Complexes of the antimicrobial ciprofloxacin with soil

COMPLEXES OF THE ANTIMICROBIAL CIPROFLOXACIN WITH SOIL, PEAT,AND AQUATIC HUMIC SUBSTANCES

LUDMILLA ARISTILDE*y and GARRISON SPOSITOzxyBiological and Environmental Engineering, Cornell University, Ithaca, New York, USA

zMolecular Toxicology Group, University of California at Berkeley, Berkeley, California, USAxDivision of Ecosystem Sciences, University of California at Berkeley, Berkeley, California, USA

(Submitted 9 January 2013; Returned for Revision 11 February 2013; Accepted 22 February 2013)

Abstract: Natural organic matter (NOM) is implicated in the binding of antibiotics by particles in soils and waters. The authors’ previouscomputational study revealed structural rearrangement of both hydrophilic and hydrophobic moieties of NOM to favor H-bondingand other intermolecular interactions, as well as both competition with ion-exchange reactions and bridging interactions by NOM-bounddivalent cations. The importance of these interactions was investigated using fluorescence-quenching spectroscopy to study the adsorptionof ciprofloxacin (Cipro), a fluoroquinolone antibiotic, on 4 reference humic substances (HSs): Elliott soil humic acid (HA), Pahokee peatHA, and Suwannee river HA and fulvic acid. A simple affinity spectrum HS model was developed to characterize the cation-exchangecapacity and the amount of H-bond donor moieties as a function of pH. The adsorption results stress the influence of both pH conditionsand the type of HS: both soil HA and peat HA exhibited up to 3 times higher sorption capacity than the aquatic HS at pH ! 6, normalizingto the aromatic C content accounted for the differences among the terrestrial HS, and increasing the concentration of divalent cations led toa decrease in adsorption on aquatic HA but not on soil HA. In addition, the pH-dependent speciation models of the Cipro–HS complexesillustrate an increase in complexation due to an increase in deprotonation of HS ligands with increasing pH and, at circumneutral andalkaline pH, enhanced complexation of zwitterionic Cipro only in the presence of soil HA and peat HA. The findings of the present studyimply that, in addition to electrostatic interactions, van der Waals interactions as facilitated by aromatic structures and H-bond donatingmoieties in terrestrial HS may facilitate a favorable binding environment. Environ Toxicol Chem 2013;32:1467–1478.# 2013 SETAC

Keywords: Antibiotic Soil Adsorption

INTRODUCTION

Antibiotics are introduced into soils through land applicationof manures containing antibiotics [1–3] and into surface watersvia direct discharge of contaminated waters [2,4–6]. Theantibiotics may then exert ecotoxicological effects on sensitiveorganisms [5,7,8], including aquatic photosynthetic spe-cies [9,10] and nitrifying and sulfate-reducing soil bacte-ria [11,12]. Fluoroquinolone (FQ) antibiotics, which areinhibitors of DNA synthesis and replication, are an importantclass of broad-spectrum antibacterial agents widely used in bothhuman and veterinary medicine and often detected in surfacewaters [6]. The strong affinity of antibiotics including FQs fordivalent metal cations and organic fractions in soils can influencetheir bioavailability and toxicity [13–15]. Retention of FQswithin organic matter particles has been implicated in both theirpersistence in organic-rich horizons of soils [3] and theirdecreased photodegradation in natural waters [16–18]. Ofimportance in this respect are the mechanisms of adsorptionof FQs and related antibiotics onto organic matter. Elucidation ofthese mechanisms is essential to environmental risk assessment.

Experimental adsorption studies of the interactions of othertypes of antibiotics (tetracycline [19], clarithromycin [20],sulfathiazole [21]) with a soil humic acid (HA) stress theimportance of electrostatic interactions, whereby a decrease insorption as a function of pH correlates with a decrease in theconcentration of cationic antibiotic species. Moreover, involve-

ment of H-bonding interactions and the formation of ternarycation-bridging complexes between tetracycline and soil HA [19]are implied, respectively, by a decrease in adsorption as the H-bond forming carboxylic OH groups of HA deprotonate and anincrease in adsorption in the presence of Ca. Another study [22]reports an increase in the adsorption of oxytetracycline onAldrichHA amendedwith Al and Fe(III), corroborating the formation of acation-bridging complex, but a decrease in the adsorption in thepresence of Ca, suggesting competition from Ca.

A report by Vasudevan et al. [23] on the adsorption ofciprofloxacin (Cipro; Figure 1), a widely prescribed FQantibiotic commonly detected in rivers, on 30 different naturalsoils, demonstrated 2 different types of profile for Ciproadsorption as a function of pH—a decrease in adsorption as pHincreases and an initial increase in adsorption up to pH 5.5followed by a decrease. These trends were attributed mainly todifferences in the cation-exchange capacity (CEC) of the soils,with an additional contribution of surface complexation in metaloxide–rich soils. In an effort to probe Cipro binding to theorganic fraction of soils, Carmosini and Lee [24] investigatedCipro adsorption on reference terrestrial HAs and fulvic acids(FAs) and found that the HAs exhibited up to 3 times moreadsorption capacity for Cipro than did the FAs, thus indicatingthat different humic substance (HS) types may present differentbinding environments for the antibiotic. An increase in Ciproadsorption from pH 4 to pH 6, analogous to the findings byVasudevan et al. [23], was thought to be a result of an increase inthe CEC of the HS due to deprotonation of its functional groups,whereas a decrease in adsorption at higher pH is attributed torepulsion between negatively charged HS groups and zwitter-ionic or anionic Cipro species. Furthermore, Carmosini andLee [24] showed a decrease in Cipro adsorption as a function of

All Supplemental Data may be found in the online version of this article.* Address correspondence to [email protected] online 1 March 2013 in Wiley Online Library

(wileyonlinelibrary.com).DOI: 10.1002/etc.2214

Environmental Toxicology and Chemistry, Vol. 32, No. 7, pp. 1467–1478, 2013# 2013 SETAC

Printed in the USA

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increasing ionic strength under acidic pH conditions; at pH 7 andhigher, however, an increase in ionic strength did not affectCipro adsorption, suggesting that cations may compete less withthe HS binding sites for Cipro due to the greater CEC of the HS athigher pH. For other FQ compounds (flumequine and oxilinicacid) which contain a carboxyl group but no positively chargedamino group, Lützoft et al. [25] reported an increase in theiradsorption on soil HA as a function of increasing pH up to pH 5,but this pH-dependent behavior is posited to be a result of thecomplexation of the increasingly ionized carboxylic acidmoiety of the FQs by protonated amino groups in the HA.However, in the same study [25], adsorption of sarafloxacin, anFQ compound with a positively charged amino group, exhibitedthe same trend up to pH 5 but the amount adsorbed was about 10times more, indicating that cation-exchange reactions dominatedat acidic pH values. At higher pH, sarafloxacin adsorptiondecreased and approached the values obtained for the adsorptionof flumequine and oxolinic acid, thus implying that theadsorption may be predominantly mediated by the carboxylicmoiety of the FQ compounds following the deprotonation of theamino moiety. Given the high dissociation constant (pKa) of theprotonated amino group in Cipro (Figure 1) compared withsarafloxacin, it is reasonable to posit that, under environmentallyrelevant pH conditions, Cipro adsorption by organic fractions ofsoil will be primarily dominated by cation-exchange reactions,mediated by electrostatic attraction between the positivelycharged amino group of the antibiotic and the negatively chargedHS groups, in agreement with the proposed mechanism put forthby Carmosini and Lee [24] for Cipro, by Gu et al. [19] for atetracycline antibiotic, by Sibley and Pedersen [20] for amacrolide antibiotic, and by Richter et al. [21] for a sulfonamideantibiotic. However, other interactions, such as H-bonding andcation-bridging interactions, which were suggested previouslyby Gu et al. [19] and MacKay and Canterbury [22] fortetracycline antibiotics, may also be important in the interactionof FQs with HS particles, especially when cation exchange isimpeded by the ubiquitous competing presence of cations innatural soils and waters.

