potential in the pH range 3-9. This is in line with previous observations of natural particulate matter [20]. The electrophoretic mobilities of the three samples are broadly similar, but there are significant differences. The mobilities of the R. Ouse particles vary most with pH, and attain the highest (most negative) values at high pH, while the Ouse estuary particles maintain relatively high mobilities at low pH. The R. Trent SPM displays intermediate behaviour. The mobilities reflect the potentials of the particles at the plane of electrokinetic shear, which are governed by a combination of factors, notably the sorption reactions of protons, metal ions and natural organic matter, and the surface topography [21,22]. A different, more quantitative, picture of the particle surface chemistry can be obtained by acid-base titrations, which give information about proton dissociation from the particles in their entirety, not just at the shear plane. However, interpretation of the results needs to take into account the presence of base cations and/or acid anions associated with the particulate matter under study. From the electrophoresis results, it is evident that each of the particle preparations bears a negative charge. Therefore, at the start of a titration, there will be an excess of balancing base cations in the suspension. The charge balance for the system during the acid-base titration can therefore be represented by Ta + T* + [ H +] - TA -- [ O H - ] + Z [ S P M ] = O (1)

where TB is the concentration of added base, T] is the concentration of base already present, TA is the concentration of added acid, Z is the particle proton charge (eq g 1), [SPM] is the concentration of particles (g 1 1) and square brackets indicate concentrations. To estimate T*, the known amounts of the particles were extracted with acid, and base cations in solution were determined (Table 2) note that this will contain any base cations liberated by the dissolution of mineral phases, e.g. CaCOa. The concentrations of protons

J.R. Ferreira et al./Colloids Sur/ilces A. Phvsicochem. Eng. ,4spect,~ 120 (1997) 183 19,~

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and hydroxyl ions at each point in the titration were obtained from the pH readings, taking activity coefficients into account using the extended Debye Ht~ckel equation. Since the added amounts T~, TA and [ S P M ] are known, Z can be calculated from Eq. ( 1 ). It should be noted that the Z values estimated and plotted are in terms of protons. They do not include the contribution from adsorbing base cations. At the start of the titration (low pH) these will be in solution but, as the pH is increased, they (divalents especiallyl will adsorb to the particles, displacing protons. Thus the net surface charge will be somewhat lower (less negative) than those calculated. However, the total concentrations of divalent cations are small, and so only a small influence on particle charge is expected. Values of Z as a function of pH are plotted in Fig. 3. It is seen that the R. Ouse and ()use estuary particles bear a negative proton charge over the pH range studied. However, the samples differ markedly in that the R. Ouse material has ~ 2 meq g 1 of pH-dependent charge titrating between pH 4 and 9, approximately four times that of the estuarine sample. It can be concluded that the river material bears a range of weakly acidic groups, while the estuarine particles are character-

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190

J.R. Ferreira et al./Colloids Surfaces d." Physicochem. Eng. Aspects 120 (1997) 183 198

ised by a fixed negative charge. The sample from the R. Trent has low net charge, and appears to bear a positive charge at low pH (Fig. 3). However, given that the correction for base cations described above depends upon extraction and analysis, and bearing in mind the small charges measured by the titration, it is perhaps not justifiable to put too much weight on the absolute values of Z. More important is the observation that the R. Trent particles bear approximately 1 meq g 1 of pH-dependent charge, due to weak acid groups. In this respect, the R. Trent particles are more like the R. Ouse material than the estuarine material. Each of the three natural SPM samples gave a measurable dependence of charge on ionic strength (I), the magnitude of Z being greater at I = 0 . 1 M than at I=0.001 M for a given pH (Fig. 3). This is usual behaviour for charged surfaces in aqueous solutions, and arises because the indifferent electrolyte shields the electrostatic attraction of the surface for the dissociating protons [23]. The dependence of Z on I is best defined for the R. Ouse particles, and is of similar magnitude to that seen for oxides and humic substances. The interactions of the particles with orthophosphate were investigated by determining the release of orthophosphate from the isolated samples as a function of pH. The results are shown in Fig. 4, expressed in terms of the total contents of orthophosphate, determined by extraction with 0.01 M N a O H . The same general trends were found for all three samples, orthophosphate being released into solution at both low and high pH values. A plot representing the data of Edzwald et al. [24] for a sample of illite is included for comparison (see Section 5). Results from metal adsorption experiments, carried out by the centrifugation-depletion method, are represented by the points in Fig. 5. Generally, it can be seen that each SPM sample gives similar trends, the fraction of metal in solution falling with pH, in some cases rising again at pH values greater than 7. The observed metal adsorption behaviour is qualitatively similar to results for simpler surfaces such as metal oxides (see e.g. Ref. [2]), and can be interpreted in terms of the interaction of metal aquo ions (i.e. Cs +, Co 2+, Eu 3+, etc.) with deprotonated sites on the surfaces of the SPM.

