Yu Chen
email: yu.chen@spin.cnr.it

Maria D'antuono

Nicholas B Brookes

Gabriella M De Luca

Emiliano Roberto Di Capua

Di Gennaro

Giacomo Ghiringhelli

Cinthia Piamonteze

Daniele Preziosi

Benoit Jouault

Emiliano Di Gennaro

Mariona Cabero

José María González-Calbet

Carlos León

Jacobo Santamaria

Alessia Sambri

Daniela Stornaiuolo

Marco Salluzzo
email: marco.salluzzo@spin.cnr.it

Ferromagnetic Quasi-Two-Dimensional Electron Gas with Trigonal Crystal Field Splitting

Keywords: ferromagnetic q2DEG, trigonal crystal field, XMCD, EELS, Ti-3d orbitals, (111) STO, LAO/ETO/STO

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Introduction

The discovery of a quasi-two-dimensional electron gas (q2DEG) at the (001) interface between LaAlO3 (LAO) and SrTiO3 (STO) band insulators 1,2 generated a widespread interest due to the remarkable and sometimes unexpected properties, [START_REF] Brinkman | Magnetic Effects at the Interface between Non-Magnetic Oxides[END_REF] alike gate tunable (unconventional) superconductivity, [START_REF] Caviglia | Electric Field Control of the LaAlO3/SrTiO3 Interface Ground State[END_REF][START_REF] Richter | Interface Superconductor with Gap Behaviour like a High-Temperature Superconductor[END_REF][START_REF] Stornaiuolo | Signatures of Unconventional Superconductivity in the LaAlO3 / SrTiO3 Two-Dimensional System[END_REF][START_REF] Ben Shalom | Tuning Spin-Orbit Coupling and Superconductivity at the SrTiO3/LaAlO3 Interface: A Magnetotransport Study[END_REF] arising from the interplay between a rather strong Rashba-like spin-orbit coupling [START_REF] Caviglia | Tunable Rashba Spin-Orbit Interaction at Oxide Interfaces[END_REF][START_REF] Stornaiuolo | Weak Localization and Spin-Orbit Interaction in Side-Gate Field Effect Devices at the LaAlO3/ SrTiO3 Interface[END_REF] and an intrinsic multi-band transport associated to non-degenerate Ti-3d t2g bands.

Moreover, two-dimensional magnetism, coexisting with superconductivity, can be induced by opportune atomic engineering of the heterostructures. [START_REF] Stornaiuolo | Tunable Spin Polarization and Superconductivity in Engineered Oxide Interfaces[END_REF] All these phenomena are related to interfacial orbital reconstructions which dominate the oxide q2DEGs physics. [START_REF] Salluzzo | Orbital Reconstruction and the Two-Dimensional Electron Gas at the LaAlO3/SrTiO3 Interface[END_REF][START_REF] Salluzzo | Origin of Interface Magnetism in BiMnO3/SrTiO3 and LaAlO3/SrTiO3 Heterostructures[END_REF] Recently, (111) LAO/STO interfaces attracted substantial interest due to the reconstruction of the bulk quasi-octahedral crystal field into a trigonal one [START_REF] De Luca | Symmetry Breaking at the (111) Interfaces of SrTiO3 Hosting a Two-Dimensional Electron System[END_REF] and to the possible role of electronic correlations, which need to be further studied to explain their intriguing electronic properties. [START_REF] Monteiro | Band Inversion Driven by Electronic Correlations at the (111) LaAlO3/SrTiO3 Interface[END_REF][START_REF] Rout | Link between the Superconducting Dome and Spin-Orbit Interaction in the (111) LaAlO3/SrTi O3 Interface[END_REF] At the (111) interface, the octahedral (Oh) and quasi-octahedral (D4h) symmetry of bulk STO and of (001) or (110) STO 2DEGs, respectively, turn into a D3d symmetry, where the crystal field is trigonal (Fig. 1a) owing to the hexagonal surface lattice. [START_REF] De Luca | Symmetry Breaking at the (111) Interfaces of SrTiO3 Hosting a Two-Dimensional Electron System[END_REF] Within the D3d symmetry, the t2g orbitals (3dxy, 3dyz, 3dzx) mix and transform into a1g and eg  states (Fig. 1b). The latter can be further split through distortion along the (111) axis and by the spin-orbit coupling. [START_REF] Trama | Strain-Induced Topological Phase Transition at (111) SrTiO3-Based Heterostructures[END_REF] Consequently, the band structure of (111) oxide heterostructures is susceptible to being more significantly altered at the interfaces with respect to the (001) and (110) cases. Indeed, (111) LAO/STO q2DEGs exhibit some intriguing properties, like a reduced superconducting critical temperature and intrinsic anomalous Hall effect related to the topology of the band structure and their Berry curvature. [START_REF] Lesne | Designing Berry Curvature Dipoles and the Quantum Nonlinear Hall Effect at Oxide Interfaces[END_REF][START_REF] Trama | Gate Tunable Anomalous Hall Effect at (111) LaAlO3/SrTiO3 Interface[END_REF] In this manuscript we report on the electronic properties and the magnetism in the artificial spinpolarized q2DEG realized by introducing few unit cells (uc) of EuTiO3, an antiferromagnetic insulator isostructural to STO, between a LAO epitaxial film and a Ti-terminated (111) STO single crystal.

By using x-ray magnetic circular dichroism (XMCD) at the Ti-3d L2,3 and Eu-4f M4,5 edges, complemented by quantum interference device (SQUID) magnetometry, we provide evidence of a ferromagnetic (FM) order originating from the correlation between Ti-and Eu-magnetic moments in this system. Furthermore, we show a sizeable Ti-orbital moment oriented primarily along the in-plane direction that, together with an anisotropic magnetization, suggests an intriguing spin-orbit texture stemming from the interplay between the trigonal crystal field splitting of the low lying a1g and eg  bands, the spin-orbit coupling and the magnetism in this oxide q2DEG.

Experimental Results

LAO/ETO/STO heterostructures were prepared by pulsed laser deposition (PLD) assisted by specular reflection high energy electron diffraction (RHEED). In a background oxygen pressure of 0.8 10 -5 mbar at 720 °C and laser fluence of 1.3 J/cm 2 , 3 uc-thick ETO film and 10 uc-thick LAO film were sequentially grown epitaxially on Ti-terminated (111) STO single crystals preprocessed by buffered hydrofluoric acid etching and oxygen annealing. [START_REF] Torrelles | Electronic and Structural Reconstructions of the Polar (111) SrTiO3 Surface[END_REF] The average number of layers along the [111] direction was monitored by the oscillations of the specular RHEED intensity (see Figure 2a). Each oscillation corresponds to a thickness equal to the distance (d=a/√3, a is the lattice constant) between two Ti/Eu (for ETO, d ~ 0.23 nm) and Al/La (for LAO, d ~ 0.22 nm) layers along [111] direction. Real time RHEED oscillations (Fig. 2a) show that the film growth proceeds layer-by-layer, while Atomic Force Microscopy (AFM, Fig. 2b) confirms that the LAO/ETO/STO surface is composed of step-terraces with a height of ~0.22 nm, similar to the pristine STO (see Supporting Information, SI, Note 1 and SI Figure S1).

High resolution aberration corrected electron microscopy and spectroscopy revealed the high structural quality of the LAO/ETO/STO heterostructures. Fig. 2c displays an annular dark field (ADF) image of a (111) LAO/ETO/STO sample. The layers look homogeneous with no sign of mismatch dislocations, evidencing the epitaxial growth of the heterostructure. The images reveal a sharp interface between ETO and STO, while the LAO/ETO interface is less defined due to similar (atomic number) Z-contrast of La and Eu.

On the other hand, both interfaces are clearly discerned in the right panel of Fig. 2c, which shows false colour La, Eu and Ti elemental maps obtained by electron energy loss spectroscopy (EELS) at Ti L2,3, La M4,5 and Eu M4,5 absorption edges. Composition profiles and EELS show some La/Eu intermixing at the LAO/ETO interface, and a negligible Ti interdiffusion into the LAO film, as shown in SI Figures S2 andS3. The effective average thickness of the ETO layer estimated from the Eu-EELS line scans is about 3-4 Eu rows, thus comparable to the average number of layers as expected from the deposition rate and RHEED analysis (See Fig. 2a). At the LAO/ETO interface, we do see an intermixed unit cell where Eu and La are partially exchanged. Some Eu migrate also in the LAO layer, which can be clearly distinguished from ETO in the line scans by the abrupt reduction of the titanium signal in the EELS maps (SI Figures S2 andS3). The Eu-fraction replacing La, from the EELS line scan, is roughly 20% in the first nm at the interface.

Changes of the local Ti oxidation state across the LAO/ETO/STO interfaces have been investigated by O-K near-edge EELS fine structure (see SI Figure S3). Fig. 2d shows a laterally averaged Ti-valence profile as a function of the spatial coordinate along the interface direction (the error bars in the figure correspond to the standard deviation of the measurement). While in the STO substrate the Ti valence is near to the nominal +4 value, starting from the first two STO unit cells at the interface with ETO it is gradually reduced to a minimum of +3.8 in the first ETO layer at the interface with LAO. This reduction of the Ti oxidation state is related to the partial La/Eu intermixing, which can provide electron doping in the ultrathin ETO layer [START_REF] Katsufuji | Transport and Magnetic Properties of a Ferromagnetic Metal: Eu1-xRxTiO3[END_REF] , and to the presence of a conducting q2DEG.