In a recent computational study on the favorable interactionsin the complexes of Cipro with a model HS [26], wedemonstrated that both the hydrophobic and hydrophilic HSmoieties undergo structural rearrangements to facilitate stableCipro–HS complexes via the formation of multiple H-bondsbetween HS OHmoieties and Cipro keto and carboxylate groupsand between HS keto groups and the Cipro amino group. Ourmolecular modeling results also indicated that divalent cationsbound to HS carboxylates can impede binding of the positivelycharged Cipro amino group and mediate outer-sphere ternarycomplexes, whereby awater molecule serves as a bridge betweenthe Cipro carboxylate and the HS-bound cation. It is important tonote that HS-bound Mg and Fe(II) mediate this ternary complexbut not HS-bound Ca [26], indicating that the high ionic potential(radius/valence) of the HS-bound cation facilitates the bridginginteraction, in agreement with the aforementioned increase inoxytetracycline adsorption on Aldrich HA amended withtrivalent cations [22]. In a review on the challenges underlyingpredictive models of the environmental fate of polyfunctionalionic compounds, MacKay and Vasudevan [27] pointed out theneed to consider the potential multiple interactions of thesecompounds with environmental surfaces: hydrophobic partition-ing, electron donor–acceptor interactions, cation exchange, andcation bridging.

In sum, previous experimental and molecular modelingstudies lead us to propose the following: 1) in addition toelectrostatic interactions between the Cipro amino group and HScarboxylates, van derWaals andmultiple H-bonding interactionsas facilitated by optimal interactions with hydrophobic andhydrophilic moieties in soil organic fractions may contribute tothe high affinity of FQ antibiotics, and 2) multivalent cations caneither increase or decrease this affinity. An experimentalinvestigation that considers all of these hypotheses with differentHS adsorbents and varying pH can thus provide insight into thefate of FQs in natural settings having differing aqueouschemistry and adsorbent chemistry.

Here, we report the adsorption of Cipro on 4 reference HSsisolated from both terrestrial and aquatic sources: Elliot soil HA,Pahokee peat HA, and Suwannee River HA and FA. Weemployed fluorescence quenching spectroscopy to monitorCipro adsorption on the HS samples equilibrated in a syntheticfreshwater under varying pH and alkalinity. We interpret theadsorption results in terms of the chemical properties of each HSand the ampholytic properties of Cipro.We alsomonitored Ciprocomplexation under differing concentrations of Ca and Mg toexamine the influence of divalent cations, typically presentwithin soil particles, on Cipro–HS interactions. Our studyappears to be the first to employ explicitly the acid-basechemistry of HSs from both terrestrial and aquatic sources todescribe experimental adsorption data of a cationic organiccontaminant, thus offering a comparative insight essential in theassessment of the retention of FQs and other related antibioticsby organic fractions of differing chemistry over a broad range ofenvironmental conditions.

MATERIALS AND METHODS

Materials

Ciprofloxacin hydrochloride was obtained from MP Bio-medicals. Standard Elliott soil HA (1S102H), Pahokee peat HA(1S103H), Suwannee River HA (2S101H), and Suwannee RiverFA (2S101F)—henceforth termed soil HA, peat HA, aquaticHA, and aquatic FA, respectively—were obtained from theInternational Humic Substances Society and used in base

pH4 5 6 7 8

Frac

tion

of to

tal s

peci

es

0.0

0.2

0.4

0.6

0.8

1.0Cipro+ Cipro+/-

Cipro-

pKa1 = 6.1 ± 0.2

pKa2 = 9.2 ± 0.5

Figure 1. Structure and pH speciation of Cipro. Cipro ¼ ciprofloxacin;pKa ¼ acid dissociation constant.

1468 Environ Toxicol Chem 32, 2013 L. Aristilde and G. Sposito

titration experiments without further treatment. The chemicalcomposition of the samples is given in Table 1. All otherchemicals were analytical grade, purchased from FisherScientific. Absorbance measurements were made with aShimadzu UV-160 spectrophotometer and fluorescence meas-urements, with a Fluorolog-3 Model FL3-11 fluorescencespectrometer.

Base titrations

To prepare for conducting a base titration of HS, a 100-mLsolution containing 360 mg/L sample dissolved in 0.05 MNaClbackground electrolyte was stirred moderately in a 125-mLErlenmeyer flask in the absence of light for 24 h at laboratorytemperature. As noted by Ritchie and Perdue [28], pretreatmentwith 0.002 M NaOH was necessary to obtain completedissolution of the soil HA and peat HA samples—an equivalentamount of HCl was subsequently added to neutralize the addedNaOH. The sample solution was poured through a 0.2-mm filterinto a 150-mL glass beaker, which was then securely capped,allowing access only for a pH electrode, an N2 bubbler, and a20-mL pipette. To minimize dissolved CO2 species, the solutionwas purged with N2 for approximately 1 h before andcontinuously during a base titration. Titrations were conductedunder low-light conditions to prevent photo-oxidation sidereactions. With the reaction vessel submerged in a water bath(24.9 # 0.2 8C), base titrations were performed under continualmagnetic stirring by adding 0.02-mL increments of 0.1 MNaOH, raising the solution pH gradually from its initial value(3.20–3.51) to approximately 8.0. The titration pH values wererecorded continuously (with a precision of 0.01 pH unit) at 2- to5-min intervals using a Beckman pH meter equipped with anOrion-Ross combination electrode that was previously calibrat-ed using Fisher standard buffer solutions. The electrode was leftin the sample solution throughout the course of a titration.

Blank solutions were performed in the same manner as thetitrations of the sample solutions. A small volume of 0.1 M HClsolution was first added to the blank to bring its pH value to thatmeasured initially for each sample solution. For both the blankand the sample solutions, the dilution-corrected concentration ofbase added was calculated at each titration step. Thus, the sample

titration data were blank-corrected at each titration step bysubtracting the amount of base required to bring the blanksolution to a chosen pH value from the amount of base added tothe sample solution to achieve the same pH value. Blankcorrection was facilitated by fitting the blank titration data withpower-law regression curves describing the experimentallyobserved relationship between the concentration of base addedand the resulting pH value.

