Methylamine and dimethylamine photocatalytic degradation—Adsorption isotherms and kinetics

Applied Catalysis A: General 402 (2011) 201–207

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Methylamine and dimethylamine photocatalytic degradation—Adsorption isotherms and kinetics Sihem Helali a,b , Eric Puzenat a , Nathalie Perol a , Mohamed-J. Safi b , Chantal Guillard a,∗ a b

Université Lyon 1, CNRS UMR 5256, IRCELYON, 2 avenue Albert Einstein, F-69626 Villeurbanne, France Université Tunis El Manar, Ecole Nationale d’Ingénieurs de Tunis, ENIT, BP 37, le Belvédère 1002, Tunis, Tunisia

a r t i c l e

i n f o

Article history: Received 4 January 2011 Received in revised form 2 June 2011 Accepted 5 June 2011 Available online 13 June 2011 Keywords: Adsorption Dimethylamine Methylamine Photocatalytic degradation TiO2

a b s t r a c t The photocatalytic degradation of two nitrogenous organic compounds, methylamine CH3 NH2 and dimethylamine (CH3 )2 NH, used in pharmaceutical and chemical industries was investigated in the presence of UV-irradiated TiO2 aqueous suspensions. Different parameters were studied: adsorption under dark and UV-A conditions, photolysis, kinetics of degradation, and chemical pathway of methylamine (MA) and dimethylamine (DMA) degradation. The percentage of covered OH in the dark was equal for different concentrations of MA and DMA. The adsorption isotherms of these two amine compounds MA and DMA follow the Langmuir model. The photocatalytic oxidation kinetics of MA and DMA are described by the Langmuir–Hinshelwood model with a first order kinetics for concentrations below 0.5 mmol L−1 , and then reach a plateau. For both MA and DMA, the coverage rates of the TiO2 surface in the dark and under illuminated conditions were different. The nitrogen atoms were decomposed mainly to ammonium (NH4 + ). Nitrite (NO2 − ) was also formed but was rapidly oxidized to nitrate (NO3 − ). MA was detected as an intermediate product during the degradation of DMA. Formic acid (HCOOH), and two other products not identified were detected. These non-identified products do not correspond to formamide. Total Organic Carbon (TOC) analysis shows the presence of final slightly mineralised intermediate compounds. © 2011 Elsevier B.V. All rights reserved.

1. Introduction Photocatalysis involving TiO2 has proved to be a good alternative for eliminating chemical water contaminants that are hardly degradable by conventional biological treatment even at low concentration making these contaminants hazardous for the environment and the human health. Photocatalysis is based on the interaction between UV-light and catalyst producing highly reac• • • tive oxygen species (OH , O2 − , HO2 ) that can mineralize dissolved organic pollutants to CO2 , H2 O and inorganic constituents. Photocatalysis with TiO2 is considered as an advanced and green technology [1] for its non-toxic, clean and safe properties. It is widely applied in air [2–4] and water purification systems. In particular, photocatalysis using TiO2 was successfully used for degradation of nitrogen compounds [5,6], pesticides [7–10], insecticides [11], dyes [12,13] which are difficult to be treated by biological methods. Photocatalysis was also promising for microorganism elimination [14,15]. Even if presenting a high photocatalytic activity, titanium dioxide remains difficult to separate from treated water in practical

∗ Corresponding author. Tel.: +33 4 72 44 53 16; fax: +33 4 72 44 53 99. E-mail address: [email protected] (C. Guillard). 0926-860X/$ – see front matter © 2011 Elsevier B.V. All rights reserved. doi:10.1016/j.apcata.2011.06.004