The LAO/ETO/STO q2DEG exhibits a metallic temperature dependence of the sheet resistance down to 4.2 K with low temperature values in the 0.6-0.8 K range, comparable to standard (001) LAO/STO 6 and (001) LAO/ETO/STO heterostructures. [START_REF] Stornaiuolo | Tunable Spin Polarization and Superconductivity in Engineered Oxide Interfaces[END_REF][START_REF] Di Capua | Orbital Selective Switching of Ferromagnetism in an Oxide Quasi Two-Dimensional Electron Gas[END_REF] The sheet resistance of the q2DEG can be strongly modulated by back-gating using electric field effect (Fig. 3a). At zero gate voltage the resistance shows a small upturn, possibly related to weak-localization effects, like in (001) LAO/STO and LAO/ETO/STO heterostructures [START_REF] Caviglia | Tunable Rashba Spin-Orbit Interaction at Oxide Interfaces[END_REF][START_REF] Stornaiuolo | Weak Localization and Spin-Orbit Interaction in Side-Gate Field Effect Devices at the LaAlO3/ SrTiO3 Interface[END_REF][START_REF] Stornaiuolo | Tunable Spin Polarization and Superconductivity in Engineered Oxide Interfaces[END_REF] and (111) LAO/STO interfaces. [START_REF] Monteiro | Band Inversion Driven by Electronic Correlations at the (111) LaAlO3/SrTiO3 Interface[END_REF] When a positive gate voltage is applied, on the other hand, the resistance shows a metallic temperature dependence with a drop below 8-10 K, possibly related to a FM ordering (see Fig. 3b), similar to (001) LAO/ETO/STO q2DEGs. [START_REF] Stornaiuolo | Tunable Spin Polarization and Superconductivity in Engineered Oxide Interfaces[END_REF][START_REF] Di Capua | Orbital Selective Switching of Ferromagnetism in an Oxide Quasi Two-Dimensional Electron Gas[END_REF] In order to study the crystal field splitting and the magnetism in our heterostructures, we used polarization dependent x-ray absorption spectroscopy (XAS) across the Eu-M4,5 and the Ti-L2,3 edges. The experiments were performed at the beamline X-Treme of the Swiss Light Source [START_REF] Piamonteze | X-Treme Beamline at SLS: X-Ray Magnetic Circular and Linear Dichroism at High Field and Low Temperature[END_REF] and at the beamline ID32 of the European Synchrotron Research Facility (ESRF). [START_REF] Brookes | The Beamline ID32 at the ESRF for Soft X-Ray High Energy Resolution Resonant Inelastic X-Ray Scattering and Polarisation Dependent X-Ray Absorption Spectroscopy[END_REF][START_REF] Kummer | The High-Field Magnet Endstation for X-Ray Magnetic Dichroism Experiments at ESRF Soft X-Ray Beamline ID32[END_REF] XAS performed with circularly or linearly polarized photons can detect the magnetic moments and the 3d orbital crystal field splitting, respectively. The two techniques, usually known as XMCD and x-ray linear dichroism (XLD) are so sensitive that they can be used on a single interface. The samples were mounted with the (111) surface in the vertical laboratory plane, and the vertical axis parallel to the [101 ̅ ] crystallographic direction determined by x-ray diffraction, i.e., parallel to one of the three equivalent Ti-Ti bonds within the hexagonal surface lattice (Fig. 3c). The energy resolution at both beamlines was around 100 meV. XAS, XLD and XMCD spectra were compared to theoretical spectra obtained by atomic multiplet scattering calculations using the CTM4XAS code. [START_REF] Stavitski | The CTM4XAS Program for EELS and XAS Spectral Shape Analysis of Transition Metal L Edges[END_REF] The orbital crystal field symmetry and splitting, of the (111) LAO/ETO/STO q2DEG, were obtained from the analysis of XLD spectra, determined as the difference of XAS spectra acquired with linear polarizations along the horizontal (H) and vertical (V) directions in the laboratory frame, XLD = I H -I V . In Fig. 3d we present a typical XLD spectrum measured in grazing incidence conditions (GI, = degrees from the surface normal). The XLD of (111) heterostructures differ notably from the typical one measured on (001) interfaces (see SI Figure S4). [START_REF] Salluzzo | Orbital Reconstruction and the Two-Dimensional Electron Gas at the LaAlO3/SrTiO3 Interface[END_REF][START_REF] De Luca | Symmetry Breaking at the (111) Interfaces of SrTiO3 Hosting a Two-Dimensional Electron System[END_REF] It can be reproduced by atomic multiplet scattering calculations only by using a trigonal crystal field symmetry (D3d), and a splitting of 20 meV between the (lowest) a1g and eg  orbitals.

Since the total electron yield (TEY) probing depth is between 2-3 nm at the Ti-edge, the results suggest that not only the interface layer, but also the ETO and the first STO unit cells exhibit a trigonal crystal field, in spite of the quasi-Oh symmetry in bulk STO. The trigonal crystal field reflects in the presence of at least two non-degenerate bands originating from the trigonal a1g and eg  orbitals, corresponding to J=1/2 and J=3/2 doublet and quartet respectively.

To further elucidate the effect of the trigonal symmetry on the magnetic properties of this heterostructure, we studied the magnetism of the (111) LAO/ETO/STO q2DEG by using XMCD at the Eu-M4,5 and the Ti-L2,3 edges. In Fig. 4a we show typical large field (6.5 Tesla) XMCD spectra acquired in GI conditions. Alike (001) LAO/ETO/STO q2DEG, [START_REF] Stornaiuolo | Tunable Spin Polarization and Superconductivity in Engineered Oxide Interfaces[END_REF][START_REF] Di Capua | Orbital Selective Switching of Ferromagnetism in an Oxide Quasi Two-Dimensional Electron Gas[END_REF] (111) LAO/ETO/STO shows a strong XMCD at both Eu-M4,5 and the Ti-L2,3 edges. The Eu-spin moment, obtained from the sum rules, is of the order of 4 B/Eu, lower than the 7 B/Eu value expected for Eu 2+ states, suggesting the presence of a fraction of Eu in the non-magnetic Eu 3+ configuration. This result agrees with the EELS data, which show some Eu/La substitution inside the LAO layer, where Eu get an Eu 3+ valence, explaining the reduced value of the XMCD.

The overall shape of the XMCD spectra can be reproduced by atomic multiplet simulations for Eu 2+ and Ti 4+ ions in C3 symmetry. However, while a positive exchange is needed to simulate the Eu 2+ magnetization, parallel to the magnetic field, only a negative one correctly reproduces the Ti-XMCD peaks at L3 and L2, analogous to the (001) interface (see refs. [10, 20]).

Furthermore, we also find a finite and positive Ti-XMCD integral which, according to the sum rules, [START_REF] Piamonteze | Accuracy of the Spin Sum Rule in XMCD for the Transition-Metal L Edges from Manganese to Copper[END_REF] indicates a sizeable Ti-3d orbital moment, morb, up to -0.025 B/Ti. The finite XMCD integral, unexpected for empty Ti-3d 0 orbitals, is related to a small fraction of Ti 3+ (Ti-3d 1 ) electronic states which contribute to the finite magnetic (orbital and spin) moment. The result qualitatively agrees with the Ti valence profile measured by EELS across the LAO/ETO/STO interface, suggesting the presence of Ti-3d 1 electrons within the ETO and in the first unit cells of the STO.

In Fig. 4b we report Ti-XMCD spectra and XMCD-integral at 0.1 Tesla as function of the angle . The data show a sizeable XMCD also at this low magnetic field, and an orbital moment, which progressively increases with the angle  and changes sign at normal incidence (=) . In Figs. 4c,d we show the magnetic field dependence of the Eu-M5 and Ti-L3 XMCD intensity, proportional to the magnetization, measured at 2 K in normal and GI conditions. The data show that the XMCD is anisotropic, with a preferred in-plane orientation, and saturates above 2-3 Tesla. Moreover, the XMCD intensities acquired at the two edges are clearly correlated, since they have a similar magnetic field dependence, suggesting a magnetic coupling between Eu-4f and Ti-3d states.

In order to establish if the system is ferromagnetic, we show in Fig. 4e a magnetic hysteresis loop at 2 K obtained by SQUID magnetometry. The data shows the presence of a hysteresis at low field when the field is parallel to the interface, which suggests FM correlations in the plane of the q2DEG. On the other hand, no hysteresis was observed for a magnetic field perpendicular to the interface.

In Fig. 4f we show the angular dependence of the Eu (𝑚 𝑠𝑝𝑖𝑛 ) and Ti (𝑚 𝑜𝑟𝑏 ) orbital magnetic moments at 0.1 Tesla and 2 K obtained using the sum rules. The data confirm that the Ti-3d orbital moment decreases as a function of the angle , and changes sign in normal incidence, indicating that the Ti and Eu spin moment (according to the Hund's rules) are parallel (and directed along the magnetic field) in GI conditions, but antiparallel in normal incidence. One of the possible Euspin configurations, which can explain the absence of a finite out-of-plane magnetization, is shown in the inset of Fig. 4f. The idea is that Eu-spin moments are canted out-of-plane by the perpendicular magnetic field in alternating up and down sequences, so that no net out-of-plane magnetization is observed at zero field.

These results suggest that the magnetism in (111) LAO/ETO/STO is rather complex and characterized by possibly non-collinear spins, with a FM ordering of the in-plane magnetic moment component, in agreement with the presence of a clear hysteresis only for a field parallel to the interface, and a FM coupling between Ti-3d and Eu-4f moments in the plane. On the other hand, the coupling between Ti and Eu magnetic moments switches to antiferromagnetic in the outof-plane direction, where also no hysteresis is observed in the XMCD and SQUID magnetization loops, suggesting a complex spin-texture which deserves further theoretical and experimental studies.

Conclusions

We have demonstrated the formation of a ferromagnetic q2DEG at the (111) LAO/ETO/STO interface characterized by a reconstruction of the bulk quasi-octahedral crystal field into a trigonal field. Atomic resolved electron energy loss spectroscopy shows that the q2DEG is located in the EuTiO3 layers and extends into the first layers of SrTiO3. Furthermore, we found an intriguing angular dependence of the Ti-orbital moment, suggesting a sizable in-plane moment and in-plane ferromagnetic correlations. The study on the (111) LaAlO3/EuTiO3/SrTiO3 q2DEG reveals an intriguing interplay between ferromagnetism, spin-orbit coupling and trigonal crystal field.

Methods

Samples preparation

LAO(10 uc)/ETO(3 uc)/STO(111) heterostructures were fabricated by PLD assisted by RHEED from sintered Eu2Ti2O7 and crystalline LAO targets onto TiO2-terminated (111) STO substrates pre-processed by buffered hydrofluoric acid etching and oxygen annealing. The samples were deposited using an excimer laser (Lambda Physics, 248 nm wavelength) and 1.3 J/cm 2 fluence, 1 Hz of repetition, in a background O2 pressure of 8×10 -5 mbar at a temperature of 720 °C and cooled down in the same conditions to room temperature with a rate of 5 °C/min.