Base titration experiments measure the changes in net protoncharge of a sample (dsH,titr) caused by the addition of strongbase, given by the conventional expression [29]

csdsH;titr ¼ ½OH%& % ½Hþ& % ½NaOH&added ð1Þ

where cs is the sample concentration. The true net proton charge(sH) is the sum of this measured change and the initial net protoncharge (sHo) before any titrant is added [30]:

sH ¼ sHo þ dsH;titr ð2Þ

Although it is commonly assumed to be equal to 0 [31,32], infact sHo must be determined independently of the titrationexperiment because it represents the net proton charge due toacidic functional groups that have reacted during equilibration ofthe sample prior to the titration experiment. In their comprehen-sive analysis of titration experiments involving clay minerals,Bourg et al. [31] obtained the value of sHo through explicitchemical modeling of sample equilibration. Following theirapproach, we note that the initial blank 0.05 M NaCl solution,after 1 h of N2 bubbling, was at pH 7.1 # 0.2 and that, followingequilibration of a sample solution, the initial pH value was 3.20to 3.51. Given the significant drop in pH caused by introductionof the HS sample, we conclude that sHo reflects protons desorbedby exchange with Naþ in the background electrolyte solutionduring sample equilibration and accordingly that it can be setequal to minus the measured initial concentration of protons,[Hþ]o, divided by the sample concentration:

sHo ¼ %½Hþ&o=cs ð3Þ

Table 1. Chemical composition of the humic substances

Soil HA Peat HA Aquatic HA Aquatic FA

Element a (%)C 58.13 56.37 52.63 52.34H 3.68 3.82 4.28 4.36O 34.08 37.34 42.04 42.98N 4.14 3.69 1.17 0.67S 0.44 0.71 0.54 0.46P 0.24 0.03 0.013 0.004

Carboxylb (mol kg%1 C) 5.85 # 0.52 7.27 # 0.71 5.89 # 0.57 9.36 # 0.76Phenolic OH moietiesc (mol kg%1 C) 1.87 1.91 1.72 2.84Other OH moietiesd (mol kg%1 C) 1.90 0.93 0.92 0.40NH moietiesd (mol kg%1 C) 4.39 2.15 0.55 0.15Carbonyla 0.06 0.05 0.06 0.07Aliphaticitya 0.16 0.19 0.29 0.35Aromaticitya 0.50 0.47 0.37 0.24

aElemental composition and fraction of C content found in carbonyl, aliphatic, and aromatic structures as reported by the International Humic Substances Society(http://humicsubstance.org).bMeasured by a modified calcium acetate method [26].cDetermined by Ritchie and Perdue [28] based on analysis of titration data as reported by the International Humic Substances Society.dBased on the amount of OH (excluding carboxyl and phenolic OH) and NH moieties from the content in amino acids and carbohydrates as reported by theInternational Humic Substances Society.HA ¼ humic acid; FA ¼ fulvic acid.

Adsorption of ciprofloxacin on humic substances Environ Toxicol Chem 32, 2013 1469

http://humicsubstance.org/

This estimate of sHo is reasonable, given the highconcentration of Naþ in the background electrolyte solutionand the well-known strong acidity of HS.

HS chemical modeling

The HS model put forth by Westall et al. [33], wherein HSacid-base chemistry is described by 4 proton affinity classes withpredesignated pKa, was applied recently to describe theadsorption edges for cationic antibiotics by a soil HA [19,20].Noting the parameter proliferation issue raised by Westallet al. [33] and expanded on in the context of the NICA-Donnanmodel by Lenoir and Manceau [34], we have developed asimplified version of the Westall et al. [33] model with just 3proton affinity classes to describe the acid-base chemistry of HS.The 3 predesignated pKa values in our model reflect protondissociation constants inferred from recent comprehensivestudies of the acid-base chemistry of HAs and FAs [28,34–38]. Importantly, our model is not optimized on publishedtitration data, which often contain unreported systematic errors(e.g., no blank correction and improper renormalization to theinitial value of the net proton charge on the HS sample beingtitrated) [38]. Instead, we use new blank-corrected base titrationdata that have been properly renormalized, withmodel validationachieved through comparing the model-optimized relativecontent of each affinity class with independent quantitation ofthe principal acidic functional groups in the HS. We apply oursimplified model to account for the pH-dependent speciation ofHS in the formation of Cipro–HS complexes under a range ofenvironmentally relevant pH conditions (pH 4–8).

We now follow the approach ofWestall et al. [33] in applyingthe optimization program, FITEQL [29], while employing thefollowing chemical-equilibrium (Equations 4–6) and mass-balance (Equations 7 and 8) constraints, in addition to a totalproton balance condition (TH, see below):

HLi ¼ Li% þ Hþ Ki i ¼ 1% 3 ð4Þ

Naþ þ Li% ¼ NaLi KNa i ¼ 1% 3 ð5Þ

H2O ¼ Hþ þ OH% Kw ð6Þ

½Naþ&total ¼ ½Naþ& þX

½NaLi& i ¼ 1% 3 ð7Þ

½HLi&total ¼ ½HLi& þX

Li%½ & þX

NaLi½ & i ¼ 1% 3 ð8Þ

with HLi representing a protonated model HS acidic functionalgroup whereby HL1, HL2, and HL3 have assigned respectiveapparent pKa values of 3, 6, and 8, and the summations are overall i species. The equilibrium constant for complexes with Naþ,KNa, is assumed to be the same for all 3 functional groups [33].The choice of pKa values was made in the spirit of the Westallet al. [33] model, such that they would adequately represent theacidic functional groups commonly present in HS, primarilycarboxyl (pKa < 7) and phenolic (pKa ! 8) [28,29,36–38]. Weadded a pKa value lying between 3 and 8 to ensure that thetitration data would be captured accurately by our simplifiedapproach, with the validation of this choice performed bycomparing the sum of [HLi]total for i ¼ 1,2 obtained fromoptimization with the measured carboxyl content of the 4 HSs(Table 1).

Based on their analysis of a large number of titration data sets,Milne et al. [39] reported generic NICA-Donnan model pKa

values of 3.09 # 0.51 (2.65 # 0.43) and 7.98 # 0.96(8.60 # 1.06) for the 2 classes of HA (FA) acidic functionalgroups they ascribed nominally to carboxyl and phenolic OH,respectively. Matynia et al. [37] compiled similar values of these4 intensity parameters from a variety of published modelingapproaches, while offering the opinion that the values below 7deduced by Milne et al. [39] seemed “low for carboxylic typegroups.” They then applied HS mixture models to 10 sets of“high-quality” proton titration data taken from Milne et al. [39],finding average pKa values of 3.77 # 0.20 and 9.85 # 0.33 forcarboxyl and phenolic OH groups, respectively. Noting that theirresults led to NICA-Donnan fits comparable to those achievedby Milne et al. [39], they concluded with Westall et al. [33]that the “data window” of titration experiments is narrower thanthe “spectral window” of parametric modeling, thus necessitat-ing that some parameters be estimated independently ofoptimization.