applications. This point is however poorly taken into account in the literature. Kagaya et al. [16] reported that a coagulation technique was recommended to separate titanium dioxide from solution after organic contaminant degradation using basic aluminium chloride as a coagulating agent for rapid flocculation and sedimentation of the titanium dioxide. Malato et al. [17] have demonstrated a sedimentation technique for separating catalyst from treated water using a new granulated version of the well-known P-25 titanium dioxide named Degussa Aeroperl presenting a large particle size which facilitates sedimentation. Ultrafiltration [18] and cross-flow microfiltration [19] were also investigated to explore the separation of titanium dioxide from water. According to the theory of photocatalysis, the efficiency of photocatalysis decreases due to the recombination of photo-generated electron/hole (e− /h+ ) pairs. Jiang et al. [20] reported that photocatalysis can be enhanced by an electric field across a photocatalyst which can promote the electron/hole separation and can prevent the recombination process. Moreover, Wang et al. [21] reported that O2 can be added to the water solution to capture the excited electrons reaching the semiconductor surface. Two amine compounds, methylamine (MA) and dimethylamine (DMA), were chosen in the present research to be degraded by heterogeneous photocatalysis using TiO2 Degussa P25 as photocatalyst. MA and DMA are widely used in the chemical industry to

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Table 1 Physicochemical properties of the model compounds. Compounds

Empirical formula

Molecular weight (g mol−1 )

pKa

Solubility (g 100 mL−1 )

Density at 4 ◦ C

Methylamine Dimethylamine

CH5 N C2 H7 N

31.06 45.08

10.63 10.73

108 354

0.897 0.890

make pesticides (ziram, thiram, diuron), surfactants (alkyldimethylamine oxide), photographic chemicals, explosives, dyes and pharmaceuticals. They are also used in nylon industry to improve the tensile strength. Asatoor et al. [22] suggested that short-chain aliphatic amines such as methylamine and dimethylamine, if present in large amounts, are able to readily access the brain and spinal tissues and interfere with neurological function owing to their low molecular weight, ease of solubility in aqueous and lipid phases and their electron-rich amine function. Dimethylamine has been suspected as a possible neurotoxin in uraemic patients where it is sequestered intracellularly [23] and occurs in higher concentrations than normal in the intestine, blood, and brain tissues [24]. Lee and Yoon [25] reported that dimethylamine can form a carcinogenic nitrosamine, nitrosodimethylamine, by reaction in the body with nitrosant agents like nitrites from food or with bacterial or endogenous agents. Methylamine decomposition under oxidative conditions has been reported by Kantak et al. [26] indicating that at 623 K, a greater part of MA is initially converted to NH3 while a small NOx fraction is increasingly formed with the increase of the O2 concentration. Choi [27] has compared the photocatalytic degradation of dimethylamine ((CH3 )2 NH) using pure and platinized TiO2 suspensions. The photodegradation of dimethylamine on Pt/TiO2 was much faster and yielded products different from those obtained with pure TiO2 . About 30% of dimethylamine was converted into trimethylamine in the deaerated Pt/TiO2 suspension within 1 h of irradiation whereas no conversion into trimethylamine was observed in the presence of pure TiO2 . Photocatalytic oxidation of methylamine (MA) on titanium dioxide in aqueous and gas-phases was studied by Kachina et al. [28] indicating a maximum aqueous photocatalytic oxidation efficiency in alkaline media. Kachina et al. also studied the gas-phase photocatalytic oxidation of dimethylamine [29] showing the formation of volatile products such as ammonia, formamide, carbon dioxide and water. The objectives of the present research will therefore be focussed on a better understanding of the photocatalytic mechanism for the degradation of the MA and DMA amine compounds widely used in the chemical industry through the study of their adsorption kinetics under dark and UV conditions and through the evolution of their kinetics of photocatalytic degradation at different concentrations. The fate of carbon and nitrogen atoms will also be analyzed.