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Fig. 4. Sorption of orthophosphate by the three SPM samples. The experiments were conducted without the addition of extra orthophosphate. The dotted line represents the results of Edzwald et al. [24] for illite. Concentrations: R. Ouse, SPM 0.21 g 1~1 particles, 19 gM orthophosphate; R. Trent, SPM 0.26g 1 1 37p.M orthophosphate; Ouse estuary, SPM 0.46g 1 ~, 10HM orthophosphate, illite 5 g 1 i 650gM orthophosphate.

Competition between the metal ions and protons for the sites accounts for the pH dependence. We have used a simple model, similar to that of Mouvet and Bourg [25], in order to compare the results among metals for adsorption. The metal binding sites on the particle surfaces are represented by S, and are assumed to interact with protons and metal aquo ions as follows: S + H + = SH

(2)

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(3)

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(4)

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(5)

If the number of S sites per gram of SPM is n, and [ S P M ] denotes the particle concentration (g 1- 1 ),

J.R. Ferreira et al./Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 183 198 1.2

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Fig. 5. Results of metal SPM adsorption experiments. The fraction of total metal present in the dissolved phase (centrifugation supernatant) is plotted against pH: (OI R. Ouse SPM, ( e ) R. Trent SPM, (~2) Ouse estuary SPM. The concentration of SPM was 100 mg I ~ in each case, and the starting concentration of metal was 1 pM (see text). The lines are model fits ( Eqs. 2 71: see Table 3.

J.R. Ferreira et aL /Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 183-198

192

then it follows that: [SM] -

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n K [ S P M ] [M=+]/[-H + ] 1 + K [M=+]/[H + ]

Sr,Cs

(6)

When the occupancy of sites by the metal is low, the term K [ M = + ] / [ H +] is much less than one, and Eq. (6) simplifies to [SM]=nKESPM]

[ M = + ] / [ H +3

(7)

Thus, the concentration of adsorbed metal can be calculated from the three variables [M=+], [ H + ] and [ S P M ] , together with the constants n and K. Under these trace conditions n and K can be combined into a single constant, which we shall denote nK. Because this parameter is a combination of site density and affinity, it can only give an overall representation of the metal-binding properties of the particles, so that a sample with a few strong sites might be indistinguishable from one with many weak sites. It is admitted that the model does not take into account a number of potentially important factors, including competition amongst metals, the adsorption of hydrolysed metal species, e.g. C u O H +, site heterogeneity, the presence of adsorption sites with different proton exchange stoichiometries, significant site occupancy (nontrace binding) and electrostatics. However, if these deficiencies are borne in mind, the model is helpful for making comparisons among different samples of SPM and different metals. In the experiments, the total concentrations of dissolved metals were determined, and from these it is necessary to extract values of [ M =+ ], in order to apply Eq. (7). For the experimental conditions of the present work, complexation reactions with CI-, O H - , H C O 3 - and CO 2 have to be taken into account. This was done with the WnAM chemical speciation model [26], which was used to calculate the fraction of dissolved metal present as M =+ ; the results of such calculations are shown in Fig. 6. It is seen that complexation can make this fraction very small for some of the metals (Cu, Eu, Pb) at higher pH. To fit the data in Fig. 5 with the model, residuals (differences between observed and calculated values) in the fraction of total metal in the dissolved form were considered. The value of log n K was

1 I~ ,~ "~ o

0.1 0.01

0.o01 0.0o01

I

[

I

I

6

7

8

9

10

pH Fig. 6. Dependence on pH of the fraction of dissolved metal present as the aquo ion. The full curves were calculated with WHAi~ [26], taking account of complexation by O H - , CI and carbonate species. The dotted curve refers to a calculation in which fulvic acid (FA), at a concentration of 5 mg I ~, was included (see text).