The optimization procedure in FITEQL [29] minimizes thedifference between the experimentally determined TH (TH

exp)and a model TH (TH

model) based on Equations 4 through 8. Sincedeprotonation of HS functional groups and the dissociation ofwater are the only contributors to the model total proton balance,TH

model is computed according to Equation 9 [33]

THmodel ¼ ½Hþ& % ½OH%& %

XLi%½ & %

XNaLi½ & ð9Þ

The model is optimized on the data by varying the 4 adjustableparameters, [HLi]total and i ¼ 1,2,3, as well as KNa, constrainedby the mass-balance equations (Equations 7 and 8). Takentogether, the 2 summation components in Equation 9 are equal tothe negative net proton charge

THmodel ¼ ½Hþ& % ½OH%& þ sHcs ð10Þ

In the context of a titration, sH is given by Equation 2,supplemented by Equation 3. It follows that the equation forTH

exp as determined by a base titration is obtained by substitutingthese 2 expressions into Equation 10 and referring to thedefinition in Equation 1

THexp ¼ %½Hþ&o % ½NaOH&added ð11Þ

At the start of the experiment, THoexp is determined from the pH

value of the sample solution prior to base addition and a protonacitivity coefficient given by the Davies equation [33,40].

As discussed in the Introduction, the availability of bothionized functional groups and H-bond donor moieties in the HSparticles is essential in the complexation of Cipro. We employedthe HS chemical models developed in the present study tocompute the amount of deprotonated ligands and H-bond donormoieties in each HS as a function of pH. The first quantity wascalculated by taking into account the ionization of the model HSacidic functional groups, HLi (i ¼ 1–3), as a function of pH. Thesecond was computed by considering, in addition to the OH fromthese model HS groups, the reported amount of OH and NHmoieties in carbohydrates and amino acids (Table 2)—theselatter moieties would not deprotonate under the experimental pHvalues due to their high pKa values. To quantify the H-bonddonating capabilities of each HS based on its total content of NHand OH groups, we normalize to the strength of an HSOH or NHmoiety as an H-bond donor to a Cipro carbonyl O atom using thefollowing estimated intramolecular H-bond energies, %24 kJmol%1 for OH…O between alcohols and carboxylates [41] and6 kJ mol%1 for NH…O between amino acids [42].


Adsorption experiments

Adsorption batch reactions were conducted in 250-mL glassErlenmeyer flasks using a 200 # 2-mL suspension of each HSsample at 10 mg/L. The flasks were kept in the dark by beingcovered with aluminum foil to prevent light-mediated sidereactions. The background solution was a synthetic freshwatersolution typically used in aquatic toxicological testing [43–46]:0.520 mM CaCl2, 0.205 mM MgSO4, 0.0147 mM KCl, and0.155 mM NaHCO3. Following pH adjustments, pH 4 to pH 8(0.5 or 1.0 increments), the HS sample was allowed toequilibrate in the freshwater solution for 12 to 18 h beforeintroducing Cipro. Acetate/phosphate (NaC2H3O2/Na2HPO4)buffer at 10 mM was used to maintain constant pH, which wasmeasured using a Beckman pH meter equipped with an Orion-Ross combination electrode previously calibrated using Fisherstandard buffer solutions. Experiments were initiated by addingto each flask an aliquot (100 mL) from concentrated solution ofCipro, initially dissolved in 0.02 M NaOH, to attain an initialCipro concentration of 0 mM to 15.1 mM. For each adsorptionisotherm performed at pH 4 to pH 8 (with1.0 pH increments), atotal of 30 to 46 experiments were carried out using 11 to 14different initial Cipro concentrations (0.0, 0.15, 0.30, 0.76, 1.0,1.5, 3.0, 4.5, 5.3, 6.0, 9.1, 10.6, 12.1, or 15.1 mM Cipro), witheach experiment performed in 2 to 4 replicates. To determine theeffect of divalent cations on Cipro (3.0 mM) adsorption onto HS(10 mg/L), additional adsorption experiments were conductedwith soil HA and aquatic HA equilibrated at various pHvalues, pH 4 to pH 8 at 1.0 increments, with a backgroundsolution similar to above but with the total charge concentrationof Ca2þ and Mg2þ increased by a factor of 2 (from 1.94 to3.88 meq/L)—the solutions were determined to be undersatu-rated with respect to calcite (CaCO3(s)) [47]. Preliminary kineticsexperiments indicated that Cipro adsorption reached equilibriumafter 6 h [47].

Fluorescence quenching

The amount of Cipro adsorbed in these adsorption experi-ments was quantified using fluorescence quenching spectrosco-py after 48 h of reaction time. An appropriate and robust methodto quantify the complexation by dissolvedHS should not involvea physical separation of the adsorptive from the adsorbent toavoid disturbing the adsorption equilibrium and quantitativeuncertainties associated with separation [46]. The ability of HSto quench fluorescing organic contaminants on bindinginteractions has been used previously to monitor HS complexa-

tion of aromatic organic compounds [46,48]. To ensure that thismethod is appropriate for our experiments, we performed adetailed examination of the fluorescence properties of HS andCipro as a function of pH and tested the underlying assumedrelationship [49] between the formation of a Cipro–HS complexand loss of Cipro fluorescence intensity. Our examination alsowill be instrumental in the use of the fluorescence-quenchingmethod to monitor HS complexation with related ionizableorganic contaminants.

It is assumed that quenching is due to neither loss of Ciprofluorescence over time nor inner-filtering effects but rather is aresult of true molecular contact between Cipro and HS.However, these assumptions need to be validated.We performeda series of evaluation and validation steps [47] whereby, at theexperimental pH values mentioned above, both absorbance andfluorescence measurements of Cipro (3.0 mM) and HS (2 mg/L)were recorded from aliquots in a 1-cm path length quartzcuvette. Total luminescence spectra were obtained as a series ofemission spectra, at 0.1-nm increments in emission wavelengthsfrom 400 nm to 500 nm, each recorded at a single excitationwavelength from 250 nm to 400 nm in 1-nm increments; the slitwidth was set to 2.5 nm, and the scan speed was 250 nm/min.Each absorbance or fluorescence spectrum was blank-correctedfor the background solution in the case of Cipro alone or HSalone and for the HS solution in the case of Cipro–HS mixtures.

Examination of the absorbance spectra and the totalluminescence spectra plots revealed no overlapping peaks inabsorbance and 1 excitation–emission wavelength pair from theCipro total luminescence spectra plots, distinct from that of HS,which was therefore used to monitor the quenching of Ciprofluorescence during the adsorption experiments [47] (Table 2).No appreciable change in the fluorescence intensity of Cipro andHS over 100 h [47] indicated that quenching was not affected byloss of fluorescence over time. Also, the absorbance spectra ofCipro–HS solutions represented the sum of the spectra of CiproandHS alone, indicating that the excited states of the fluorophore(Cipro) are not affected by the presence of the quencher (HS);and quenching will be thus a result of changes in the vibrationalstates of Cipro due to molecular contact with HS, as would becharacterized by a decrease in the emission intensities of excitedCipro molecules [49].