Purity (w/w% in water) 41% 40%

2.2. Adsorption experiments A volume of 30 mL of methylamine and dimethylamine wih different initial concentrations at natural pH was stirred under dark condition during 30 min. This time was necessary to reach the equilibrium between adsorption and desorption. A concentration of 1 g L−1 of TiO2 was used. The pH of the solutions were 9.7 ± 0.7 for methylamine and pH 10.7 ± 0.57 for dimethylamine varying with concentrations. (Table 4). To determine the quantities of methylamine and dimethylamine adsorbed on TiO2 , samples were filtered using a 0.45 ␮m Millipore filters and analyzed. 2.3. Photocatalytic experiments The photoreaction was carried out in a Pyrex cylindrical reactor using an optical window with a 12.5 cm2 area (Fig. 1). The radiant flux was provided by a high-pressure mercury lamp Philips HPK 125 W which provides maximum energy at 365 nm. An optical filter Corning 0.52 in conjunction with a circulating water system was installed to cut-off wavelength below 340 nm and to avoid from heating the solution under UV irradiation. The radiant flux was measured with a VLX-3W radiometer with a CX-365 detector (UV-A). A value of 6.5 mW cm−2 was found corresponding to about 1.5 × 1017 photons s−1 . To absorb all the photons transmitted by the lamp, a concentration of 1 g L−1 of TiO2 was used with a 30 mL volume of solution [30]. At different times of photodegradation, samples were taken and filtered using a 0.45 ␮m Millipore filter. 2.4. Analyses The concentrations of MA and DMA remaining after the adsorption and during the photocatalytic degradation process were determined using a Dionex DX-120 ion exchange chromatograph. For cation analysis, a Dionex Ion-Pac CS16 (5 × 250 mm) column was used with H2 SO4 (11 mmol L−1 ) as eluent at a flow rate of 1 mL min−1 . The analysis of anions was performed by using a Dionex Ion-Pac AS14A (4 × 250 mm) column with Na2 CO3 (8 mmol L−1 ) and NaHCO3 (1 mmol L−1 ) as eluent (1 mL min−1 ). A HPLC coupled with UV–vis detector (VARIAN Prostar 230) was used to identify the intermediate products using a SARASEP CAR-H (7.8 × 300 mm) column with H2 SO4 (5 mmol L−1 , 0.7 mL min−1 ) as eluent. Total Organic Carbon (TOC) was determined with a Shimadzu model TOC-Control V provided with an autosampler.

2. Experimental 3. Results and discussion 2.1. Materials 3.1. Adsorption The two amine compounds, methylamine (MA) and dimethylamine (DMA), were purchased from Fluka and employed as received. Solutions were prepared under different concentrations in Ultra-Pure water from Millipore water Milli-Q Plus185 equipment. The physico-chemical properties of MA and DMA are given in Table 1. The photocatalyst used was titanium dioxide Degussa P-25 with a surface area of 50 m2 g−1 composed of anatase (80%) and rutile (20%). This titanium dioxide compound presents non-porous particles with a mean crystallite size of 30 nm.

To determine the effect of the initial concentration of methylamine and dimethylamine on the adsorption kinetics, different

Table 2 Langmuir parameters for methylamine and dimethylamine. Compounds

Kads (L mmol−1 )

Qmax (mg g−1 )

Qmax (molecule nm−2 )

Methylamine Dimethylamine

0.8 0.85

5.27 7.47

2 1.9

S. Helali et al. / Applied Catalysis A: General 402 (2011) 201–207

7

DMA

6

Qe(mg/g TiO2)

203

(a)

MA

5 4 3 2 1 0 0

1

2

3

4

5

Ce(mmol/L) 2.0

DMA Fig. 1. Photocatalytic set-up for the degradation of methylamine and dimethylamine.

(b)

MA

Qe =

Kads Qmax Ce 1 + Kads Ce

Qe is the adsorbed quantity of pollutant on the photocatalyst at the equilibrium (mg g−1 ), Kads is the adsorption constant (L mmol−1 ), Qmax is the maximum amount to be adsorbed (mg g−1 ), and Ce is the concentration of the compound at the adsorption equilibrium (mmol L−1 ). The values of the Langmuir parameters for MA and DMA are represented in Table 2.