adjusted until the sum of the squares of the residuals was minimised. The fitted lines are shown in Fig. 5. Generally, the model represents the data adequately, with nearly all root-mean-squared deviations being less than 0.1. Values of log n K are presented in Table 3. They vary somewhat amongst the particle types, adsorption by the R. Ouse sample being greater than by the other two, which Table 3 Values of the - l o g n K for the samples of SPM. The italicised values for the R. Trent refer to fits taking into account the presence o f 5 m g 1 l offulvic acid (2.5rag 1 1 DOC) in the dissolved phase; the value for Eu estimated in this way is unreliable (see text). The smaller - l o g nK, the stronger is the binding of the metal in question Metal

Co Ni Cu

Zn Sr

Cd Cs Eu Pb

R. Ouse 5.9 5.9 4.2 8.4 5.8 >10 2.3 2.8

R. Trent 7.2 6.9 5.5 6.3 9.0 6.7 >10 2.9 3.7

7.1 6.7 3.8 6.0 9.0 6.6 >10 (0.7) 3.5

Ouse estuary 7.1 7.1 4.6 8.3 6.9 >10 3.3 3.0

J.R. Ferreira et aL/Colloids Sur~wes A: Physicochem. Eng. Aspects I20 (1997) 183 19~'

are approximately equal (but see below). It is found that in each case the extent of adsorption of the metals follows essentially the same trend, i.e. C s < S r < C o m N i ~ C d < Z n < C u < P b < E u . The weak binding of Cs indicates the absence of sufficient specific edge sites on e.g. illite [27] in any of the three samples. As an alternative way of looking at adsorption, an ion-specific electrode (ISE) was used to study copper uptake by the R. Trent and Ouse estuary particles. This method provided direct estimates of the concentration of Cu 2+, rather than the total concentration of dissolved metal as obtained in the centrifugation-depletion method. The results, together with model fits, are shown in Fig. 7. Again the model works reasonably well, but the pH dependence is imperfectly described, a consequence of adopting a metal-proton exchange ratio of unity (Eq. (411. In the case of the Ouse estuary SPM, the value of log t~K obtained from the ISE measurements ( - 4 . 9 ) was close to that from centrifugation-depletion (-4.6). However. for the R. Trent material, the ISE results value was substantially less negative ( - 3 . 6 as against - 5 . 5 ) , indicating apparently stronger binding. A possible explanation for this is

I

I

q

I

.-£ o Q.

o

._.Trent

•

Ouse

,

I

%"

estuary

o \ o

10

4

i

I

i

5

t

6

i

I

7

,

8

pH Fig. 7, Adsorption of Cu by SPM, studied with an ISE; p[Cu] is the negative logarithm of the concentration of Cu e+ . The total concentrations of SPM and Cu were 1 0 0 m g 1 * and 1 2 laM respectively. The lines are model fits (Eqs. 2 7), with IognK3.6(R. Trent S P M ) a n d - 4 . 9 ( O u s e e s t u a r y S P M } .

19~

the release from the R. Trent particles of dissolved organic matter. Thus, supernatants from the centrifugation depletion experiments with the R. Trent particles, but not with either the R. Ouse or Ouse estuary materials, had measurable optical absorbance ( ~ 0.02 at 340 nm, in a 1 cm pathlength cell p. If all the absorbance were due to humic substances, this value would correspond to a concentration of 1 mg C I 1, based on results for isolated humic substances [28] and for dissolved organic matter in surface waters [29]. However, the organic matter extracted from the R. Trent sample with base (see Table 2) gave a lower absorbance/carbon ratio than was obtained from these previous studies, and the concentration of DOC estimated on this basis is ~ 4 nag 1 1 in order to determine the influence on the value of log nK of Cu complexation by the dissolved organic matter, it was assumed that the concentration of carbon due to fulvic acid (FAt in the supernatant was 2.5 mg 1 1 a value intermediate between the two estimates of the DOC concentration. The data for Cu adsorption by the R. Trent SPM were re-analyzed using WHAM to take metal FA interactions into account. This led to a less negative log ~TK value of 3.8 for the centrifngation-depletion data, but the ISE value was unchanged at -3.6. Thus, complexation of Cu by dissolved organic matter seems to explain the discrepancy between the resuhs from the two experimental methods. The values of log ~IK l\)r other metals v~ere also affected (Table 3) bul, with the exception of Eu, less than those found for Cu. The new log nK value for Eu cannot be considered reliable, however, because this metal is calculated to be very strongly complexed by FA, making estimates of the concentration of Eu 3 " unreliable. The metal-proton exchange model was used to compare adsorption by the three samples studied in the presenl work with results tk~r pure solids, and other natural SPM. To do this, we determined log i~K values from published resuhs iTable 4). The data used for this exercise were restricted to those obtained at high solid metal concentration ratios, in order to be relevant to our experiments and to natural conditions. The use of the model allows data obtained under different conditions of concentration, pH and background electrolyte to be compared. Howe~er, because the model is so simple,