A 1:1 molecular contact between adsorbate and adsorbent is aprerequisite for the use of fluorescence quenching spectroscopyin monitoring HS-contaminant binding interactions [49]. Weevaluated the fluorescence quenching of Cipro (3.0 mM) withincreasing concentrations of soil HA or aquatic HA (1–50 mg/L)

Table 2. Excitation–emission wavelength pairs corresponding to peaks in total luminescence plots for Cipro and humic substances

Cipro pH 4 pH 5 pH 6 pH 7 pH 8

lex,1 282 282 280 276 276lem,1 450 448 442 420 420lex,2 335 n.d. 332 n.d. 334

lem,2 450 n.d. 431 n.d. 418


pH 4 lex 347–389 358–385 338–353 332–350lem 465–487 463–495 451–468 440–446

pH 8 lex 363–385 363–400 349–37 330–353lem 463–493 466–500 457–474 444–469

Cipro ¼ ciprofloxacin; HA ¼ humic acid; FA ¼ fulvic acid; n.d. ¼ not determined.


at different pH values (Figure 2). A linear relationship in theStern-Volmer plots, which are plots of HS concentration versusextent of quenching (i.e., ratio of Cipro fluorescence intensity inthe presence of the HS sample to the intensity in the absence ofthe adsorbent), corroborates the presence of a single type offluorophore equally accessible to the quencher [49] and thussuggests that the quenching of Cipro fluorescence was due tospecific interactions with HS. However, deviation from linearityat higher adsorbent concentration (>25 mg/L) may mean thatthe fluorescence quenching method is not appropriate at thesehigh HS concentrations.

We note that nonradiative fluorescence quenching between 2interacting molecules results from either collision (dynamicquenching) or formation of a nonfluorescent complex (staticquenching) between them, and both of these quenchingprocesses can yield a linear relationship for the plot mentionedabove [49]. A relationship between the slope of the Stern-Volmerplots, termed the Stern-Volmer constant (KSV), and thebimolecular quenching rate constant, KQ [50,51], was employedin the present study to gain insight into which type of quenchinginteraction is likely between Cipro and HS. The value of KQ isestimated by dividing KSV by the average lifetime of thefluorophore in the absence of the quencher, t0 [49]. Given anupper limit of KQ on the order of 1010 M%1s%1 for a diffusion-controlled bimolecular process, a greater value is indicative thata binding interaction was likely the cause of the fluorescencequenching [49]. Using t0 ¼ 2.3 * 10%9 s for Cipro [52] andpublished relative molecular masses for the HA samples (15 850daltons for soil HA [53] and 3790 daltons for aquatic HA [54]),

the calculatedKQ values exceeded the diffusion-controlled upperlimit by 2 to 3 orders of magnitude and thus suggested stronglythe occurrence of a Cipro–HS complex. The relative molecularmasses for the HAs are only estimates and can vary considerablydepending on the analytical methods [55] used to determinethem, but our conclusion that the quenching efficiency is notmerely due to diffusive collision still stands after introducing alarge uncertainty (up to 3 orders of magnitude) in the estimatedHS molecular masses.

Taken collectively, these encouraging results indicatethat the fluorescence quenching method is adequate formonitoring Cipro–HS interactions in our adsorption experi-ments. Therefore, the loss of Cipro fluorescence intensitymeasured at a specific excitation–emission wavelength pairdependent on experimental pH (see Table 2) is attributed to adecrease in the concentration of free (or unbound) Cipro insolution. The amount of adsorbed Cipro was then calculated asfollows

Q ¼½Cipro&0 % ½Cipro&eq

CHS¼½ &mmol

gð12Þ

where Q is in micromoles of Cipro per gram of HS, [Cipro]0 isthe total initial concentration of Cipro in micrometers, [Cipro]eqis the equilibrium Cipro concentration after 48 h of reaction,and CHS is the total mass (grams) of dry HS adsorbent perliter of suspension. Due to the background fluorescence intensityof HS in solution (at 10 mg/L), the quantitation limit for Ciprowas 0.06 mM, 0.05 mM, 0.06 mM, and 0.13 mM in the presenceof, respectively, soil HA, peat HA, aquatic HA, and aquaticFA [47].

Modeling complexation by HS

Gu et al. [19] applied a spectrum affinity model, following themethod of Westall et al. [33], to account for both pH and ionicstrength effects (NaCl background electrolyte) on the adsorptionof tetracycline by a soil HA, and Sibley and Pedersen [20] havedone the same for the adsorption of clarithromycin by the sameHS. In both of these applications, adjustable equilibriumconstants for antibiotic adsorption by 2 of the 4 model acidicfunctional group ligands in the HS sample were sufficient todescribe all available experimental data accurately. Using asimilar approach to model the complexation of Cipro by eachHS, we considered the complexation of Naþ (to account for ionicstrength effects) and the positively charged moiety (i.e., theprotonated secondary amino group) of the cationic andzwitterionic species of Cipro (respectively, Ciproþ and Cipro0)by the deprotonated HS ligands of our simplified 3-protonaffinity HS model (i.e., HL1, HL2, or HL3)—the HS parametersare listed in Table 3. The modeling setup for the component-species matrix for the Cipro–HS complexation reactions isdetailed elsewhere [47]. Briefly, the modeling exercise includedthe equilibria describing the acid-base chemistry of the HSligands (HL1, HL2, HL3), Cipro

þ, and Cipro0; and the complexesNaLi (I ¼ 1–3), Ciproþ-L1, Ciproþ-L2, Cipro0-L1, andCipro0-L2. The binding affinity of each antibiotic species wastaken to be the same for L1% and L2% because, as mentioned inHSchemical modeling, these 2 model functional groups collectivelyrepresent the carboxyl content. The Cipro–HS complexationmodels were optimized on the adsorption data obtained from thereaction of 3.0 mM Cipro with 10 mg/L of soil HA, peat HA,aquatic HA, and aquatic FA. The adsorption experiments wereperformed as previously described at pH 4 to pH 8 (at 0.5increments)

A

Wavelength, nm360 420 480 540

Inte

nsity

C

Caquatic HA, mg/L5 20 35 50

F/F o

1

3

5

7

B

Wavelength, nm360 420 480 540

Inte

nsity

D

Csoil HA, mg/L5 20 35 50

F/F o

1

8

15

22

Figure 2. Fluorescence quenching of ciprofloxacin (Cipro) by humicsubstance (HS) and associated Stern-Volmer plots. Representative spectraof Cipro fluorescence quenching in the presence of (A) aquatic humic acid(HA; from top to bottom 1 mg/L, 5 mg/L, 10 mg/L, 12.5 mg/L, 20 mg/L,25 mg/L) and (B) soil HA (from top to bottom 5 mg/L, 10 mg/L, 15 mg/L,20 mg/L, 25 mg/L, 30 mg/L) at pH 4. Plots of F/Fo versus CHS for Ciproreactionwith (C) aquatic HA and (D) soil HA at pH 4 (black symbols) and pH8 (white symbols), with F and Fo representing the fluorescence intensity,following blank corrections (see Fluorescence quenching for details),of Cipro in the presence and in the absence of HS, respectively.Caquatic HA ¼ concentration of aquatic HA inmg/L; Csoil HA ¼ concentrationof soil HA at mg/L.