Table 3 Langmuir and Langmuir–Hinshelwood parameters values under dark and UV conditions. Dark

UV −1

Qmax (mg g MA DMA

5.27 7.47

)

Kads (L mmol 0.80 0.85

−1

)

k (␮ mol−1 L−1 min−1 )

K (Lmmol)

26.34 12.05

1.5 1.1

40

1.0

OH covered (%)

solutions of MA and DMA were maintained in the dark under stirring during 30 min until equilibrium was reached. Two different methods were used to characterize the adsorption of methylamine and dimethylamine on the TiO2 surface. Fig. 2(a) represents the amounts (mg g−1 ) of MA and DMA adsorbed per gram of TiO2 while Fig. 2(b) gives the area density of the adsorbed pollutants (molecule nm−2 ). All results are represented as a function of the MA and DMA equilibrium concentration (Ce ). Fig. 2(a) shows that, in dark conditions, the amount of MA and DMA adsorbed on the TiO2 surface (Qe ) increases with the equilibrium concentration until it reaches a plateau. This amount seems to be more important for DMA due to its molecular weight higher than for MA. However, Fig. 2(b) shows that for different concentrations of MA or DMA, the area density (molecule nm−2 ) of adsorbed MA and DMA are almost the same leading to similar percentages of covered OH for different concentrations of MA and DMA. Covered OH represents almost 30% for MA and DMA in the plateau region (Fig. 2(c)) corresponding to 1.5 molecule nm−2 . This value is important compared to values, less than 1%, found for molecules containing aromatic cycles such as tryptophan, phenylalanine [31,32] and anisole-type compounds [33]. Like the majority of organic compounds [31], the adsorption isotherms of MA and DMA can be modelized by Langmuir model according to the following equation:

molecule/nm2

1.5

0.5

(c)

30 20 10 0 0

1 2 3 4 Ce (mmol/L)

5

0.0 0

1

2

3

4

5

Ce(mmol/L) Fig. 2. Isotherms of adsorption of MA and DMA in TiO2 Degussa P25 (1 g L−1 ).(a) Amounts of MA and DMA adsorbed per gram of TiO2 as a function of the equilibrium concentration Ce . (b) Area density of the adsorbed pollutants (molecule nm−2 ) as a function of the equilibrium concentration Ce . (c) covered OH as a function of the equilibrium concentration Ce .

3.2. Photocatalytic degradation of methylamine and dimethylamine To highlight the importance of the TiO2 photocatalyst under UV-light conditions for the photocatalytic degradation of the methylamine and dimethylamine pollutants, experiments of adsorption, photolysis and photocatalysis were carried out as shown in Fig. 3. Low adsorption values of MA and DMA are obtained in the presence of TiO2 suspension under dark condition. The direct photolysis of MA and DMA at  > 340 nm was negligible because their main absorbance occurs below 290 nm. The disappearance observed in the presence of TiO2 under UV conditions at this wavelength ( > 340 nm) is only due to the photocatalytic process. After 30 min of agitation in the dark, about 0.63 molecule nm−2 (3 × 1019 molecules g−1 ) of MA and DMA are adsorbed at equilibrium. Taking in to account that adsorption of MA or DMA occurs on OH• groups which represent about 5 OH• nm−2 [34], this amount corresponds to a coverage of 12.6% of OH• present on the TiO2 surface. The photocatalytic degradation of different initial concentrations of methylamine and dimethylamine were obtained upon illumination at  > 340 nm with a radiant flux ˚ of 6.5 mW cm−2 , in the presence of 1 g L−1 of TiO2 and under natural pH. The kinetics of disappearance at different MA and DMA concentrations is

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UV

0.6

30

MA 0.5

DMA

r0 (µmol/L/min)

24

C (mmol/L)