194

J.R. Ferreira et al./Colloids Surfaces A." Physicochem. Eng. Aspects 120 (1997) 183 198

Table 4 Comparison of - l o g n K values estimated from published data and from the results of this study. The data all refer to high particle/metal ratios (> 3 x 104 g tool-~). The oxide specific areas are all normalized to 100 m 2 g-~ (see text). A surface area for the illite sample was not quoted, and so the - l o g nK value refers to the published data on a weight basis. The value in parentheses for Pb adsorption to MnO2 is estimated from results at a lower particle/metal ratio (6 x 103g t o o l - l ; Ref. [32]). The value in parentheses for Cd adsorption to MnO2 is estimated on the basis that Zn and Cd show similar adsorption affinities at low particle/metal ratios (6 x 10/ g mol-~; Ref. [33]), and from the derived value for Zn. Values for HA were calculated with the WHAM speciation code [26]. The illite and Susquehanna River S P M results refer to an artificial river water containing Mg and Ca (I~0.004 M). Two sites on the R. Glatt were sampled, and the ranges of values refer to six sampling dates. The smaller -lognK, the stronger is the binding of the metal in question Sample

Ref.

Medium

c~-SiO2 AI203 MnO 2 ~/-MnO2 MnO 2 ct-FeOOH Fe(OH)3 Illite HA

Chlamydomonas

[30] [30] [31] [32] [33] [30J [34] [35] [26] [36]

0.1 M NaNO3 0.1 M NaNO3 0.01 M NaNO3 0.01 M NaNO3 0.02 M NaNO3 0.1 M NaNO3 0.1 M NaNO3 SRW 0.1 M NaNO3 0.01 M KNO3

R. Glatt (1) R. Glatt (2) R. Rhine Susquehanna R.

[7] [7] [37] [35]

R. Ouse R. Trent Ouse estuary

Cu

Zn

1.9

3.3 2.9

4.9

6.3

2.7 4.7

4.2

0.01 M K N O 3 0.01 M K N O 3 0.05 M NaNO3 SRW

-

4.4-5.2 4.5-5.0 -

0.1 M NaCI 0.1 M NaCI 0.1 M NaC1

4.2 3.8 4.6

care must be taken not to over-interpret the derived values of log nK. For the pure oxides, actual solids concentrations (g1-1) were "normalised" by assuming a specific surface area (SSA) of 100 m 2 g-1. Thus, an oxide with S S A = 4 0 m 2 g-1 at a concentration of 0.5g 1-1 would have a "normalised" concentration of 40 x 0.5/100= 0.2g 1-1. By this procedure, data for solids of different surface area can be compared directly, on a mass-for-mass basis.

5. Discussion

Qualitatively, the samples obtained from the Rivers Ouse and Trent are similar. They both contain substantial amounts of organic matter, have a net negative surface potential, bear groups from which protons dissociate in the pH range

6.0 -

Cd 8.5 7.0

Pb (1)

(3) 6.8 6.5 7.5 6.0 6.3

3.9 3.4

3.4 4.4 3.4 3.9 5.7 7.0 5.8 6.6 6.9

2.8 3.5 3.0

4-9, and release orthophosphate at low and high pH. Their selectivities towards metal binding are the same. Quantitatively, they are somewhat different, the R. Ouse material having about twice the content of ionisable groups, and a greater overall affinity for metals. Consideration of the results for the sample of SPM from the turbidity maximum of the Ouse estuary reveals further differences among natural particles. The organic content is much lower, and the organic matter is much more humified and resists extraction by base. Clay minerals are the dominant particulate material, and this makes for a significant difference in surface charge properties, with very little dependence on pH. Metal binding by the estuarine sample is found to be the weakest of the three samples examined, and phosphate release at low pH is relatively less than that found for the river samples.