RESULTS AND DISCUSSION

Modeling HS chemistry

We modeled the ionizable groups in each HS sample in aneffort to account explicitly for their pH-dependent speciation.Model fits for each of the HS samples are shown in Figure 3 asplots of pH versus TH, and optimized values of the totalconcentrations (normalized to the C content of the HS) of the3 classes of acidic functional groups together with log KNa arelisted in Table 3. The first 2 rows of results in Table 3 indicatethe overall high quality of our optimization fits in matching theexperimental values of TH. This quality was maintainedthroughout the pH range of the titration experiments, asillustrated in Figure 3. To evaluate parameter sensitivity in themodel fits, we used the peat titration data to examine the changesin [HLi]total (i ¼ 1–3) as a result of varying log KNa by 0.5 andthe changes in both [HLi]total (i ¼ 1–3) and log KNa from eitherincreasing pK3 to 9.0 or decreasing it to 7.5 (Figure S1,Supplemental Data). The shifts in log KNa and pK3 did not affectthe values of [HLi]total (i ¼ 1,2) outside their reportedimprecision in Table 3. Likewise, a change in pK3 from 8.0 to9.0 resulted in no statistically significant change in log KNa.However, a change in pK3 from 8.0 to 7.5 resulted in a 32%increase in log KNa (i.e., a more acidic HL3 results in a higheraffinity of the model functional groups for Naþ), and eitherincreasing log KNa by 0.5 or pK3 by 1.0 led to statisticallysignificant increases in [HL3]total (Figure S1, SupplementalData), suggesting that there is some parameter correlationbetween [HL3]total and log KNa.

The model estimates of the carboxyl content (sum of HL1 andHL2) for soil HA, peat HA, aquatic HA, and aquatic FA were,respectively, 5.50 # 0.37 mol kg%1 C, 5.84 # 0.43 mol kg%1

C, 6.19 # 1.35 mol kg%1 C, and 9.51 # 1.28 mol kg%1 C.These values are not statistically different from the measuredvalues of carboxyl content given in Table 1 (2-tailed unpairedt test, p ¼ 0.05). On the other hand, Gu et al. [19] reported thevery large carboxyl content of 8.5 mol kg%1 C for Elliott soilHA as the sum of their optimized [HLi]total (i ¼ 1,2) withassumed pKa values of 3 and 5. We believe this high value is theresult of optimizing the model of Westall et al. [33] on a titrationcurve for Elliott soil HA that was not blank-corrected. Indeed,our uncorrected titration data for soil HA was congruent withtheir published titration curve [47]. Our estimated concentrationof the third model acidic functional group (HL3, pKa ¼ 8) inTable 3 compares favorably with the phenolic group content of

Table 3. Model parameters for HSs and Cipro–HS complexes


HSa

–THo (mM)–THo

exp 252.26 # 27.43 283.04 # 25.99 364.63 # 54.28 515.05 # 45.32–THo

model 247.25 # 29.50 261.06 # 6.09 286.12 # 43.42 476.67 # 84.76[HLi]total (molc kg

%1 C)[HL1]total, pKa ¼ 3.0 3.14 # 0.36 3.63 # 0.21 4.68 # 0.46 6.82 # 0.66[HL2]total Ka ¼ 6.0 2.36 # 0.07 2.21 # 0.38 1.51 # 1.27 2.69 # 1.10[HL3]total pKa ¼ 8.0 1.54 # 0.42 1.29 # 0.33 2.31 # 0.47 2.39 # 0.07

log KNa 1.93 # 0.06 1.82 # 0.20 2.41 # 0.63 2.49 # 0.10Cipro–HSb

log KCiproþ 5.23 # 0.03 5.24 # 0.02 5.21 # 0.00 4.95 # 0.05log KCipro0 5.16 # 0.04 5.07 # 0.04 5.07 # 0.00 4.73 # 0.03

aModel parameters (mean # standard deviation) determined by optimization performed on 3 replicates of each titration experiment.bBinding affinity for each Cipro–HS species deduced from optimization performed on adsorption data using the HS model parameters.HS ¼ humic substance; Cipro ¼ ciprofloxacin; HA ¼ humic acid; FA ¼ fulvic acid.

Figure 3. Experimental and theoretical total proton balance condition (TH)values for (A) soil humic acid (HA), (B) peat HA, (C) aquatic HA, and (D)aquatic fulvic acid. Points and error bars represent, respectively, the meanand standard deviation of the data obtained from 3 replicates. Shaded areasrepresent the span of TH values from the model fit (mean # standarddeviation) at each point on a titration curve.


the 4 HS samples as estimated by Milne et al. [39] using theNICA-Donnan model (Qmax2,H) and Ritchie and Perdue [28]based on their analysis of titration data (Q'2, Q2).

In terms of the trend in the HS ionizable groups as a functionof pH, we note that aquatic FA has the highest amount ofdeprotonated ligands as a result of a relatively high amount ofacidic groups, notably HL1 (Figure 4A). Soil HA and peat HAexhibit similar trends in the amount of deprotonated ligands as afunction of pH and have about 3 to 4 mol less than aquatic FA,whereas aquatic HA possesses a slightly greater amount thanterrestrial HS (Figure 4A). Based on the amount of ionizablegroups, aquatic FA would seem to have the greatest capacity tocomplex Cipro. However, as noted above, H-bonding inter-actions are also important to the affinity of Cipro for HS. Inaccordance with the proposal [26] that the ability of HS to beengaged in multiple H-bonding interactions will be optimalwhen the ionizable ligands are protonated under acidicconditions, Figure 4B illustrates that indeed all 4 HSs exhibitthe highest H-bond donating capabilities at lower pH values.Also as expected, the H-bond donors decrease as a function ofincreasing pH, primarily due to the ionization of carboxylicgroups, assigned to the HL1 and HL2 ligands (Table 3 andFigure 3B). However, this trend differs among the HS samples.Soil HA exhibits the highest amount of H-bond donating

moieties, implying a higher capacity to be engaged in H-bondinginteractions with Cipro to confer a more favorable bindingenvironment. Peat HA and aquatic HA should have a relativelysimilar decreasing trend in the amount of H-bond donors as afunction of increasing pH.While aquatic FA initially possessed ahigher total amount of H-bond donating moieties than both peatHA and aquatic HA, this amount decreased rapidly at pH >3(Figure 3B) because a significant contribution comes fromcarboxylic groups with very low pKa values, represented in thepresent study as HL1 (Tables 1 and 3 and Figure 3B).

Ciprofloxacin adsorption as a function of pH

According to the respective acid-base speciation of the HSadsorbents and the Cipro adsorptive, we would expect thatelectrostatic attraction between the positively chargedCipro (protonated amino group) and deprotonated HS acidicfunctional groups would predominate at pH < pKa1, whereas athigher pH there would be, in addition to these latter interactions,repulsion between the Cipro carboxylate anion and thenegatively charged moieties of HS and possible cation-bridginginteractions between HS-bound cations and the Cipro carboxyl-ate. Hydrogen-bonding interactions would be favored atlower pH, as previously suggested [26]; and according to thetrend in H-bond donors as a function of pH, discussed inModeling HS chemistry and illustrated in Figure 4B, we positthat the soil HA would provide the most favorable H-bonddonating environment among the 4 HSs used in our adsorptionexperiments.

To examine Cipro adsorption as a function of pH, weobtained adsorption isotherms, shown in Figure 5A, measuredat pH 4 to pH 8 at 1-pH increments for soil HA, peat HA, aquaticHA, and aquatic FA, wherein Q and [Cipro]eq are as previouslydefined. Interestingly, at pH 4, when about 99.5% of the Ciprospecies are cationic (Figure 1), the slopes of the adsorptionisotherms were relatively similar for soil HA, peat HA, andaquatic HA, implying a similar extent of sorption. At higher pH,however, the extent of complexation by each HS differedappreciably (Figure 5A). As pH increased, both soil HA and peatHA exhibited increasingly steeper slopes than the ones obtainedwith aquatic HA and aquatic FA (Figure 5A).