0.4 0.3 0.2 0.1

18

12

6

0 -30

0

30

60

90 120 150 180 210 240 270 300 330

photolysis MA and DMA photocatalysis MA

0

represented as a function of irradiation time in Fig. 4. For both MA and DMA, the photodegradation is initially rapid but gradually decreases. The time necessary for the maximum degradation of MA and DMA increases with the concentration of both aliphatic amines. The initial disappearance rates r0 are determined for all concentrations of methylamine and dimethylamine and are represented as a function of the concentration at the equilibrium (Ce ) (Fig. 5). Fig. 5 shows that the initial disappearance rates r0 for MA and DMA increase as a function the Ce concentration, before reaching

dark

UV-irradiation

5 mmol/L

5

C (mmol/L)

MA

2 mmol/L

0,5 mmol/L

3

0.25 mmol/L

2 1

0 30

60

90

120

150

180

210

240

270

300

t (min)

(b)

dark

UV-irradiation

5

5 mmol/L

DMA

C (mmol/L)

3 mmol/L

4

2 mmol/L 1 mmol/L

3

0.5 mmol/L 0.25 mmol/L

2

1

0 -30

0

30

60

90

120

3

4

5

Fig. 5. Evolution of the initial rate of disappearance of MA and DMA as a function of the concentration at the equilibrium (Ce ).

a plateau. The initial disappearance rate of methylamine is higher than for the dimethylamine case. The higher reactivity of MA could be due to the higher reactivity of the cationic form. Actually, due to the natural pH (Table 4) present in solution, the concentration of CH3 NH3 + is expected to be more important than the concentration of (CH3 )2 NH2 + suggesting that electrons would react with protonated amines as proposed by Möning et al. [35]. Until a concentration of about 0.5 mmol L−1 of MA or DMA, the initial disappearance rate r0 is proportional to the amine concentration according to the equation r0 = kKCe before reaching a plateau. The kK values for MA and DMA are respectively 0.032 min−1 and kK = 0.011 min−1 . Under these conditions, Langmuir–Hinshelwood model can be used according to the following equation: r0 =

1 mmol/L

0

2

3 mmol/L

4

-30

1

Ce(mmol/L)

Fig. 3. Adsorption, photolysis, photocatalysis of methylamine and dimethylamine (V = 30 mL, [TiO2 ] = 1 g L−1 , ˚ = 6.5 mW cm−2 ).

(a)

0

Adsorption MA and DMA photocatalysis DMA

150

180

210

240

270

300

t (min) Fig. 4. Disappearance of (a) MA and (b) DMA at different initial concentrations (V = 30 mL, [TiO2 ] = 1 g L−1 , ˚ = 6.5 mW cm−2 , natural pH).

kKCe 1 + KCe

r0 is the initial rate of MA and DMA disappearance (␮mol L−1 min−1 ), k is the rate constant (␮mol L−1 min−1 ), K is the adsorption constant under UV conditions (L mmol−1 ), and Ce is the concentration of the compounds at the adsorption equilibrium (mmol L−1 ). Rate and adsorption constants under UV conditions and Langmuir adsorption constants obtained in the dark for MA and DMA are summarized in Table 3. The rate constant appears higher for MA than for DMA. The fitting of the curves using the Langmuir–Hinshelwood model was in agreement with the result of Kachina et al. [28] who found that the MA conversion inversely varies with the concentration indicating that the process is not of the first order while fitting well with the Langmuir–Hinshelwood model. The TiO2 surface coverage rate () was calculated for both MA and DMA in the dark ( dark = Kads Ce /1 + Kads Ce ) and under UV condition ( uv = KCe /1 + KCe ) and was represented in Fig. 6. We note that the TiO2 surface coverage rate increases with the concentration of both amine compounds reaching 90% and 84% respectively for 5 mmol L−1 of MA and DMA concentrations. The surface coverage obtained in the dark is almost the same for MA and DMA which is not the case under UV condition where a higher coverage rate was observed for MA. For both amine compounds, coverage rates /  dark ), if considin the dark and UV condition are different ( uv = ering the difference between adsorption constant in the dark (Kads ) and under UV condition (K) (Table 3). This phenomenon has already been observed by us [30,31] and by several other authors [36,37]. Several hypotheses can be suggested to explain this behavior:

S. Helali et al. / Applied Catalysis A: General 402 (2011) 201–207

205

Table 4 pH variation after 300 min of photodegradation of MA and DMA and their molecular fraction. C (mmol L−1 )

pH (MA)

0.25 0.5 1 2 3

CH3 –NH3 +

pH (DMA)

CH3 –CH3 –NH2 +

t initial

t 300 min

t initial

t initial

t 300 min

t initial

9.06 9.15 9.64 9.9 10.42

6.4 6.67 6.9 7.5 7.11

83% 81% 73% 67% 55%

10.11 10.52 10.81 11.16 11.25

7.3 6.85 7.86 8.02 8.1

65% 55% 48% 39% 37%

• Modification of active sites under UV-irradiation considering the variation of the electronic proprieties of the TiO2 surface under UV where Ti4+ → Ti3+ and O2− → O•− • Recombination of active species (electron/hole or radicals) during the irradiation process causes heat generation (h + TiO2 → e− cb + h+ vb → heat) resulting in a modification of the thermodynamic equilibrium at the semiconductor surface. • The reaction takes place not only on the surface but also near the surface. • We can also notice that coverage rate of MA under UV is higher than for DMA under UV condition in agreement with the higher reactivity of MA.

3.4. Organic carbon mineralization Mineralization of organic carbon was followed by measuring the TOC evolution during the photocatalytic degradation of methylamine and dimethylamine as shown in Fig. 8. For different concentrations of MA and DMA, TOC decreases with irradiation time. The TOC evolution depend of the initial pollutants concentrations where less time was needed for a total mineralization at low concentration. For high concentration of MA and DMA, organic compounds persist in the solution suggesting the formation of products which are not easily mineralised. Below 1 mmol L−1 of an initial concentration of MA or DMA, TOC can be removed. This behavior is in agreement with the evolution of MA or DMA as a function of time (Fig. 4)

3.3. pH evolution during the MA and DMA degradation 3.5. Nitrogen mineralization Fig. 9 shows that during the photocatalytic degradation of MA and DMA, amines were predominantly converted to ammonium (NH4 + ) rather than nitrite (NO2 − ) or nitrate (NO3 − ). It is also interesting to note that nitrite ions are the precursors of the nitrate ion formation. Actually, the concentration of nitrite ions begins decreasing when nitrate ions appear due to the photooxidation of NO2 − to NO3 − [38]. MA was detected from the beginning of degra-

(a)

2.0 10 1.6 8 1.2

6

pH

C (mmol/L)

The pH of the solution depends on the initial concentration of MA and DMA, it increases with the concentration as reported in Table 4. For all concentrations the pH decreases during the degradation of MA and DMA. The values of pH obtained after 300 min of degradation are given in Table 4. In order to reach the total degradation, a concentration of 2 mmol L−1 of MA and DMA was chosen to be degraded until 15 h and 18 h respectively. The pH was measured for MA and DMA during this period of degradation as shown in Fig. 7. For a concentration of MA of 2 mmol L−1 the pH drop from 9.9 to 7.5 after 5 h of photodegradation. For 2 mmol L−1 of DMA the pH drop from 11.16 to 8.02. After this time (5 h) the pH of both MA and DMA stay almost constant until the end of degradation. This drop in the pH may be partially caused by the formation of nitrite, nitrate ions coming from the amine groups but also from organic acid formed as observed in Fig. 9.