195

J.R. Ferreira et al./Colloids Surfaces A. Physicochem. Eng. Aspect~s 120 (1997) 183 19~¥

The present results, and those reported by others for particulates from other rivers [7,35,37] show that the chemical properties of natural SPM vary from site to site and with time. This must be due to variations in the types and/or relative amounts of the different particles that comprise the natural particle assemblage. It is therefore of interest to consider which particle types that comprise natural SPM might be responsible for observed chemical properties, in this instance with regard to charge characteristics, phosphate binding, and metal binding. Likely contributors to the observed pHdependent charge are the oxides of A1, Si, Mn and Fe, humic acid, and algae. Literature data show that for oxides with surface areas of 100 m 2 g L, changes in surface charge between pH 3 and pH 8 would typically be in the range 0.5-1 meq g [38-43], while for humic acid the value is 3-4 meq g i [44]. In the cases of 8i02, MnO2 and humic acid the charge would be negative at both pH values, whereas for iron oxides the particles would bear a net positive charge at pH 3 and a zero to negative one at pH 8. Xue et al. [36] reported that the proton charge on the green alga C h l a m y d o m o n a s varied from - 0 . 5 meq g-~ at pH 4 to - 2 . 3 meq g--i at pH 9. Other algae studied by Crist et al. [45] showed changes in (negative) charge between pH 4 and 9 of 0.5-3 meq g - l , although absolute values were not reported. Aluminosilicate clay minerals will contribute charge that is independent of pH. Examples are given by Rowell [46], and range from 0.1 meq g for kaolinite and chlorite to 1 meq g- t for smectite and 1.5 meq g ~ for vermiculite. The charges per gram quoted here are similar to the net charges of the three samples of SPM that we have studied, and so a variety of combinations of different "pure phases" can be envisaged to account for the observed values. The high organic contents of the two river samples point to major contributions to the charge from algae and other organic components, including humic matter and particles derived from sewage treatment. Aluminosilicates and SiO2 may also be important. Because manganese and iron oxides comprise between them only a few percent of the total weight, it is unlikely that they contribute significantly. The sample from the Ouse

estuary has little pH-dependent charge, and it can be concluded that the observed charge is principally due to aluminosilicate clay minerals. Our results for the pH-dependent adsorption of orthophosphate by the SPM samples are more consistent with the results of Edzwald et al. [24] for illite (Fig. 4) than with the model of Fox [47], in which riverine-dissolved orthophosphate concentrations are controlled by the equilibrium with a solid solution of ferric phosphate in ferric hydroxide. Fox's theory, which has been shown to account for orthophosphate concentrations in a variety of rivers, and the more general process of phosphate adsorption by iron oxides (see e.g. Ref. [43]), require that adsorption strength increases as the pH is decreased, while the reverse is the case for the particles under study here. This is illustrated by Fig. 8, which shows the predictions of Fox's theory for the three samples, based on their orthophosphate and iron oxide contents (Table 2). Beck & Van Riemsdijk [48] state that clay minerals adsorb orthophosphate via ligand exchange with OH groups coordinated to surface AI atoms, while Edzwald et al. [24] suggested that adsorption of orthophosphate by illite also depended upon Fe

i

i

...

~

i

R.Ouse R.Trent Ouse estuary

".,...

[

]

i

o --e-.... o . .

'-....

7

' " '.......

IEL a. 6

h

I

k

I

I

I

I

3

4

5

6

7

8

9

10

pH Fig. 8. Comparison of experimental results for orthophosphate adsorption by SPM samples (points) with predictions from the model of Fox [47] {lines);PPT is the negative logarithm of the total dissolved phosphate concentration. The predicted values were calculated taking into account the varying Fe/P ratio of the solid phase.