The sorption data were fitted well (r2 ! 0.93) with the vanBemmelen-Freundlich isotherm [37]

Q ¼ Kf ð½Cipro&eqÞb ð13Þ

where Kf is the Freundlich coefficient, with its units being theratio of the units ofQ to the units of [Cipro]eq raised to the powerb, and the exponent b can be interpreted as a heterogeneityparameter, reflecting the spectrum of sorption affinities in the HSsamples [35]. As reported in Table S1 (Supplemental Data), thevalues of bwere near 1.0 (on average 0.9–1.1), indicating a verynarrow affinity spectrum. We took advantage of the (approxi-mately) linear relationship between Q and [Cipro]eq to use Kf,effectively the slope of the lines in Figure 4A, normalized to thecarbon content of each HS (henceforth termed KOC), to comparethe extent of binding of Cipro by each HS as a function of pH(Figure 5B).

The KOC values computed for both soil HA and peat HAfollowed an ascending trend as pH increased to 6, whichcorresponds approximately to the first pKa of Cipro; athigher pH, however, there was either a decrease or no significantchange in KOC. This pH-dependent trend in the Cipro adsorptionon soil HA and peat HA mirrors the results of Carmosini and

A

pH3 4 5 6 7 8 9

Depr

oton

ated

liga

nds,

mol

/L

0

2

4

6

8

10

12

14

B

pH3 4 5 6 7 8 9

H-bo

nd d

onor

s, m

ol O

H Kg

-1 C

0

2

4

6

8

10

12

Figure 4. Computed amount of (A) deprotonated ligands and (B) H-bonddonors in each humic substance (HS) as a function of pH. (A) The amount ofdeprotonated ligands was derived from the model HS parameters in Table 3as the sum of L1% , L2% , and L3% . (B) The amount of H-bond-donatingmoieties was calculated by taking into account both the amount of NH groups(Table 1) and the speciation of the protonated model HS acidic functionalgroup ([HLi], i ¼ 1–3), which represent carboxylic acid and phenolic OHgroups (see HS chemical modeling for more details). Soil humic acid (HA;black solid line), peat HA (gray solid line), aquatic HA (black dashed line),aquatic fulvic acid (gray dashed line).


Lee [24], whereby an increase in adsorption is caused by anincrease in cation-exchange sites on HS due to increasingionization of the HS functional groups. On the other hand, ouradsorption results with the aquatic HS revealed a differentbehavior. With both aquatic HA and aquatic FA, we obtained adecrease or minimal change in the KOC values as a functionof pH, thus implying a more favorable binding environment forCipro with the terrestrial HS than with the aquatic HS.

To examine the importance of HS cation-exchange sitesavailable for Cipro adsorption, we monitored Cipro adsorptionon soil HA and aquatic HA following a doubling of the chargeconcentration of Ca2þ and Mg2þ equilibrated with the HSsamples, from 1.94 meq/L to 3.88 meq/L. We observed astatistically significant decrease only in the amount of Ciproadsorbed on aquatic HA (Figure 6). In accordance with ourmolecular modeling results [25], the present findings demon-strate that divalent cations commonly present in waters and soilsindeed may compete with the positively charged moiety of theFQ antibiotics for binding with the HS ligands, therebydecreasing FQ sorption. However, the extent of this competitionseems to be dependent on the type of HS. As mentioned in theIntroduction, both Gu et al. [19] andMcKay and Canterbury [22]reported enhanced complexation of a tetracycline antibiotic by asoil HA, respectively, in the presence of Ca and of Al and Fe(II).This trend is in accordance with the proposal put forth from ourprevious modeling studies [25,56] that the high ionic potential ofthe HS-bound cation facilitates the formation of a ternarycomplex, outer-sphere bridging complexation of HS-bound

metal cations with the Cipro carboxylate being obtained onlywith Fe(III) and Mg but not with Ca. By contrast to theresults obtained by Gu et al. [19], McKay and Canterbury [22]reported a decrease in tetracycline adsorption when the Caconcentration was increased, results which agree with bothour experimental andmolecular modeling results [25]. However,the lack of appreciable decrease in Cipro adsorption with

pH 4

[Cipro]eq, mol/L0 2 4 6 8

Q,

mol

/g

0

200

400

600

800

1000

1200pH 5


pH 6


pH 7


pH 8


A

Soil HA

pH4 5 6 7 8

K oc, 10

6 L/K

g oc

0.0

0.8

1.6

2.4

3.2

4.0Peat HA

pH4 5 6 7 8

Aquatic HA

pH4 5 6 7 8

Aquatic FA

pH4 5 6 7 8

B

Figure 5. Adsorption isotherms for ciprofloxacin (Cipro) adsorption on soil humic acid (HA;+), peat HA (&), aquatic HA (5), aquatic fulvic acid (FA;,) at pH4, 5, 6, 7, and 8. (A) Amount of adsorbed Cipro (Q, micromoles of Cipro per gram of humic substance sample) as a function of equilibrium Cipro concentration insolution ([Cipro]eq); lines represent van Bemmelen-Freundlich isothermmodel fits (see Equation 1 in text and Table S1 in the Supplemental Data). (B) Organic C-normalized Freundlich coefficient plotted as a function of pH for Cipro adsorption on soil HA, peat HA, aquatic HA, and aquatic FA; error bars represent standarddeviations for 4 to 6 replicates.

A

pH4 5 6 7 8

D oc

(L/K

g oc)

0.0

0.8

1.6

2.4

3.2

4.0B

pH4 5 6 7 8

0.0

0.6

1.2

1.8

2.4

3.0

*

* *

Figure 6. Adsorption envelope for ciprofloxacin on (A) soil humic acid (HA)and (B) aquatic HA. Symbols and error bars represent, respectively, the meanand standard deviation of the experimental data (Doc is a conventionaldistribution coefficient [32] expressed in the present study as 105 L kg%1 C)obtained with sorption experiments performed in synthetic freshwater with1.94 meq/L (black symbols) and 3.88 meq/L (white symbols) of divalentcations (Ca2þ þ Mg2þ). Unpaired 2-tailed t test analysis comparing 2 valuesat the same pH (n ¼ 2–8): - p < 0.01.


soil HA on an increase in alkalinity suggests that the bindingmechanisms of Cipro to soil HA are not affected as muchby the additional amount of divalent cations, which in turnimplies that other mechanisms may be involved in addition toelectrostatic interactions to strengthen Cipro–HS interactionswith terrestrial HA.