Electron microscopy

Specimens were prepared by grinding and Ar ion milling. Sample structure was characterized by STEM at Centro Nacional de Microscopia Electronica, Universidad Complutense, Madrid. Zcontrast images were obtained in a NionUltraSTEM operated at 100 kV and equipped with a 5th order corrector and a Gatan Enfina spectrometer capable of simultaneous AFD imaging and EELS at Ti L2,3, La M4,5 and Eu M4,5 absorption edges. Ti L2,3 edge XAS and XLD spectra at 5K, and simulation by atomic multiplet scattering calculations using the CTM4XAS code with a D3d crystal field. [START_REF] Stavitski | The CTM4XAS Program for EELS and XAS Spectral Shape Analysis of Transition Metal L Edges[END_REF] The crystal field parameters are reported in ref. [13]. ) corresponds to the out-of-plane magnetic field orientation. In the inset we show a sketch of a possible non-collinear arrangement of the Eu-spin moments which leads to an in-plane ferromagnetic and an out-of-plane antiferromagnetic alignment of the spin components in the two directions (blue spheres Eu ions, red spheres Ti ions, grey spheres oxygen ions). All the magnetic measurements were performed by zero field cooling the sample down to base temperature.

Figure 1 .

 1 Figure 1. Schematics of crystal symmetry and orbitals. (a) A reconstruction of octahedral crystal

Figure 2 . 20 Figure 3 .

 2203 Figure 2. (a) Oscillation of RHEED intensity during the growth of a LAO/ETO/(111)STO heterostructure(upper panel) and RHEED image along the [101 ̅ ] crystallographic direction for

21

 21

Figure 4 .

 4 Figure 4. (a) Ti L2,3 edge (left panel) and Eu M4,5 edge (right panel) XMCD spectra acquired at 2

ACKNOWLEDGMENTS

This project received funding from the ERA-NET QUANTERA European Union's Horizon H2020 project QUANTOX under Grant Agreement No. 731473, from the Ministero dell'Istruzione, dell'Università e della Ricerca for the PRIN project TOP-SPIN (Grant No. PRIN 20177SL7HC). DP has benefited support from the initiative of excellence IDEX-Unistra (ANR-10-IDEX-0002-02) from the French national program Investment for the future. We acknowledge discussion with R. Citro, M. Cuoco, A. Filippetti and A. Perroni.

XAS, XLD and XMCD spectroscopy measurements

The experiments of polarization dependent x-ray absorption spectroscopy (XAS) across the Eu-M4,5 and the Ti-L2,3 edges were performed at the beamline X-Treme of the Swiss Light Source [START_REF] Piamonteze | X-Treme Beamline at SLS: X-Ray Magnetic Circular and Linear Dichroism at High Field and Low Temperature[END_REF] and at the beamline ID32 of the European Synchrotron Research Facility (ESRF). [START_REF] Brookes | The Beamline ID32 at the ESRF for Soft X-Ray High Energy Resolution Resonant Inelastic X-Ray Scattering and Polarisation Dependent X-Ray Absorption Spectroscopy[END_REF][START_REF] Kummer | The High-Field Magnet Endstation for X-Ray Magnetic Dichroism Experiments at ESRF Soft X-Ray Beamline ID32[END_REF] XAS performed with circularly or linearly polarized photons can detect the magnetic moments and the 3d orbital crystal field splitting, respectively. The samples were mounted with the (111) surface in the vertical laboratory plane, and the vertical axis parallel to the [101 ̅ ] crystallographic direction determined by x-ray diffraction, i.e., parallel to one of the three equivalent Ti-Ti bonds within the hexagonal surface lattice (Fig. 3c). The energy resolution at both beamlines was around 100 meV. XAS, XLD and XMCD spectra were compared to theoretical spectra obtained by atomic multiplet scattering calculations using the CTM4XAS code. [START_REF] Stavitski | The CTM4XAS Program for EELS and XAS Spectral Shape Analysis of Transition Metal L Edges[END_REF]

SQUID magnetometry measurement

SQUID data were collected by using a Quantum Design MPMS3 magnetometer at the Université de Strasbourg, IPCMS. Magnetization measurements as a function of magnetic field were acquired on LAO/ETO//STO heterostructure. Data were corrected for the diamagnetism of the substrate substrating the linear contribution acquired at high magnetic fields.

Author Contributions

Conceptualization


Nicolas Cimetiere

Joseph De Laat
email: joseph.delaat@esip.univ-poitiers.fr

Henry's law constant of N,N-dichloromethylamine. Application to the contamination of the atmosphere of indoor swimming pools

Keywords: Henry's law constant, N,N-dichloromethylamine, Temperature dependence, Headspace technique, Swimming pool

The volatility of N,N-dichloromethylamine (DCMA), a disinfection by-product formed during chlorination of swimming pool water, has been investigated in the present work. The Henry's law constants for DCMA were experimentally determined at five temperatures (5, 15, 25, 35 and 45 °C) using the single equilibrium technique. The volumic ratio between the gas and the water phases in the headspace vessels ranged from 1 to 60 and the initial concentration of DCMA in the aqueous solutions was  2 mM. The values obtained for the dimensionless Henry's law constant varied from 0.047 at 5 °C to 0.312 at 45°C. The temperature dependence of the Henry's law constants followed the van't Hoff equation. Trace levels of DCMA (16-70 µg.m -3 ) were detected in the atmosphere of an indoor swimming pool whereas the concentrations of DCMA in water pools were in the range 8.8-15 µg.L -1 .

Introduction

Chlorination with chlorine gas or sodium hypochlorite is the most common process used for the disinfection of public swimming pool water. Along with its disinfectant action, chlorine also leads to the formation of numerous disinfection by-products such as chloramines, trihalomethanes, haloacetic acids, chloral hydrate and haloacetonitriles by the reaction with pollutants introduced by bathers. It has been estimated that a bather introduces in the water 25-80 mL.h -1 of urine and 200-1000 mL.h -1 of sweat [START_REF] Gunkel | The problem of urea in bathing water[END_REF][START_REF] Erdinger | Potassium as an indicator of anthropogenic contamination of swimming pool water[END_REF]. These human materials contain nitrogenous compounds (urea, ammonia, creatinine, amino-acids, amines, …) and the amount of nitrogenous pollution released by a bather is about equal to 0.85 g N.h -1 [START_REF] Seux | Evolution de la pollution apportée par les baigneurs dans les eaux de piscines sous l'action du chlore[END_REF].

It is well-known that reactions of aqueous chlorine with ammonia at neutral pH lead to the formation of monochloramine and dichloramine as intermediates and to nitrogen gas as end-product [START_REF] Jafvert | Reaction scheme for the chlorination of ammoniacal water[END_REF]. In addition to N 2 , NO 3 -is a stable end-product of reactions between chlorine and ammonia. The reactions are largely controlled by the chlorine:ammonia nitrogen ratio (Cl:N), pH, buffer and ionic strength . At Cl:N less than 7.6:1 by weight, chlorine is present as combined chlorine (inorganic chloramines) and at Cl:N higher than 7.6:1, free chlorine is present in water. Small amounts of nitrogen trichloride (or trichloramine) and nitrate ions are also produced.

The reactions of chlorine with nitrogenous components present in urine and sweat can also lead to the formation of organic and inorganic chloramines, nitrate as well as many potentially harmful DBPs [START_REF] Judd | The fate of chlorine and organic materials in swimming pools[END_REF]. By using a new analytical method (membrane introduction mass spectrometry), [START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF] have shown that nitrogen trichloride can also be formed by chlorination of four model compounds present in urine (urea, creatinine, histidine and arginine). The reaction of chlorine with creatinine leads to the formation of N,N-dichloromethylamine (CH 3 NCl 2 , DCMA) [START_REF] Tachikawa | Occurrence and production of chloramines in the chlorination of creatinine in aqueous solution[END_REF][START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF] and [START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF] first reported the occurrence of CH 3 NCl 2 in a swimming pool water.

Chloramines and in particular inorganic chloramines have been suggested to largely contribute to the typical smell and to the irritant properties of the atmosphere of indoor swimming pools [START_REF] Héry | Exposure to chloramines in the atmosphere of indoor swimming pools[END_REF]. Nitrogen trichloride is the most volatile inorganic chloramine [START_REF] Holzwarth | The fate of chlorine and chloramines in cooling towers. Henry's law constants for flashoff[END_REF] and is easily released into the air atmosphere of indoor pools. It has been shown to be a powerful irritant for upper respiratory tracts [START_REF] Gagnaire | Comparison of the sensory irritation response in mice to chlorine and nitrogen trichloride[END_REF]. Therefore, most of the studies related to the health risks associated with an exposure to DBPs present in the atmosphere of swimming pools have focused on nitrogen trichloride. Published values for nitrogen trichloride concentrations in the atmosphere of swimming pools range from below the detection limit ( 0.05 mg.m -3 ) to about 2 mg.m -3 [START_REF] Jacobs | Exposure to trichloramine and respiratory symptoms in indoor swimming pool workers[END_REF]. For swimming pool workers, a dose/response relationship was obtained for eye irritation, nasal and throat symptoms (Massin et al., 1998a;Massin et al., 1998b).

In France, NCl 3 concentrations in the atmosphere of indoor pools are not regulated. However, the French institute of Occupational Health and Safety (INRS) has proposed a comfort limit value of 0.5 mg.m -3 [START_REF] Héry | Exposure to chloramines in the atmosphere of indoor swimming pools[END_REF] even 0.3 mg.m -3 [START_REF] Gérardin | Réduction de l'exposition des travailleurs au trichlorure d'azote par action sur les procédés dans deux secteurs d'activités[END_REF] for the concentration of NCl 3 in the atmosphere of indoor pools. Carbonelle et al. (2002) have shown that trichloramine in air increases the lung epithelium permeability at concentration levels as low as 0.3 mg.m -3 . In Belgium, the Government of the Brussels-Capital Region has set a recommended maximum value of 0.3 mg.m -3 and a maximum admissible value of 0.5 mg.m -3 for the concentration of NCl 3 in the atmosphere of indoor swimming pool (Gouvernement de la Région de Bruxelles-Capitale, 2002).