196

ZR. Ferreiraet al./Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 183-198

associated with the clay surface. It is not clear, however, how these mechanisms would account for the observed pH dependence of orthophosphate adsorption by the three samples studied here. It should also be noted that the pH dependences for the three solids are somewhat different, and all of them exhibit maximum adsorption at a higher pH than does the illite sample. Another candidate surface for phosphate adsorption is CaCO3, but adsorption in this case increased with pH above pH 7 [49]. We conclude that clay minerals are the most likely adsorbents for phosphate in the samples studied, in terms of controlling solution concentrations; the high levels of phosphate in the rivers and estuary may mean that iron oxide phases are present in amounts too small to be important in this respect. Table 4 presents values of log n K for four commonly-studied metals adsorbing to a variety of solids. If we consider first Cd, it is seen that the natural SPM sa'mples are stronger adsorbents on an equal weight basis (less negative log nK values) than both S i O 2 and illite, which suggests that neither of these pure phases is significant in the net metal adsorption properties of the natural materials. A similar conclusion can be drawn for alumina and iron oxides; these phases generally comprise only a few percent of natural SPM, so unless they have much greater surface areas than the I00 m 2 g 1 assumed for the calculations of Table 4, they will not be significant in the adsorption of Cd. In the case of humic acid (HA), surface area considerations do not apply, since different isolates of this material display very similar binding capacities towards metals E44], and so unless HA is present at high levels in the particulate matter it will not be significant. It is not possible to rule out algae, as represented by Chlamydomonas in Table 4, as significant sorbents of metals, although their contribution must vary with the season. We estimate that MnO2 adsorbs Cd quite strongly (Table 4), and this phase is likely to be a significant binder of Cd in particle assemblages. Turning to Cu, and applying the same argument as for Cd, it is seen from Table 4 that both MnO 2 and HA bind this metal significantly more strongly than do the natural SPM samples. Therefore, if one or both comprise a few percent of natural

materials, the overall metal adsorption properties might be accounted for. Again, adsorption by Fe oxides appears too weak, while the contribution of algae, if present in sufficient amounts, may be significant. For Zn, the values of log nK suggest that MnO2 may be important, and possibly HA. For Pb, MnO2 is the most likely candidate, although this is based on an extrapolation of adsorption data at relatively low oxide/metal ratios (see Table 4). The above reasoning suggests, on the basis of results for pure components studied in isolation, that MnO2, HA and algae are the most likely components of SPM to explain metal-binding properties, and that silica, alumina, iron oxides and clay minerals are unlikely to be significant. This is advanced as a reasonable, and hopefully testable, hypothesis rather than a definite conclusion. Modelling studies, along the lines suggested by Luoma and Davis [9], are needed to predict the contributions of different solid phases to the metal-sorbing properties of SPM assemblages. The issue of interactions among the different solid phases must also be addressed, although there is evidence to suggest that mixtures behave additively E50,5l ]. An important goal is to explain the variation of log n K among different samples of natural SPM (Table 4), a feature which is likely to make simple models of the kind used here unsatisfactory for general environmental application. Taken together, the above arguments suggest that there is probably no strong relationship between surface charge and trace metal binding, although this might not be so if organic materials (algae, sewage particulates, HA) are dominant. If MnO z were the dominant metal-binding phase, then its low levels in SPM would rule out a relationship between the sites responsible for overall charging behaviour and those involved in metal adsorption. This does not mean, however, that modelling would be able to ignore the highcapacity, low-affinity sites since they would be important for major cations, especially Mg z+ and Ca 2+, the interactions of which need to be understood in order to take competition reactions into account. Furthermore, the interpretation and prediction of particle aggregation requires knowledge of surface charge characteristics.