According to the proposal put forth by Carmosini andLee [24], an increase in Cipro adsorption as a function ofincreasing pH is expected due to the increasing deprotonation ofthe HS carboxylic acid groups with increasing pH, as illustratedin Figure 4B. However, this trend was observed only with soilHA and peat HA in our experiments, not with aquatic HS(Figure 5B). As discussed in Modeling HS chemistry anddemonstrated in Figure 4, soil HA has the highest amount ofH-bond donor moieties and they may be responsible for its highadsorption capacity. However, peat HA, which exhibits anadsorption capacity similar to soil HA, reflects the same trend inH-bond donating capability as aquatic FA, on which thelowest amount of Cipro was adsorbed. This suggests thatother characteristics in peat HA may be responsible for thisdifference. Interestingly, peat HA also has about the sameamount of C in aromatic structures as does soil HA, but thearomaticity of peat HA is nearly twice that of aquatic FA. Thus,aromatic structures may favor adsorption of Cipro since H-bonddonating moieties of HS decrease as pH increases. In fact,despite the higher amount of the acidic functional groups (i.e.,HL1 and HL2) in aquatic HA and aquatic FA than in soil and peatHA (Table 3), soil and peat HA exhibited significantly higherCipro adsorption than both aquatic HSs as pH increased.Furthermore, we notice that there was less Cipro adsorbed onaquatic FA than aquatic HA at pH 4, despite the greater amountof deprotonated ligands in aquatic FA at pH 4 (Table 3 andFigure 3). It is important to note that a clear difference betweenthe 2 aquatic HSs is that aquatic HA has significantly more of itscarbon content in aromatic structures than aquatic FA. Onceagain, the aromaticity of HS plays a role in stabilizing Cipro–HScomplexes.

The aromatic moieties of FQs have been proposed to mediatep–p interactions with DNA molecules, the toxicological targetsof FQs [57–59]. Interactions between electron-rich or p-donoraromatic moieties of organic contaminants and electron-deficientor p-acceptor quinone rings of HS have been suggestedpreviously [60]. Thus, we hypothesize that similar p–pinteractions may favor Cipro binding by terrestrial HS, aphenomenon perhaps reflected in the structural rearrangement ofthe hydrophobic moieties of a model HS around the Ciprobenzene ring observed in our recent molecular dynamics study ofCipro–HS complexes [25]. These interactions may be moresignificant for terrestrial HAs, whereas electrostatic interactionsmay be more prevalent for the aquatic HS, in accordance with aninverse relationship previously proposed [61,62] between theability of HS to engage in H-bonding interactions and proton-transfer reactions and their ability to engage in donor–acceptorp–p interactions. In agreement with this relationship, we recallthe decrease in Cipro adsorption on aquatic HA but not on soilHA, on increasing the concentration of divalent cations. Takencollectively, our experimental adsorption results imply that,while an increase in cation-exchange sites can lead to higherCipro adsorption as a function of pH, these electrostaticinteractions seem to be stabilized, perhaps even enhanced, byboth H-bond donating moieties and the aromaticity of theadsorbents.

Modeling Cipro–HS complexes

As a preliminary assessment of our HS model, we shallsummarize its success in describing published data on theadsorption of tetracycline [19] and clarithromycin [20], byInternational Humic Substances Society standard Elliott soil HA(Figure S2, Supplemental Data). Using our data on Ciproadsorption by the 4 HS samples, we obtained complexationconstant reactions involving Ciproþ and Cipro0 with L1% andL2% , the most acidic classes of functional groups, which weresufficient to model adequately the adsorption envelopes ofCipro–HS interactions (Figure 7). Both log KCiproþ and log

Figure 7. Model calculations of the adsorption envelope of ciprofloxacin (Cipro) on (A) soil humic acid (HA), (B) peat HA, (C) aquatic HA, and (D) aquatic fulvicacid. Symbols and error bars represent, respectively, the mean and standard deviation of the experimental data (Doc) for 4 to 8 replicates. Shaded areas representthe span of Doc values from the model fits (mean # standard deviation) at each data point (Doc is a conventional distribution coefficient [32] expressed in thepresent study as 105 L kg%1 C). Curves in the figures represent the contribution (average value) of each antibiotic adsorbate species: Ciproþ-L1 (black solid line),Ciproþ-L2 (black dashed line), Cipro0-L1 (gray solid line), Cipro0-L2 (gray dashed line).


KCipro0 were much larger than log KNa (Table 3), suggestingthat the Cipro species mainly compete with Hþ ions for thebinding sites. We note that, while both KCiproþ and KCipro0

implicitly refer to the binding of the positive charge from bothCiproþ and Cipro0 to the same type of HS acidic functionalgroup (i.e., carboxylates), their values differ among the HSsamples.

Accordingly, the speciation of the Cipro–HS complexesdiffered for Cipro interactions with the HS samples (Figure 7). Ingeneral, the most acidic HS ionizable group (HL1, pK1 ¼ 3.0)contributed most to the interaction of Ciproþ with all 4 HSsamples. The contribution of Ciproþ–L2 complexes to theoverall sorption of Ciproþ with the terrestrial HS (soil HAand peat HA) was approximately twice that for the aquaticHS (Table 3 and Figure 7). The aquatic HS had up to 2 timesmore HL1 species than the terrestrial HS (Table 3), whichmay explain the greater contribution of the HL1 species to theaquatic HS.

In the case of Cipro0–HS complexes, both L1% and L2%

contributed significantly to complexation by soil HA andpeat HA, whereas L1% was the main contributor for complexeswith aquatic HS (Figure 7). On the hypothesis that thearomaticity of the HS may be important in stabilizing theformation of Cipro–HS complexes, the KCipro0 bindingconstants (Table 3) were subjected to aromatic C-normalization;but this normalization accounted for the differences in KCipro0

only for the adsorption of Cipro onto soil HA and peat HA,whereas the KCipro0 values for the aquatic HS were stillsignificantly different [47]. This finding corroborated the ideathat the hypothesized p–p interactions may be significantonly for terrestrial HS, whereas electrostatic interactions aremore important for aquatic HS, with little or no contributionfrom aromatic structures.

Environmental chemodynamics

Addressing the parameter proliferation issue raised previ-ously [32,33], we demonstrated that proton titration data for HS,following proper charge renormalization, can be describedquantitatively by a model with only 4 adjustable parameters. It ispatent that natural organic matter plays an important role in thebehavior of cationic contaminants, including antibiotics andherbicides, so our simplifiedmodelmay provide a predictive toolfor assessing the potential hazard of these contaminants tosensitive ecosystems.

Previous studies [3,15,16] have suggested that interactionsof FQs with organic matter fractions of soil may be a controllingfactor in the environmental fate of these compounds: theamount of organic matter in rivers is a major factor indetermining the rate of photodegradation of FQs, and FQspersist in the organic-rich horizons of soils [3]. Our findings alsoindicate that the high Cipro binding capacity with soil HA andpeat HA correlates with the high aromaticity typical of terrestrialHS, facilitating the accumulation of FQs in soils. Bycomparison, the much lower binding capacity of aquatic HSmay imply that significant photodegradation of FQs in rivers ispossible, which raises in turn questions about the need toquantify and assess the fate and effects of photodegradationproducts.

Acknowledgment—The authors express gratitude to the US Department ofAgriculture’s Multistate Research Project W-1082, “Evaluating the Physicaland Biological Availability of Pesticides and Pharmaceuticals in AgriculturalContexts,” for financial support and to A. Yang and E. Joy-Pelen fortechnical support. The authors acknowledge 3 anonymous reviewers for theirinsightful comments.

SUPPLEMENTAL DATA

Table S1.Figures S1 and S2. (225 KB PDF).

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Complexes of the antimicrobial ciprofloxacin with soil