Organic chloramines formed by the reaction of chlorine with organic nitrogen contaminants may i) contaminate the atmosphere of swimming pools, ii) give a positive interference in the determination of NCl 3 concentration and iii) contribute to the irritant properties of the swimming pool air. Among the organic chloramines, low molecular weight compounds such as DCMA might be volatile. To our knowledge, the Henry's law constant of this chlorination by-product of creatinine is unknown. Therefore, the main objective of the present work was to determine the Henry's law constant of DCMA in the water/air system at various temperatures ranging between 5 and 55 °C. A second objective of this work was to determine the concentration of DCMA in a pool water and in the atmosphere of an indoor swimming pool of Poitiers.

Material and methods

Reagents and preparation of standard solutions of DCMA

All solutions were prepared from reagent-grade chemicals and purified water delivered by a

Millipore system (Milli RX75/Synergy 185). Stock solutions of methylamine (methylamine hydrochloride, Acros Organics, 99%) and of free chlorine (prepared from sodium hypochlorite, Acros Organics, 13 %) were stored in the dark at 4 °C.

Stock solutions of N,N-dichloromethylamine (2 mmol.L -1 ) were prepared a few minutes before use by slowly adding free chlorine (pH  8) to a solution of methylamine chloride in a well-stirred reactor and using a chlorine:methylamine chloride molar ratio of 2:1. pH of the final solution was adjusted to pH 3.5  0.5 with hydrochloric acid (37 % by weight). The change of the UV/Visible absorbance with reaction time indicates that the reaction was achieved within a few seconds at pH < 4, in agreement with the previous works of [START_REF] Poncin | Kinetic studies on the chlorination of methylamine by sodium hypochlorite in dilute aqueous medium[END_REF].

Henry's law constant determination

Air/water partition coefficients or Henry's law constants were experimentally determined by using a static method with variable headspace. This method can only be applicable for sufficiently volatile compounds to be analytically admissible [START_REF] Mcaullife | GC determination of solutes by multiple phase equilibration[END_REF][START_REF] Gossett | Measurement of Henry's Law Constants for C1 and C2 Chlorinated Hydrocarbons[END_REF]. In this study, a volume V L of a freshly prepared aqueous solution of DCMA (C 0,L  2 mmol.L -1 ) was put into a vial (Total volume : V T : 44  1 mL) or into a bottle (V T : 260  2 mL). The total volume of each vial or bottle was exactly determined. The stock solution, the vessel and the air in the vessel were thermostated at the desired temperature for at least one hour prior to the introduction of the DCMA solution. The vessel was hermetically sealed with a PTFE-faced-silicone septum, immediately after introduction of the DCMA solution. For each run, various ratios of inner gaseous headspace volume to the liquid volume (1 < V G /V L < 60) were used in order to show the dependency of C eq,G with C eq,L . The vessels were set on a rotary tumbler for agitation in a thermostated room (Lovibond) for at least one hour.

Preliminary experiments demonstrated that the gas/liquid equilibrium was achieved within 10 minutes (Fig. S4 in Supplementary Content). After an agitation time of 1 hour, the vial or the bottle was opened, the liquid phase was immediately transferred in a quartz cell (without using a pipette) for a UV-spectroscopic determination of equilibrium DCMA concentration in the liquid phase (C eq , L ) at 302 nm. For each run, some vials were filled without headspace in order to confirm that DCMA was not decomposed under our experimental conditions.

Determination of DCMA level in the atmosphere of an indoor swimming pool

An in situ measurement method was developed to determine DCMA level in the atmosphere of an indoor swimming pool located at Poitiers (France). The sampler consisted of two 250-mL pyrex® glass traps, containing 200 mL sodium sulfite (10 mmol.L -1 ) and phosphate buffer (10 mmol.L -1 , pH = 3.5-4) (Fig. S1 in Supplementary Content). The sampling flow rate was 1.5 L.min -1 and the sampling period varied from 2 to 4 h. The two traps were connected in series and no significant breakthrough of DCMA from the first trap to the second one has been observed during any sampling. Under these conditions, DCMA present in the atmosphere of the swimming pool is absorbed in the traps and instantaneously reduced into the non volatile methylammonium ion (pKa = 10.62; [START_REF] Hall | Correlation of the Base Strengths of Amines[END_REF] by sodium sulfite:

CH 3 NCl 2 + 2 SO 3 2-+ H 3 O + + H 2 O  CH 3 NH 3 + + 2 SO 4 2-+ 2 HCl
It should be noted that other volatile chloramines such as trichloramine can also be reduced by sodium sulfite in the traps.

In order to determine the amounts of DCMA which have been trapped, the trapping solution was transferred into a 200 mL volumetric flask and completed to 200 mL with purified water (Fig. S1 in Supplementary Content). A volume of 20 mL of the trapping solution was introduced into a 25 mL volumetric flask (five replicates for each trap). Methylammonium ion present in the trapping solution was converted into DCMA with sodium hypochlorite in a 10 % molar excess (calculated from initial sulfite concentration) and the solution was then adjusted to pH 3.5  0.5. DCMA concentration was then determined by HPLC and UV detection at 210 nm.

Determination of DCMA in pool water

A volume of 10.0 L of water was collected from a spa and from a basin located in the indoor swimming pool of Poitiers. DCMA was then stripped with an air flow rate of 1.5 L.min -1 during 10 h at 45°C and the desorbed DCMA was then reduced into methylammonium ion with sodium sulfite by passing the gas through the two same traps than those used for air sampling. It should also be noted that a slow oxidation of sodium sulfite by oxygen occurred in the traps but the concentration of sodium sulfite was always greater than 2 mmol.L -1 after 10 h of air sparging.

To assess this method, 10.0 L of a standard aqueous solution of DCMA ([DCMA] = 1.9 µmol.L -1 ) has been analyzed under identical conditions. The data showed that DCMA was quantitatively desorbed and trapped in the first absorption vessel. The recovery yield was 104  3 % and DCMA was not detected in the second trap.

Analytical methods

Free chlorine and DCMA concentrations in the stock solutions were determined iodometrically by were determined iodometrically in order to confirm the concentrations of DCMA.

DCMA concentration in the traps was determined by HPLC because of the possible presence of other chloramines. HPLC analyses were performed with a Waters system (W600 controller;

W717plus Autosampler) equipped with a Zorbax 300Extend-C18 column and a UV detector at 210 nm (Waters W486). DCMA was eluted with a methanol (20%) / phosphate buffer (0.1 mol.L -1 ; pH = 4) (80%) mobile phase at a flow rate of 0.8 mL.min -1 . All HPLC analyses were made on five replicates with a sampling volume of 200 µL. Under these analytical conditions, the absolute detection and quantification limits were evaluated as 3 x 10 -12 and 9 x 10 -12 mol, respectively.

Relative quantification limit depends on the concentration factor applied during sampling collection and on the volume of reactants added before analyses.

Background and expressions used

Data format

The two-phase equilibrium between air and water phases for a volatile solute in a dilute aqueous solution can be represented by Henry's law coefficients or by air/water partition coefficients. Since these coefficients can be expressed by a multitude of forms and associated units, Table 1 reports the symbols and the expressions that will be used in this study.

Calculations of Henry law coefficients for DCMA

If DCMA is not decomposed under our experimental conditions, equilibrium concentration in the gas phase (C i , G ) can be deduced from the mass balance expression :

C i,G = G L V V (C 0,L -C i,L ) (1)
The molar fractions of DCMA in water and in air were calculated as follows:

x i = /MW C C L L L i, L i,    = /M W C L L L i,



for dilute aqueous solutions (2)

y i = /RT P C T G i,  = G T L V P RTV (C 0,L -C i,L ) (3)
where  L is approximately equal to 1000 kg.m -3 for the temperature range studied in this work, MW L is the molecular weight of water (0.018 kg.mol -1 ), R is the universal gas constant (R = 8.314 Pa.m 3 .mol -1 K -1 = 8.206 x 10 -5 atm.m 3 .mol -1 .K -1 ).

The various Henry's law constants listed in Table 1 were then calculated from C i,L , C i,G , x i and y i values. Assuming that P T = 1 atm, the relationship between the various Henry's law constants The HPLC chromatograms of stock solutions of DCMA prepared at pH 3.5  0.5 with a chlorine:methylamine molar ratio of 2:1 showed only one peak corresponding to DCMA (Fig. S2 in Supplementary Content). Iodometric titrations showed no loss of oxidants because the total amount of chlorine was equal to the amount of free chlorine which has been introduced into the solutions of methylamine. These data indicate that stock solutions of DCMA only contain DCMA and that the concentration of DCMA is equal to the half of the initial concentration of free chlorine. (1 mole of DCMA is equivalent to 2 moles of free chlorine).

UV/Visible spectra of dilute solutions of DCMA (3.4 x 10 -4 M < [DCMA] < 3.5 x 10 -3 M) were recorded by using quartz cells (optical length: 1, 5 or 10 cm). As shown the spectrum presented in Fig. S3 in Supplementary Content, DCMA absorbs in the UV region. The spectrum presents an absorption band at 302 nm with a molar extinction coefficient equal to  302nm = 309 ± 18 mol - 1 .L.cm -1 (mean value of 39 absorbance values determined from different standard solutions of DCMA). Spectrophotometric analyses also demonstrated that the UV spectra of DCMA solutions stored in hermetically closed vials and filled without headspace did not change over a storage time of 15 h at 5 -55 °C and pH 3.5  0.5. These results confirm that DCMA did not decompose in water under our experimental conditions. Fig. S3 in Supplementary Content also shows that the UV/visible spectrum of DCMA is similar to the spectrum of dichloramine (NHCl 2 ) which presents an absorption band at 293 nm ( 293nm = 289 ± 5 mol -1 .L.cm -1 , mean value of 9 UV-absorbance values).

Henry's law constant of N,N-dichloromethylamine

The Henry's law constant was determined at five different temperatures by measuring the equilibrium concentration of DCMA in the aqueous phase (C i,L ). All the experimental conditions (C 0,L , V L , V G , T) and the equilibrium concentrations (C i,L and C i,G ) have been reported in Supplementary Content (Tables S1-S5 and Fig. S5).