J.R. Ferreira et aL/Colloids Surfaces A: Physicochem. Eng. Aspect.s 120 1997) 183 19,'¢

One of the chief aims of LOIS is to measure and model the fluxes of materials (including SPM, orthophosphate and metals) from land to sea [ 10], and the programme includes monitoring exercises to determine riverine and estuarine fluxes. The present findings are relevant to the interpretation of the results that will emerge. From measurements of dissolved phosphorus at or near the times of sampling, the contribution of adsorption by SPM to total phosphate concentrations was calculated. We find that in the Rivers Ouse and Trent only 8% and 2% respectively of the phosphate is present adsorbed to particles, whereas for the estuary the value is 70%. This calculation suggests that SPM may exert control over phosphorus transport in the low-salinity region of the estuary, but this will depend upon sediment dynamics, flushing times and possibly chemical kinetics [4,52,53]. With regard to metal behaviour, the chemistry is complicated not only by changes in the adsorption properties of SPM, but also by complexation reactions in the dissolved phase, notably those involving chloride and dissolved organic matter [37,54], adsorption competition effects, especially with Ms- , and redox transformations. Thus there are a number of issues which need to be considered before the influence of SPM on transport of metals can be properly determined. Nonetheless, the much higher concentrations of SPM in the low-salinity zone of the Humber estuary suggest that most metals will tend to pass from the solution to the solid phase. Again, sediment behaviour, hydrodynamics and chemical kinetics will be influential.

Acknowledgements Thanks are due to M.R. Williams and G.E. Millward for helpful discussions and to D.V. Leach, A.C. Pinder and M.R. Williams for field assistance. J.R.F. was supported by a European Union Marie Curie Fellowship, and by the Fundagao de Amparo a Pesquisa do Estado de Sgto Paulo (Brazil). J.M.B.'s contribution was made during an M.Sc. course at the University of Reading. We are grateful to the UK Natural Environment Research Council for a Special Topic Award within the Land-Ocean

197

Interaction Study LOIS). This is LOIS publication number 84.

References [ l ] W. Salomons and U. F/3rstner, Metals in the Hydrocycle, Springer-Verlag, Berlin, 1984. [2J W. Stumm, Chemistry of the Solid Water Interface, Wiley, New York, 1992. [3] D. Eisma, Suspended Matter in the Aquatic Environmcnt, Springer-Verlag, Berlin, 1993. [4] G.E. Millward, Analyst, 120 I1995) 609. [5] W. Stumm t Ed.), Aquatic Surfitce Chemistry. Wiley, New York. 1987. [61 B.D. Hone.~man and P.H. Santschi. Environ. Sci. Technol., 22 (1988} 862. [7] B. Mtiller and L. Sigg, Aquat. Sci., 52 11990) 75. [8] W. Sung, Environ. Sci. Technol., 29 (1995} 13(}3. [9] S.N. kuoma and J.A. Davis, Mar. Chem., 12 [ 19831 159. [ 10] 1.and Ocean Interaction Study (LOISI, Implementation Plan, Natural Environment Research Council, Swindon, UK, 1994. [11] A.L.H. Gameson led.), The Quality of the Humber Estuary, Yorkshire Water Authority, Leeds, UK, 1982. [12] J.S. Pethick. in N,V. Jones (Ed.), A Dynamic Estuary Man, Nature and the Humber, ttull tlniversitv Press, Hull, UK, 1988. [13] J.J.G. Zwolsman, North Sea Estuaries as Filters for Contaminants. Delft Hydraulics Report to the National Institute for Coastal and Marine Managemenl. Delft, NL, 1994. [14J J. Murphy and J.P. Riley, AnaI. Chim. Acta. 27(1962)31. [15] L.M. Shuman, Soil Sci. Soc. Am., J., 46 119821 1099. [ 16] A.E. Martell and R.M. Smith, Critical Stability Constants, Vol. 3, Plenum Press, New York, 1977. [17] E. Tipping. C. Woof and K.J. Clarke, ltydrol. Proc., 7 (1993) 263. [ 18] C . S . Reynolds, The Ecology of Freshwater Phytoplankton, Cambridge University Press, Cambridge, UK, 1984. [19] R.E. White. Introduction to the Principles and Practice of Soil Science, 2nd edn., Blackwell, Oxford, UK, 1987. [20] K.A. Hunter and P.S. Liss. Nature, 282 ( 19791 823. [21] E. Tipping and D. Cooke, Geochim. Cosmochim. Acta. 46 ( 19821 75. [22] R. Beckelt, in B.T. Hart lEd.j, Water Quality Management: The Role of Particulate Maner in the Transport and Fate of Pollutants, Water Studies Centre, Chishohn Institute of Technology, Melbourne, Vic.. 1986, pp. 113 142. [23] C. Tanford, Physical Chemistry of Macromolecules, Wiley, Nex~ York, 1961. [24] J.K. Edzwald. D.C. Toensing and M,C. Leung, Environ. Sci. Technol., 10 (1976) 485. [25] C. MouvetandA.C.M. Bourg, Water Res.,17 (19831641.