MA pH

0.8

4

0.4

1.0

2

0.0 3

6

9

12

15

t (h)

(b)

2.0

0.6 10

1.6

θUV DMA θdark MA

0.2

θdark DMA

8 1.2 6

DMA pH

0.8

4 0.4

0.0 0

1

2

3

4

pH

θUV MA

0.4

C (mmol/L)

coverage rate (θ)

0 0

0.8

2

5

Ce (mmol/L)

0.0

0 0

Fig. 6. Evolution of the coverage rate () of the TiO2 surface in the dark and under UV conditions as a function of methylamine and dimethylamine equilibrium concentrations (Ce ).

3

6

t (h)

9

12

15

18

Fig. 7. pH evolution during the degradation of 2 mmol L−1 of (a) MA and (b) DMA.

206

S. Helali et al. / Applied Catalysis A: General 402 (2011) 201–207

2.0

36

N- theoritical

MA

C (mmol/L)

TOC (mg/L)

1.6 24

12

Nitrogen non-identified total-N identified

1.2 NH4+

MA

0.8

NO3-

0.4

HCOOH

0 0

3

6

9

12

15

18

3 mmol/L 0.5 mmol/L

NO2-

0.0

t (h) 2 mmol/L 0.25 mmol/L

0

1 mmol/L theoritical value

3

6

9

12

15

t (h) 2.0

DMA

N- theoritical

72

Nitrogen non-identified

C(mmol/L)

TOC(mg/L)

1.6

48

total-N identified

1.2

DMA

0.8

MA

24

NH4+

0.4 NO3NO2-

0

HCOOH

0.0

0

3

6

9

12

15

18

21

24

0

t (h) 3mmol/L 0.5mmol/L

2mmol/L 0.25mmol/L

3

6

9

12

15

18

t (h) 1mmol/L theoritical value

Fig. 9. Photocatalytic degradation of MA and DMA and the formation of products as a function of the irradiation time.

Fig. 8. Kinetics of TOC disappearance of five initial concentrations of MA and DMA.

dation when the % of DMA disappearance was lower than 15%. It was formed as an intermediate product in higher amount than NH4 + during the photocatalytic degradation of DMA confirming the findings of Choi [27]. The global transformation rate of MA was 86% after 3 h while complete degradation was achieved after 15 h. For DMA, the global transformation rate was respectively 79% and 94% after 3 h and 18 h. The mechanism of formation of NH4 + and NO3 − by photooxidation of nitrogen compounds has been discussed by different researchers [39–41]. NH4 + ions could be formed from the amine group of MA or DMA without any change of the oxidation state (-III) of nitrogen. NH4 + appears from the beginning of degradation, it is detectable after 1 min of irradiation when the percentage of MA degradation is about 10%. Then, NH4 + can be considered as a primary product during the degradation of MA contrary to NO2 − and NO3 − . It has been suggested by Nohara et al. [41] that nitrogen-containing compounds without hydroxylated nitrogen are predominantly converted to NH4 + and not to NO3 − , while Bui [42] shows the importance of the electronic density on the nitrogen fate. The nitrate ions (NO3 − ) can be obtained through the photocatalytic oxidation of nitrite (NO2 − ) and/or ammonium (NH4 + ) as confirmed by Low et al. [5] and Kim et al. [6]. However, this latter pathway is negligible under acid or neutral conditions while it is enhanced at higher pH [42]. As shown in Fig. 9, whatever the amine (MA or DMA), a deficit of nitrogen atoms is observed during the degradation process indicating the presence of non-identified intermediates. Kim et al. [6] reported that dimethylhydroxylamine was detected during the degradation of DMA or trimethylamine.