From the experimental data, the various forms of the Henry's law constants reported in Table 1 (H cc , H' px , H' pc and H yx ) have been calculated by linear regression analysis. As shown the data in Fig. 1, straight lines were obtained for the dependence of Ci ,G with C i,L at 288 and 308 K as well as for the other temperatures studied (Tables S1-S5 in Supplementary Content).

Values of Henry's law constants determined in this work are reported in Table 2. The values obtained for the air-to-water concentration ratio (H cc ) vary from 0.048 at 5 °C to 0.335 at 45 °C. The temperature dependence of Henry's law constants is correctly described by the van't Hoff equation (Fig. 2):

T B A ln(H)  
where B = 4298.11 K (excepted for H cc ) and A, a constant which depends on the unit used for the Henry's law constants. Values of A and B have been reported in Table 2.

The solvatation enthalpy (H°s ol ) can be calculated from the slope of the ln(H) = f(1/T) straight line (for H constant other than H cc ) :

H°s ol = B R with R = 8.314 Pa.m 3 .mol -1 K -1
The solvatation enthalpy of -35.7 kJ.mol -1 obtained for DCMA is very close to the values published for NH 2 Cl, NHCl 2 and NCl 3 [START_REF] Sander | Compilation of Henry's law constants for inorganic and organic species of potential importance in environmental chemistry[END_REF] which are in the range 34.2-39.9 kJ.mol -1 (Table 3).

A comparison of the Henry's law constants (Table 3) or of the air-water partition coefficient (Fig. 3) of some disinfection by-products formed during chlorination of swimming pool water shows that DCMA is a little less volatile than chloroform, about 3 times less volatile than trichloramine but far much volatile than monochloramine and dichloramine. At the average temperature of recreation pools ( 30°C), the air-water coefficients (H cc values) of trichloramine, chloroform, DCMA, dichloramine and monochloramine are equal to 0.504, 0.195, 0.135, 1.75 x 10 -3 and 5.58 x 10 -4 respectively. If DCMA is produced in water pools at concentration levels of about 0.1 µmol.L -1 (10 µg.L -1 ), the theoretical concentration of DCMA in air in a closed system should be  1.9 mg.m -3 . This concentration might cause irritation to bathers and to pool attendance.

Concentration levels of DCMA in water pools and in the atmosphere of an indoor swimming pool.

The DCMA levels in water and in the atmosphere of an indoor swimming pool located at Poitiers have been determined between February 2009 and May 2009. The swimming pool has a 25 meter pool (436 m 3 ), a leisure pool (190 m 3 ), a children pool (10 m 3 ), a toboggan pool (26 m 3 ) and a spa (5 m 3 ). The total volume of the atmosphere of the indoor pool (15 000 m 3 ) can be contaminated by volatile compounds that escape from the five basins. Table 4 reports the main operating conditions of the swimming pool for the four campaigns (C1 to C4) made in this study.

DCMA was determined in water collected from the spa (campaigns C1 and C2) and from the 25 meter pool (campaigns C3 and C4). The spa water is totally renewed every night in order to maintain the concentration of combined chlorine below 0.6 mg Cl 2 .L -1 . The mean residence time of water in the 25 meter pool is 10 days. Analyses showed that the concentrations of DCMA in water samples ranged from 88 to 150 nmol.L -1 (8.8 and 15 µg.L -1 ) (Table 4) and that about 3 % of the combined chlorine concentration can be accounted for DCMA (Table 4). Our concentrations of DCMA in water are of the same order of magnitude than those obtained by [START_REF] Weaver | Volatile disinfection by-product analysis from chlorinated indoor swimming pools[END_REF] using membrane introduction mass spectrometry (MIMS) for the determination of DCMA in 11 water pools (1.6 µg.L -1 < [DCMA] < 50.5 µg.L -1 ). It should be noted that about 80 % of the spa water is renewed every night and there is always about 12-13 bathers in the spa.

The water samples collected from the 25 meter pool were chlorinated at an arbitrary dose of 14 mg Cl 2 .L -1 in a 10 L headspace-free bottle in order to show if the DCMA concentration could increase by increasing free chlorine concentration and reaction time. After a reaction time of 72 h, analyses showed that this postchlorination step led to a chlorine consumption of about 8 mg Cl 2 .L -1 and to an increase of the DCMA concentration by a factor of 2.5-3. The DCMA concentrations were in the range 230-270 nmol.L -1 (23-27 µg.L -1 ). This increase of the concentration of DCMA suggests that reactions leading to the formation of DCMA are not very fast, in agreement with the kinetic data obtained by [START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF] for the formation of DCMA from creatinine. Methylamine and creatinine are known to produce DCMA by chlorination [START_REF] Poncin | Kinetic studies on the chlorination of methylamine by sodium hypochlorite in dilute aqueous medium[END_REF][START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF]. As the concentration of methylamine in urine is about 10 times less than the concentration of creatinine [START_REF] Mitchell | Methylamine in urine[END_REF] most of the DCMA formed in swimming pool is probably formed from creatinine. On the basis of a production of 0.66 mole of DCMA per mole of creatinine [START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF]2007), the DCMA formation potentials measured in our study (250  20 nmol.L - 1 ) should correspond to a concentration of creatinine in the water samples of 350-400 nmol.L -1 . This is fully consistent with the concentrations of creatinine which can be estimated from the concentrations of urea of 30-50 µmol.L -1 (1.8-3.0 mg.L -1 ) which have been measured in this swimming pool water (unpublished data) and from a mean molar ratio of urea to creatinine of about 100 in urine.

As far as air analyses are concerned, air was sampled at a distance of about 4 m from the two main basins and at a height of 1.5 m. DCMA concentrations in air were in the range 0.16-0.70 µmol.m -3 (16-70 µg.m -3 ) and far below the theoretical DCMA concentration in air (1.5-2.0 mg DCMA m -3 ) that should be in equilibrium with the aqueous concentrations of DCMA at 30 °C (11-15 µg.L -1 ) at 30 °C (Table 4). The concentration of DCMA was higher for the first campaign (70 ± 30 µg.m -3 ) than for the second one (16 ± 2 µg.m -3 ). This was consistent with the observations made by the attendant pools who reported that the atmosphere was found to be more irritant during the first campaign of analyses. It should also be noted that the lowest value of DCMA concentration obtained during the second campaign might be attributed to the fact that the ventilation rate was higher and that all the roof windows were opened.

Nowadays, the concentrations of DCMA in water and in the atmosphere of indoor swimming pools are not controlled. As trichloramine exposure is considered to be at the origin of eye and airway irritations, only trichloramine concentrations have been determined in the atmosphere of indoor pools (Massin et al., 1998b;[START_REF] Jacobs | Exposure to trichloramine and respiratory symptoms in indoor swimming pool workers[END_REF]. In these studies, the analytical method used for the determination of trichloramine in the atmosphere of indoor pools is based on the reduction of trichloramine by arsenic trioxide and subsequent determination of the chloride ion formed [START_REF] Héry | Exposure to chloramines in the atmosphere of indoor swimming pools[END_REF]. The concentration of trichloramine in water pools can also be determined by this method after a gas/liquid desorption of trichloramine. The data obtained in the present work suggest that the use of this method might lead to an overestimation of the concentration of trichloramine in the presence of other volatile organic chloramines such as DCMA. The membrane introduction mass spectrometry (MIMS) is a more convenient method for the differentiation and the quantification of inorganic and organic chloramines in water and in air [START_REF] Weaver | Volatile disinfection by-product analysis from chlorinated indoor swimming pools[END_REF].

Conclusion

The Henry's law constants determined in the present work show that N,N-dichloromethylamine (DCMA) is quite as volatile than trichloramine. This work confirms the presence of DCMA in pool water previously reported by [START_REF] Li | Volatile disinfection byproduct formation resulting from chlorination of organic -Nitrogen precursors in swimming pools[END_REF] and [START_REF] Weaver | Volatile disinfection by-product analysis from chlorinated indoor swimming pools[END_REF] and this work demonstrates for the first time the presence of DCMA in the atmosphere of indoor swimming pools. These data suggest that the health risks associated with the formation of this volatile organic chloramine during chlorination of swimming pool water must be considered in further epidemiological studies. Additional research is also needed in order to develop specific analytical methods for the determination of DCMA and trichloramine and to evaluate the irritant properties for eyes and the respiratory tracts of DCMA and for mixture of DCMA and trichloramine.

  using the US Standard Method 4500B(APHA, 2005) with sodium thiosulfate 0.01 mol.L -1 (Prolabo, > 99.9%). In the case of DCMA, two moles of free iodine is liberated by the reaction with KI :CH 3 Cl 2 + 4 I -+ 4 H +  CH 3 NH 2 + 2HCl + 2 I 2 pH was measured with a MeterLab PHM 250 pH-meter (Radiometer) equipped with a glass electrode (Radiometer, CL114 JB). UV spectra of DCMA were recorded from 200 to 350 nm with a Cary 50 (Varian) and treated with WinUV 3.0 software (Varian). DCMA concentrations in dilute solutions (0.35-3.5 mmol.L -1 )

(

  Purity, stability and UV spectrum of aqueous solutions of DCMA at pH 3.5  0.5

Acknowledgments

The authors thank the French Agency for Environmental and Occupational Health Safety (AFSSET) for funding this study (Project n° EST-2008-07).