198 [26] [27] [28] [29] [30]

[31]

[32] [33] [34] [35] [36] [37] [38]

[39] [40]

J.R. Ferreira et al./Colloids Surfaces A. Physicochem. Eng. Aspects 120 (1997) 183-198 E. Tipping, Comp. Geosci., 20 (1994) 973. B.L. Sawhney, Soil Sci. Soc. Am., Proc., 31 (1964) 183. E. Tipping, Chem. Geol., 33 (1981) 81. E. Tipping, J. Hilton and B. James, Freshwater Biol., 19 (1988) 371. M.M. Benjamin and J.O. Leckie, in R.A. Baker (Ed.), Contaminants and Sediments, Vol. 2, Ann Arbor Science, Ann Arbor, MI, 1980. B.A. Dempsey and P.C. Singer, in R.A. Baker (Ed.), Contaminants and Sediments, Vol. 2, Ann Arbor Science, Ann Arbor, MI, 1980. J.G. Catts and D. Langmuir, Appl. Geochem., 1 (1986) 255. R.R. Gadde and H.A. Laitinen, Anal. Chem., 46 (1974) 2022. M.M. Benjamin and J.O. Leckie, J. Colloid Interface Sci., 79 (1981) 209. J.D. Reid and B. McDuffie, Water, Air, Soil Pollut., 15 (1981) 375. H.B. Xue, W. Stumm and L. Sigg, Water Res., 22 (1988) 917. M.A.A. Paalman, C.H. Van der Weijden and J.P.G. Loch, Water, Air, Soil Pollut., 73 (1994) 49. C.P. Huang, in M.A. Anderson and A.J. Rubin (Eds.), Adsorption of Inorganics at the Solid Liquid Interface, Ann Arbor Science, Ann Arbor, MI, 1981. Th.F. Tadros and J. Lyklema, J. Electroanal. Chem., 17 (1968) 267. D.E. Yates and T.W. Healy, J. Colloid Interface Sci., 55 (1976) 9.

[41] J.W. Murray, Geochim. Cosmochim. Acta, 39 (1975) 505. [42] R.J. Atkinson, A.M. Posner and J.P. Quirk, J. Phys. Chem., 71 (1967) 550. [43] D.A. Dzombak and F.M.M. Morel, Surface Complexation Modelling: Hydrous Ferric Oxide, Wiley, New York, 1990. [44] E. Tipping, Colloids Surfaces, 73 (1993) 117. [45] R.H. Crist, J.R. Martin, D. Carr, J.R. Watson, H.J. Clarke and D.R. Crist, Environ. Sci. Technol., 28 (1994) 1859. [46] D.L. Rowell, Soil Science: Methods and Applications, Longman, Harlow, UK, 1994. [47] L.E. Fox, Geochim. Cosmochim. Acta, 53 (1989) 417. [48] J. Beek and W.H. Van Riemsdijk, in G.H. Bolt (Ed.), Soil Chemistry, B. Physico-Chemical Methods, Elsevier, Amsterdam, 1982. [49] W.A. House and L. Donaldson, J. Colloid Interface Sci., 112 (1986) 309. [50] R.S. Altmann and J.O. Leckie, in T.P. O'Connor, W.V. Burt and I.W. Duedall (Eds.), Oceanic Processes in Marine Pollution, Vol. 2: Physicochemical Processes and Waste in the Ocean, E. Krieger, Malabar, FL, 1987, Chapter 13. [51] J.M. Zachara, C.T. Resch and S.C. Smith, Geochim. Cosmochim. Acta, 58 (1994) 553. [52] A.W. Morris, A.J. Bale, R.J.M. Howland, G.E. Millward, D.R. Ackroyd, D.H. Loring and R.T.T. Rantala, Water Sci. Technol., 18 (1986) 111. [53] A.W. Morris, Sci. Tot. Environ., 97/98 (1990) 253. [54] A. Ledin, S. Karlsson, A. Dtiker and B. Allard, Radiochim. Acta, 66/67 (1994) 213.