In our conditions of analysis, three intermediates are observed. The first one corresponds to formic acid while the two other ones are not identified. They do not correspond to formamide. Formic acid (HCOOH) only appears after 90 min of irradiation. This result can be explained considering the initial formation of alcohol such as HOCH2 NH2 then aldehyde HCONH2 before the acid formation. Finally, formic acid is oxidized on the UV-irradiated TiO2 surface [28]. The initial formation rate of NH4 + represents 33% of the initial rate of disappearance of methylamine suggesting that deamination is not the main initial step of photocatalytic degradation while the breaking of C–N bond is the major step in the degradation of DMA. Actually, the initial formation rate of MA represents 80% of the initial rate of disappearance of dimethylamine. Considering these results, two initial steps of MA degradation could be suggested: the formation of CH2 OH–NH2 or the formation of CH3 –NH–OH. 4. Conclusion The adsorption equilibrium and photocatalytic degradation of methylamine and dimethylamine under UV-A condition and in the presence of TiO2 Degussa P25 have been modelized by Langmuir and Langmuir–Hinshelwood models. The amount of MA or DMA adsorbed per nm2 of TiO2 in the dark are similar and corresponds to about 30% of covered OH groups when TiO2 saturation is achieved. The coverage rate of the TiO2 surface by MA and DMA is modified under UV condition. Under irradiation, the coverage rate of TiO2 by MA is more important than for DMA. The MA degradation rate is more important than for DMA in agreement with the highest coverage rate of MA under UV con-

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dition but also with the higher concentration of the cationic form of MA. The initial rate of TOC seems to be independent of the initial TOC and tends to increase with the pH decrease. At high TOC concentration, some difficult-to-degrade organic compounds are formed at the end of the degradation process. Ammonium and MA are respectively the main products initially formed during the degradation of MA and DMA. Concerning the degradation mechanism, the main initial step for DMA degradation is the breaking of the C–N bond while this reaction represents only 33% in the MA case. Acknowledgements S. Helali is very grateful to Islamic Development Bank (IDBARABIE SAOUDI) who supported this work. The authors wish to thank Mr Gilles Berhault for the English language revision. References [1] J.M. Herrmann, C. Duchamp, M. Karkmaz, B.T. Hoai, H. Lachheb, E. Puzenat, C. Guillard, J. Hazard. Mater. 146 (2007) 624–629. [2] P. Pichat, J. Disdier, C. Hoang-Van, D. Mas, G. Goutailler, C. Gaysse, Catal. Today 63 (2000) 363–369. [3] J. Zhao, X. Yang, Building Environ. 38 (2003) 645–654. [4] C. Guillard, T.-H. Bui, C. Felix, V. Moules, B. Lina, P. Lejeune, Comptes Rendus Chimie 11 (2008) 107–113. [5] G.K.C. Low, S.R. McEvoy, R.W. Matthews, Environ. Sci. Technol. 25 (1991) 460–467. [6] S. Kim, W. Choi, Environ. Sci. Technol. 36 (2002) 2019–2025. [7] S. Devipriya, S. Yesodharan, Sol. Energy Mater. Sol. Cells 86 (2005) 309–348. [8] J.M. Herrmann, C. Guillard, M. Arguello, A. Agüera, A. Tejedor, L. Piedra, A. Fernández-Alba, Catal. Today 54 (1999) 353–367. [9] J.M. Herrmann, C. Guillard, C.R. Acad. Sci., Ser. IIc: Chem. 3 (2000) 417–422. [10] A. Marinas, C. Guillard, J.M. Marinas, A. Fernández-Alba, A. Aguëra, J.M. Herrmann, Appl. Catal. B: Environ. 34 (2001) 241–252. [11] A. Topalov, D. Molnár-Gábor, B. Abramovic, S. Korom, D. Pericin, J. Photochem. Photobiol A: Chem. 160 (2003) 195–201. [12] M. Karkmaz, E. Puzenat, C. Guillard, J.M. Herrmann, Appl. Catal. B: Environ. 51 (2004) 183–194. [13] K. Sahel, N. Perol, F. Dappozze, M. Bouhent, Z. Derriche, C. Guillard, J. Photochem. Photobiol A: Chem. 212 (2010) 107–112.

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