Carbonnelle, S., Françaux, M., Doyle, I., Dumont, X., de Burbure, C., Morel, G., Michel, O., Bernard, A., 2002. Changes in serum pneumoproteins caused by short-term exposures to nitrogen trichloride in indoor chlorinated swimming pools Biomarkers,7,[464][465][466][467][468][469][470][471][472][473][474][475][476][477][478]


Nicolas Cimetiere

Florence Dossier-Berne

Joseph De Laat
email: joseph.delaat@esip.univ-poitiers.fr

Effect of some parameters on the formation of chloroform during chloramination of aqueous solutions of resorcinol

Keywords: Monochloramine, Monochloramine demand, Trihalomethane, Free chlorine, Monochloramine hydrolysis. 1

The effects of various factors (N/Cl ratio used to prepare monochloramine, monochloramine doses, pH and contact time) on the monochloramine demand and on the chloroform yield during chloramination of resorcinol have been investigated. Chloramination experiments were carried out at 24  1 °C, at pH values ranging from 6.5 to 12 using a bicarbonate/carbonate buffer and preformed monochloramine solutions prepared at pH 8.5 with N/Cl ratios ([NH4Cl]0/[Total free Cl2]0 ranging from 1.0 to 150 mol mol -1 ). Monochloramine determinations ([Resorcinol]0 = 5 or 100 µM, [NH2Cl]0/[Resorcinol]0 = 20 mol mol -1 , pH = 8.5  0.1) showed a slow increase of the monochloramine consumption with reaction time. The monochloramine demands after reaction times of 7 days ([Resorcinol]0 = 100 µM) and 14 days ([Resorcinol]0 = 5 µM) were equal to 8.5 moles of NH2Cl /mole of resorcinol and were higher than the chlorine demands ( 7.3 moles/mole).

Chloroform yields from monochloramination of resorcinol were lower than 8 % (< 80 mmoles of CHCl3/mole of resorcinol) and were less than the yields obtained by chlorination (0.9 -0.95 mole/mole). Chloroform productions increased with increasing monochloramine dose and reaction time and decreased with increasing pH values within the pH range 6.5 to 10. Chloroform formation markedly decreased when the N/Cl ratio increased from 1 to 1.5 mol/mol and was suppressed at N/Cl > 100 mol/mol. The data obtained in the present work suggest that free chlorine released from monochloramine hydrolysis plays a significant role on the formation of chloroform during chloramination of resorcinol at N/Cl ratios close to unity (1.0 < N/Cl < 1.5).

Introduction

Chlorination is the most widely used process for drinking water disinfection. However, reactions between chlorine and naturally occurring organic matter, bromide and iodide produce numerous disinfection by-products (DBPs) of public health concern. Since the discovery of trihalomethanes (THMs) in chlorinated drinking waters in 1974 [START_REF] Bellar | Occurence of organohalides in chlorinated drinking waters[END_REF][START_REF] Rook | Formation of haloforms during chlorination of natural waters[END_REF], more than 600 DBPs have been reported in the literature [START_REF] Richardson | Occurrence, genotoxicity, and carcinogenicity of regulated and emerging disinfection by-products in drinking water: A review and roadmap for research[END_REF][START_REF] Hrudey | Chlorination by-products, public health risk tradeoffs and me[END_REF]. Because of health risks associated to long term exposure to DBPs, most developed nations have set regulations or guidelines to minimize the concentrations of THMs and of other DBPs in finished drinking water.

Many utilities have adopted the use of monochloramine as a secondary disinfectant in order to meet the current or future regulations concerning the concentrations of DBPs in drinking water.

Monochloramine is an efficient disinfectant. Monochloramine is less reactive with natural organic matter (NOM) and its use provides long lasting residuals in the water distribution systems. Total concentrations of THMs, haloacetic acids or of DBPs in chloraminated drinking waters or chloraminated solutions of natural organic matter (NOM) are lower than in chlorinated ones [START_REF] Diehl | DBP formation during chloramination[END_REF][START_REF] Hua | Comparison of disinfection byproduct formation from chlorine and alternative disinfectants[END_REF][START_REF] Yang | Factors affecting formation of haloacetonitriles, haloketones, chloropicrin and cyanogen halides during chloramination[END_REF][START_REF] Goslan | A comparison of disinfection by-products found in chlorinated and chloraminated drinking waters in Scotland[END_REF]. The production and the speciation of DBPs formed by chloramination depend of NOM depend on various parameters such as temperature, pH, bromide ion concentrations, nature of NOM, contact time and the application modes of monochloramine (introduction of free chlorine into ammonia containing water or introduction of preformed monochloramine solution) and the N/Cl initial ratios used for the production of monochloramine [START_REF] Qi | Formation of haloacetic acids during monochloramination[END_REF][START_REF] Yang | Factors affecting formation of haloacetonitriles, haloketones, chloropicrin and cyanogen halides during chloramination[END_REF][START_REF] Yang | Correlations between organic matter properties and DBP formation during monochloramination[END_REF]. The kinetics and the reaction pathways for the formation of DBPs by chloramination of MON are not well documented.

As demonstrated by Valentine's research group, the oxidation of MON by monochloramine can be attributed to two distinct reaction pathways: direct reactions of monochloramine with MON and indirect reactions involving free chlorine liberated from monochloramine hydrolysis [START_REF] Jafvert | Reaction scheme for the chlorination of ammoniacal water[END_REF][START_REF] Duirk | Modeling monochloramine loss in the presence of natural organic matter[END_REF][START_REF] Duirk | Modeling dichloroacetic acid formation from the reaction of monochloramine with natural organic matter[END_REF]:

NH2Cl + H2O  HOCl + NH3
(1)

The value of the equilibrium constant for reaction (1) (K = 5.1 x 10 -12 M; [START_REF] Morris | A critical review of kinetic and thermodynamic constants for the aqueous chlorine-ammonia system[END_REF] indicates that free chlorine should be present at trace levels in organic-free solutions of monochloramine. Free chlorine concentration depends on the total concentration of oxidants, pH and the initial N/Cl ratios used to prepare monochloramine solutions. In the presence of organic compounds, free chlorine can be consumed by the organic solutes. This reaction shifts the equilibrium (1) to the right and thus increases the decay rate of monochloramine. The contribution of free chlorine to the degradation of organic solutes and to the formation of DBPs during chloramination depend on various parameters such as the free chlorine concentration and the relative rate constants for the reactions of monochloramine and chlorine with the organic solutes.

In order to better understand the role played by free chlorine on the formation of DBPs during chloramination of organic compounds, solutions of resorcinol have been chloraminated in the present study and the production of chloroform have been determined at various N/Cl ratios.

Resorcinol has been used in this work as a model compound for the following reasons.

Metapolyhydroxybenzenic moieties in NOM have been assumed to be responsible for the formation of THM precursors upon chlorination [START_REF] Rook | Chlorination reactions of fulvic acids in natural waters[END_REF]Nordwood et al., 1980). Resorcinol is one of the most efficient precursors of chloroform upon chlorination with the formation of 0.9-0.95 mole of chloroform per mole of resorcinol (Nordwood et al., 1980[START_REF] De Laat | Chlorine demand and reactivity in relationship to the trihalomethane formation. Incidence of ammoniacal nitrogen[END_REF][START_REF] Boyce | Reaction pathways of trihalomethane formation from the halogenation of dihydroxyaromatic model compounds for humic acid[END_REF][START_REF] Gallard | Chlorination of phenols: kinetics and formation of chloroform[END_REF]. The rate of formation of chloroform is very fast because chloroform production is possible at very low chlorination doses ([Chlorine]0/[Resorcinol]0 < 1 mole/mole) and can compete with the very fast reaction of formation of monochloramine when solutions of resorcinol are chlorinated in the presence of ammonia [START_REF] De Laat | Chlorine demand and reactivity in relationship to the trihalomethane formation. Incidence of ammoniacal nitrogen[END_REF].

Furthermore, the absolute second-order rate constants for the reaction of hypochlorous acid [START_REF] Rebenne | Aqueous chlorination kinetics and mechanism of substituted dihydroxybenzenes[END_REF] and monochloramine [START_REF] Cimetiere | Monochloramination of resorcinol: mechanism and kinetic mechanism[END_REF] with the acid-forms (Ar(OH)2, Ar(OH)O -and Ar(O -)2) of resorcinol, 4-chlororesorcinol and 4,6-dichlororesorcinol are now known.

The apparent second-order rate constants for the reaction of free chlorine and monochloramine with resorcinol in the pH range 7-11 indicate that free chlorine is  5 orders of magnitude more reactive than monochloramine [START_REF] Cimetiere | Monochloramination of resorcinol: mechanism and kinetic mechanism[END_REF]. These kinetic data suggest that free chlorine reactions may contribute to the formation of chloroform during chloramination of resorcinol.

Therefore, this work was designed to examine the effect of some parameters (N/Cl ratio, pH, monochloramine dose, reaction time) on the production of chloroform during chloramination of resorcinol.

Experimental Section

Chemicals and Preparation of Solutions

All solutions were prepared from reagent-grade chemicals and purified water delivered by a

Millipore system (Milli RX75/Synergy 185). Stock solutions of resorcinol (Sigma-Aldrich, ACS reagent, > 99 %), ammonium chloride (Fisher Scientific, > 99%) and free chlorine (NaOCl, 13 %, Acros Organics) were stored in the dark at 4 °C. Solutions of monochloramine (0.1 to 4 mM) were prepared daily by slowly adding free chlorine into an ammonium chloride solution in a well-stirred reactor. The pH of the sodium hypochlorite and ammonium chloride solutions were adjusted to 8.5  0.1 just before use. Solutions of monochloramine were prepared by using a nitrogen to chlorine molar ratio (N/Cl = [Ammonium chloride]0/[Total chlorine]0) ranging from 1.0 to 150 mol mol -1 .

Iodometric titrations and spectrophotometric analyses of freshly prepared solutions of NH2Cl

showed that free chlorine introduced in the solution of ammonium chloride was quantitatively converted into NH2Cl at pH  8.5.

Phosphate buffer was not used in the present work because phosphate ions are known to catalyze the autodecomposition of NH2Cl [START_REF] Valentine | General acid catalysis of monochloramine disproportionation[END_REF] and are suspected to promote the resorcinol ring opening during chlorination [START_REF] Heasley | Aqueous chlorination of resorcinol[END_REF]. Due to the low catalytic effect of carbonate ion toward monochloramine disproportionation [START_REF] Jafvert | Reaction scheme for the chlorination of ammoniacal water[END_REF], a carbonate buffer (20 mM) was used in the present work. Initial pH was adjusted with hydrochloric acid or sodium hydroxide. All glassware was cleaned with a solution of monochloramine ( 1 mM) and was kept under monochloramine before the experiments.

Experimental Procedures

All chloramination experiments were conducted in batch reactors at 24 ± 1 °C.

The monochloramine demands of resorcinol were determined at pH 8.5  0.2 (carbonate buffer: 20 mM), with two initial concentrations of resorcinol ([Resorcinol]0 = 5 and 100 µM) and by using an excess of monochloramine ([NH2Cl]0/[Resorcinol]0 = 20 mol mol -1 ). Reactions were initiated by adding, under vigorous mixing, an appropriate volume of a stock solution of resorcinol (5 or 100 mM) into a freshly prepared carbonate buffered solutions of monochloramine (100 µM or 2 mM).

Blank solutions ([Resorcinol]0 = 0 mM) were systematically prepared under the same conditions in order to evaluate the monochloramine autodecomposition.

To determine the effects of pH (6.5 < pH < 12), N/Cl ratio (1  N/Cl  150 mol mol -1 ) and monochloramine dose (0

 [NH2Cl]0/[Resorcinol]0  65 mol mol -1 )

Analytical methods

Free chlorine and total chlorine concentrations in the stock solutions of sodium hypochlorite and monochloramine were determined iodometrically with sodium thiosulfate 0.01 M (Prolabo, > 99.9 %). Monochloramine concentrations were determined spectrophotometrically using a molar extinction coefficient of 446 M -1 cm -1 for NH2Cl at 244 nm.

At various reaction times, total chlorine was determined spectrophotometrically by the I3 -method.

For an initial concentration of NH2Cl of 2 mM, 0.1 mL sample of the solution was rapidly transferred into a 10 mL vial containing 4 mL of KI 0.2 M (Acros Organics, > 99 % for analysis)

and 100 µL of acetic acid 1.7 M (Riedel-de Haën, 99 -100 %). After a reaction time of 2 or 3 minutes, the mixture was then transferred into a 1 cm-spectrophotometer quartz cell and the concentration of triiodide ion was determined by measuring the absorbance at 351 nm (Shimadzu UVmini 1240 spectrophotometer). The concentration of monochloramine was calculated using a molar extinction coefficient of 2.69 x 10 4 M -1 cm -1 for I3 -at 351 nm. This value which has been determined in the present work from calibration curves obtained from various standard solutions of chlorine or monochloramine was consistent with values reported in literature [START_REF] Awtrey | The absorption spectra of I2, I3 -, I -, IO3 -, S4O6 -and S2O3 --. Heat of the reaction I3 -= I2 + I[END_REF][START_REF] Bichsel | Determination of iodide and iodate by ion chromatography with postcolumn reaction and UV/visible detection[END_REF].

CHCl3 was analyzed by GC/ECD using a gas chromatograph (Varian 3300 Star) equipped with an automatic headspace sampler (DANI HSS3950), a transfer line (heated at 250 °C) between the sampler and the injector of the chromatograph, a capillary column (J&W/DB624 30 m x 0.53 mm) and an electron capture detector. Separation was carried out with nitrogen gas as carrier gas (1 mL.min -1 ). Operating temperatures are 250°C, 40 °C and 300°C for the injector, the oven and the detector, respectively. Vials were introduced in the headspace sampler for at least 50 min at 50 °C in order to reach gas-liquid equilibrium. Calibration curves were performed with standard chloroform solutions (0.01 -1.25 µM). Detection and quantification limits were 0.01 and 0.025 µM respectively.

In a few experiments ([Resorcinol]0 = 100 µM), the total organic carbon (TOC) was determined at the beginning and at the end of the experiment by a Shimadzu TOC analyzer (TOC-VCSH). S1 andS2 in Supplementary

Results

Monochloramine demand of resorcinol

Content. At a given pH, the overall rate of autodecomposition of monochloramine follows an apparent second-order kinetic law with respect to the concentration of monochloramine:

-

dt ] Cl NH [ d 2 = kapp [NH2Cl] 2
(2)

As the contribution of the autodecomposition reactions of NH2Cl to the overall consumption of monochloramine by resorcinol can not be always neglected, the monochloramine demand of resorcinol has been determined by subtracting the autodecomposition of monochloramine. The amount of monochloramine autodecomposed ([NH2Cl]a -[NH2Cl]b) between two consecutive sampling times ta and tb have been calculated from the integrated form of Eq(2):

b 2 ] Cl NH [ 1 - a 2 ] Cl NH [ 1 = kapp (tb -ta) (3) or [NH2Cl]b = (kapp (tb -ta) + a 2 ] Cl NH [ 1 ) -1 (4)
where [NH2Cl]a and [NH2Cl]b represent the monochloramine concentrations in organic-free solution at reaction times ta and tb, respectively and kapp, the apparent second-order rate constant of decomposition of monochloramine in organic-free solution. Fig. 2 presents monochloramine demands of resorcinol at pH 8.5  0.1 as a function of reaction time. The rates of monochloramine consumption increased with increasing concentrations of reactants. For an initial concentration of resorcinol of 100 µM ([NH2Cl]0 = 2 mM), the data showed that the N/Cl ratio (N/Cl = 1.2 and 5 mol mol -1 ) did not significantly affect the rate of monochloramine consumption and the monochloramine demands. Monochloramine demands of 5.5, 7.7 and 8.6 moles of NH2Cl/mole of resorcinol were measured at reaction times of 24, 72 and 168 h, respectively. For an initial concentration of resorcinol ([Resorcinol]0 = 5 µM, [NH2Cl]0 = 100 µM; N/Cl = 5 mol mol -1 ), the monochloramine demands were equal to 4.0, 6.5 and 8.4 mol mol -1 after reaction times of 48, 168 and 480 h, respectively (Fig. 2).

Data obtained from chlorination experiments of resorcinol ([Resorcinol]0 = 20 µM; [Free Chlorine]0 = 400 µM; Fig. S1) exhibit a very fast consumption of chlorine because the chlorine demands obtained after a reaction time of 5 minutes were nearly equal to 5.5 and 6.4 moles of chlorine/mole of resorcinol in carbonate and in phosphate buffered solutions of resorcinol, respectively. For a reaction time higher than 15 hours, the measured chlorine demands (7.2  0.1 mole/mole in phosphate buffer) were in good agreement with previous works [START_REF] Norwood | Reactions of chlorine with selected aromatic models of aquatic humic material[END_REF][START_REF] De Laat | Chlorine demand and reactivity in relationship to the trihalomethane formation. Incidence of ammoniacal nitrogen[END_REF]Reckhow and Singer, 1986). A comparison between the shapes of the chlorine and monochloramine consumption curves also demonstrates that monochloramine is much less reactive than chlorine. Monochloramine demands did not reach a plateau within a reaction time of 168 h and were higher than the chlorine demand of resorcinol. The complete oxidation of one mole of resorcinol into CO2 should require a theoretical amount of 12 moles of monochloramine. Since TOC removals in chloraminated solutions of resorcinol ([Resorcinol]0 = 100 µM) were less than 30 % after a reaction time of 7 days, the unexpected high monochloramine demands obtained in this work ( 8.5 moles/mole) suggest that oxidation by-products of resorcinol catalyze the decomposition of monochloramine. It is well known that proton donator inorganic species such as bicarbonate and phosphate ions catalyze the monochloramine disproportionation. Among the chloramination by products of resorcinol identified by [START_REF] Heasley | Reactions of resorcinol and its chlorinated derivatives with monochloramine: Identification of intermediates and products[END_REF], chloro-ketones and chlorocarboxylic acids may catalyze the decomposition of monochloramine.

Effect of N/Cl ratio upon chloroform formation

Preliminary experiments showed that chlorination of resorcinol ([Resorcinol]0 = 1 -50 µM;

[Cl2]0/[Resorcinol]0 = 20 -40 mol mol -1 , 7 < pH < 9, reaction time > 1 h) leads to a very fast production of 0.93  0.03 mol of chloroform/mole of resorcinol (Fig. S2). In addition, significant productions of chloroform can also be formed at very low chlorine/resorcinol ratios (< 1 mol/mol), indicating that the initial steps of chlorination of resorcinol represent rate limiting steps for the formation of chloroform (Fig. S3). All these data are in good agreement with literature [START_REF] De Laat | Chlorine demand and reactivity in relationship to the trihalomethane formation. Incidence of ammoniacal nitrogen[END_REF]Reckhow and Singer, 1986). ([NH2Cl]0 = 1 mM) and N/Cl ratios ranging from 1 to 150 mol mol -1 . The data show that chloroform productions slowly increased with reaction time (reaction time studied: 0 -96 h) and markedly decreased when the N/Cl ratio increased from 1 to 150 mol mol -1 (Figs 3 and4). For a reaction time of 48 h, chloroform yields decreased from  3.5 % for N/Cl = 1.02  0.02 mol mol -1 to 0.4 % for N/Cl = 100 mol mol -1 (a yield of 100 % is assumed to correspond to the formation of one mole of CHCl3/mole of resorcinol). As shown the curves in Fig. 4, the chloroform yield is very sensitive to the N/Cl ratio for N/Cl values between 1 and 1.5 mol mol -1 . It should also be noted that chloroform yield which is equal to 95 % for N/Cl = 0 mol mol -1 (chlorination experiments) decreases to 3.5 % when the N/Cl increases from 0 to 1.0 mol mol -1 .

Effect of monochloramine dose on chloroform formation

Solutions of resorcinol ([Resorcinol]0 = 100 µM) have also been treated at monochloramine-toresorcinol ratios ranging from 0 to 10 mol mol -1 at pH 8.5 and using various N/Cl ratios (1 < N/Cl < 10 mol mol -1 ). For a reaction time of 48 hours, GC/ECD analyses showed that chloroform was detected at concentrations close to the quantification limit in samples chloraminated at [NH2Cl]0/[Resorcinol]0 doses lower than 2 mol mol -1 (Fig. 5). For [NH2Cl]0/[Resorcinol]0 > 2 mol mol -1 , the chloroform yields increased with monochloramine dose and did nod exceed 0.5 % for the highest monochloramine dose used.

Effect of pH on chloroform formation.

To study the effect of pH on the formation of chloroform, solutions of resorcinol ([Resorcinol]0 = 10 µM) were treated at a monochloramine dose of 0.65 mM ([NH2Cl]0/[Resorcinol]0 = 65, N/Cl = 10 mol mol -1 ) and at pH values ranging from 6.5 to 12. Fig. 6 shows that the chloroform yield decreases from 7.7 % at pH 6.5 to  0.1 % at pH 10. At pH higher than 10, the concentrations of chloroform were below the quantification limit of 0.025 µM (chloroform yield : < 0.25 %). It should be noted that the decrease of the production of chloroform when the pH increases from pH 8 to 10 has also been previously observed for chlorination of resorcinol solutions [START_REF] De Laat | Chlorine demand and reactivity in relationship to the trihalomethane formation. Incidence of ammoniacal nitrogen[END_REF].

Discussion

Contribution of free chlorine to the initial steps of degradation of resorcinol

Free chlorine formed by the reversible hydrolysis reaction of monochloramine is always present in monochloramine solutions and can therefore contribute to the oxidation of organic solutes. In organic-free solutions, the concentration of free chlorine in equilibrium with monochloramine can be calculated from the equilibrium constant of the monochloramine hydrolysis reaction (reaction (1), K1 = 5 10 -12 M, [START_REF] Morris | A critical review of kinetic and thermodynamic constants for the aqueous chlorine-ammonia system[END_REF], the acid dissociation constants of HOCl (Ka1 = 10 -7.54 , [START_REF] Morris | The acid ionization constant of HOCl from 5 to 35°C[END_REF] and NH4 + (Ka2 = 10 -9.25 , [START_REF] Morris | The acid ionization constant of HOCl from 5 to 35°C[END_REF] and the total concentration of nitrogen species

([Total Nitrogen] = [NH4 + ] + [NH3] + [NH2Cl]).
By neglecting the concentration of the aqueous chlorine ([Cl2]aq) at pH > 5, the equilibrium concentration of free chlorine is equal to:

2α Chlorine] [Total 4α - β β Chlorine] [Free 2 eq + - = (4) with         + + = + + + ]) [H ])(K [H (K K ] [H K α a2 a1 1 a2 (5) and  = 1 +  ([Total Nitrogen] -[Total Chlorine]) (6)
In organic-free aqueous solutions of monochloramine, concentration of free chlorine should be at the minimum at pH  8.4 (pH = (pKa1 + pKa2)/2) and would decrease when the concentration of free ammonia increases (or when the N/Cl ratio increases). As depicted calculated curves in Fig. S4, the theoretical concentrations of free chlorine in solutions containing 1 mM of monochloramine should be very low :  6 x 10 -7 M for N/Cl = 1.0 mol mol -1 , 7 x 10 -10 M for N/Cl = 1.5 mol mol -1 and 4 x 10 -11 M for N/Cl = 10 mol mol -1 .

In the presence of organic compounds, free chlorine present in monochloramine solutions can react with organic solutes and these reactions accelerate the overall decay of monochloramine. As demonstrated in a previous paper [START_REF] Cimetiere | Monochloramination of resorcinol: mechanism and kinetic mechanism[END_REF], the importance of these indirect reactions involving free chlorine in the overall rate of degradation of organic solutes by monochloramination depend on various parameters such as N/Cl ratio, pH and the reactivity of free chlorine and monochloramine with the organic solutes. In the case of resorcinol, the absolute rate constants for the reactions of chlorine [START_REF] Rebenne | Aqueous chlorination kinetics and mechanism of substituted dihydroxybenzenes[END_REF] and monochloramine [START_REF] Cimetiere | Monochloramination of resorcinol: mechanism and kinetic mechanism[END_REF] with resorcinol are known. Therefore, the contribution of free chlorine reactions to the degradation of resorcinol can be calculated by kinetic modelling (Table S2). Fig. 7 presents computed data calculated for the following conditions: [Resorcinol]0 = 50 µM, [Monochloramine]0 = 4 mM, 6 < pH < 12, 1 < N/Cl < 2 mol mol -1 ). Fig. 7 shows that the contribution of free chlorine to the degradation of resorcinol is highly sensitive to the N/Cl ratio in the range 1.0-1.5 mol mol -1 . For the conditions used in Fig. 7, the contributions of free chlorine reactions to the degradation of resorcinol at pH 8 are equal to 25 % for N/Cl = 1.05, 15 % for N/Cl = 1.1, 8.2 % for N/Cl = 1.2 and 1.8 % for N/Cl = 2. These data demonstrate that trace amounts of free chlorine in monochloramine solutions can play a significant role on the degradation of resorcinol at Cl/N ratio close to unity because apparent rate constants for the reaction of free chlorine are about 5 orders of magnitude higher than the rate constants of mononochloramine.

Effect of N/Cl ratio and pH on the formation of chloroform

Kinetic modelling cannot be used to simulate chloroform formation during chlorination and chloramination of resorcinol because the reaction mechanisms and the rate constants of reactions leading to the formation of chloroform are not known. The reaction pathways for the formation of chloroform by chlorination of 1,3-dihydroxybenzenic model compounds such as resorcinol and orcinol have been investigated by several researchers [START_REF] Christman | Natural and model aquatic humics: Reactions with chlorine[END_REF][START_REF] Norwood | Reactions of chlorine with selected aromatic models of aquatic humic material[END_REF][START_REF] Rook | Possible pathways for the formation of chlorinated degradation products during chlorination of humic acids and resorcinol[END_REF][START_REF] Boyce | Reaction pathways of trihalomethane formation from the halogenation of dihydroxyaromatic model compounds for humic acid[END_REF][START_REF] Heasley | Aqueous chlorination of resorcinol[END_REF][START_REF] Tretyakova | Degradative pathways for aqueous chlorination of orcinol[END_REF]). These studies have shown that resorcinol degradation involves aromatic ring-substitution, oxidation, hydrolysis, decarboxylation, ring-contraction and ring-opening reactions. Chlorination of resorcinol leads to the formation of chlororesorcinols, non-cyclic and cyclic chloroketones and chloro-keto-carboxylic acids as intermediates, chloroform ( 0.9-0.95 mole/mole of resorcinol), CO2 and short-chain chlorinated acids as end-products. The reaction of monochloramine with resorcinol in water and in ether has been investigated by [START_REF] Heasley | Reactions of resorcinol and its chlorinated derivatives with monochloramine: Identification of intermediates and products[END_REF]. These authors showed that chloramination of aqueous solutions of resorcinol (pH 7) also leads to the formation of chlororesorcinols and of a pentachloro intermediate which undergoes ring-opening with NH2Cl/H2O to produce CO2 and noncyclic chloroketones without significant productions of chloroform ( 0.04 mole of CHCl3/mole of resorcinol).

The data obtained in the present work also confirmed that little chloroform is formed in the reaction of monochloramine with resorcinol. The highest chloroform yield measured in this work (0.077 mole CHCl3/mole of resorcinol) was obtained at pH 6.5 (Fig. 6). Compared to chlorination, the rate of chloroform formation by chloramination is very slow. The large differences for chlorination and chloramination with respect to the formation of chloroform (rates of formation and chloroform yields) might be attributed to the fact that free chlorine is much more reactive than monochloramine with resorcinol, its chlorinated derivatives (chlororesorcinols) and with the ring-opening byproducts (chloroketones, chloroketo-carboxylic acids), the latter represent the precursors of chloroform through the haloform reaction. As reported above, chlorine is about 5 orders of magnitude more reactive than free chlorine with resorcinol, 4-chlororesorcinol and 4,6dichlororesorcinol but the rate constants for the reaction of monochloramine with chloroketones and chloroketo-carboxylic acids are unknown.

Conclusions

Chloramination of resorcinol leads to the formation of chloroform at much lower yields (< 8 %)

than by chlorination ( 90-95 %). Chloroform yields increased with increasing monochloramine dose and reaction time and with decreasing pH in the range 6.5 -10. The data also showed a remarkable decrease of chloroform productions when the N/Cl ratio used to prepare monochloramine increased from 1 to 1.5 mol mol -1 that is to say when the concentration of excess S2).

  on the chloroform yield from chloramination of resorcinol ([Resorcinol]0 = 10 -100 µM), the reaction was initiated by introducing under vigorous mixing an appropriate volume (typically 0.5 -2 mL) of a stock solution a resorcinol into 1000 mL of a monochloramine solution. The pH of the monochloramine solution was set up to the desired pH before the addition of resorcinol. After the addition of resorcinol to the monochloramine solution, the pH of the solution was readjusted and 10 mL of solution were immediately transferred into a series of 20 mL head-space vials which were then hermetically sealed with PTFE-faced-silicone septa. At least 250 mL of the remaining chloraminated solution of resorcinol was kept in an amber bottle at 24  1 °C and pH measurements performed at various reaction times indicate that pH did not vary by more than 0.2 pH unit. At timed intervals, the reaction of monochloramine with resorcinol in the vials was stopped by adding 100 µL of HCl (1 M) and 10 µL of sodium sulfite (2 M) through the septum. Each data point for chloroform yields represents the mean value of chloroform concentrations from two or three vials.Aqueous solutions of resorcinol have also been chlorinated in the present work in order to validate our experimental methods and to compare oxidant demands and chloroform productions from chloramination and chlorination experiments. The data obtained from chlorination experiments of resorcinol have been presented in the Supplementary Content.

Fig. 1

 1 Fig.1 presents examples of concentration-time decays of monochloramine ([NH2Cl]0 = 2 mM)obtained in the absence and in the presence of resorcinol ([Resorcinol]0 = 100 µM). Under the conditions used (carbonate buffer = 20 mM, pH = 8.5  0.1), a 18 % decay of monochloramine was observed in organic-free solutions after a reaction time of 7 days at pH 8.6. A much lower loss of monochloramine was obtained with an initial concentration of monochloramine of 100 µM ( 4 % decay after a reaction time of 14 days). The autodecomposition rates of monochloramine in organicfree solutions ([NH2Cl]0 = 100 µM and 2 mM) could be well simulated by the kinetic model of[START_REF] Jafvert | Reaction scheme for the chlorination of ammoniacal water[END_REF] which has been reported in TablesS1 and S2in Supplementary

Figs 3

 3 Figs 3 and 4 present chloroform yields obtained from solutions of resorcinol ([Resorcinol]0 = 50 µM) which have been chloraminated at pH 8.5 with a [NH2Cl]0/[Resorcinol]0 ratio of 20 mol mol -1

Fig. 1 -Fig. 2 -Fig. 7 -

 127 Fig. 1 -Concentration time profiles of monochloramine in the absence and in the presence of