Alternative sources of hydrogen for hydrodechlorination of chlorinated organic compounds in water on Pd catalysts

Applied Catalysis A: General 271 (2004) 119–128

Alternative sources of hydrogen for hydrodechlorination of chlorinated organic compounds in water on Pd catalysts Frank-Dieter Kopinke∗,1 , Katrin Mackenzie, Robert Koehler, Anett Georgi Department of Environmental Technology, Centre for Environmental Research Leipzig-Halle, 04318 Leipzig, Germany Received in revised form 28 January 2004; accepted 1 February 2004 Available online 24 June 2004

Abstract Formic acid, isopropanol and hydrazine were investigated as reductants for the Pd-catalyzed hydrodechlorination of chlorobenzene in water at ambient temperature. The intention was to find alternatives to molecular hydrogen with high water solubilities. Formic acid was found to be as reactive as H2 under acidic and neutral conditions, but less reactive under alkaline conditions. The observed kinetics imply two pH-controlled reaction mechanisms (possibly H-atom and hydride transfer). H-consumers, such as chlorinated compounds, strongly stimulate the decomposition of formic acid. The half-life of 5 mg L−1 chlorobenzene in the presence of 1 mg L−1 Pd is about 2 min under optimal reaction conditions. Rh was found to be inactive in the formic acid driven hydrodechlorination. Isopropanol is less reactive by about five orders of magnitude than H2 . Hydrazine is effective as a H-donor for the hydrodechlorination under alkaline conditions. However, the reaction is slower than with H2 by a factor of 30. From the technical and economic point of view, formic acid is a promising substitute for H2 . © 2004 Elsevier B.V. All rights reserved. Keywords: Hydrodechlorination; Palladium catalyst; Rhodium catalyst; Hydrazine; Formic acid; Chlorobenzene

1. Introduction Chlorinated organic compounds (COCs) are among the most widely distributed pollutants in wastewaters and contaminated groundwaters. Their destruction or conversion to environmentally harmless compounds is a significant challenge. From the chemical point of view, the choice lies between oxidation and reduction of COCs. The advantage of reductive transformations such as hydrodechlorination is a high degree of selectivity towards halogenated pollutants, whereas oxidative transformations are usually less selective. Strong oxidants such as OH-radicals attack a broad range of water constituents including non-pollutants (e.g. natural DOM, bicarbonate, etc.). One of the promising technologies for reductive dechlorination of COCs in contaminated aquifers is the application of zero-valent iron, which reduces a variety of COCs accord-

∗ Corresponding author. Tel.: +49 341 235 3264; fax: +49 341 235 2492. E-mail address: [email protected] (F.-D. Kopinke). 1 UFZ-Centre for Environmental Research, Department of Environmental Technology, Permoser Str. 15, 04318 Leipzig, Germany.

0926-860X/$ – see front matter © 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.apcata.2004.02.052

ing to R–Cl+Fe0 +H2 O → R–H+Fe2+ +Cl− +OH− [1–4]. Unfortunately, this reaction fails for chlorinated aromatic compounds and some environmentally relevant aliphatic pollutants, such as dichloromethane and 1,2-dichloroethane. However, in many cases the application of Pd catalysts can overcome this deficiency [5–17]. The mechanisms of the Pd-catalyzed and the Fe-driven dechlorinations are different: Pd activates molecular hydrogen (Pd · · · H∗ ) which attacks the C–Cl bond or any double bond in the substrate molecule, whereas Fe0 drives an electron transfer from the surface to the C–Cl-bond. Pd is very active as a catalyst in hydrodehalogenation reactions, but it suffers from its sensitivity to catalyst poisons, such as sulphide and sulphur organic compounds [11,12,18–20]. The preferred reductant in the Pd-catalyzed hydrodechlorination is molecular hydrogen (H2 ). It has a water solubility 15 ◦ C of SH = 0.84 mM at pH2 = 100 kPa [21]. This may be 2 not sufficient for treatment of highly contaminated waste or groundwaters. From the stoichiometry of the hydrodechlorination reaction it follows that only up to 28.6 mg L−1 of trichloroethene can be reduced with dissolved hydrogen (C2 HCl3 + 4H2 → C2 H6 + 3HCl). In practice, however, much higher COC concentrations may occur [3,17]. The

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repeated dosage or even in situ providing of gaseous hydrogen into contaminated aquifers is a difficult procedure. Furthermore, hydrogen is a very digestible electron donor for microorganisms. The microbial sulphate reduction produces sulphide which is a strong poison for Pd. Therefore, the study of alternative reductants is a worthwhile task. One promising alternative for providing the Pd with hydrogen generated in situ is to use palladised iron [22–24] in the form of nano-particles [4]. The hydrogen-generating reaction in this system is the corrosion of iron according to Fe0 + 2H2 O → Fe2+ + H2 + 2OH− . However, the iron corrosion may be the rate-limiting step in the reaction sequence, and the corrosion rate usually decreases over time. Nevertheless, iron is an environmentally friendly hydrogen source with a relatively long lifetime, depending on the water chemistry (pH and EH values, salinity) in the order of magnitude of 102 . . . 103 days. This may be an advantage for long-term in situ application in contaminated aquifers, where continuous hydrogen supply is desirable. It is not, however, an appropriate solution for fast reactions in relatively short time intervals. In the present study, three infinitely water-soluble reductants were investigated: formic acid, hydrazine, and isopropanol. Their water solubility makes it feasible to inject the pure substances or their highly concentrated aqueous solutions. From the engineering point of view this is a clear advantage compared to gaseous reactants. The question to be answered in this study is whether they are as effective as molecular hydrogen from the chemical point of view. The present study is focussed on feasibility issues. Mechanism and kinetics studies of the chemical reactions concerned were carried out to the extent necessary for answering these questions.

2. Methods Methanol, chlorobenzene, formic acid (>98%) and isopropanol were purchased in 99+ grade from Merck (Germany), hydrazine hydrate (98%) was purchased from Lancaster (UK). All reagents were used as received. The catalyst pellets (0.5 wt.% Pd on ␥-Al2 O3 , G-133D, egg-shell impregnated, from Commercia, Suedchemie, Germany) were crushed, sieved (fraction < 63 ␮m), and pre-reduced with NaBH4 in methanol, washed with deionised water and dried prior to use. The batch experiments were carried out in 250 mL screw-cap bottles equipped with mininert® valves. Between 2 and 500 mg of the pulverised catalyst (depending on the expected reaction rate) were added to 200 mL of deionised water. The pH value of the suspension was adjusted to the desired value by adding 1N H2 SO4 or NaOH. The suspension was then purged with the appropriate gas (Ar, N2 , He or H2 ) for about 30 min. After purging, the headspace (50 mL) was free of air. Chlorobenzene (c0 = 0.05 to 0.5 mM, as methanolic solution) and the liquid reductants

(e.g. 250 mg L−1 HCOOH, 500 mg L−1 N2 H4 , 1000 mg L−1 C3 H7 OH) were then injected through the mininert® valve. After adding the contaminants, the bottle was vigorously shaken manually for 1 min and then continuously on a mechanical shaker (100 rpm). The reaction temperature was 23 ± 2 ◦ C. The progress of the reaction was monitored by GC analysis of headspace samples and, at the end of the reaction period, by pH measurement and argentometric titration of the chloride formed. The gas samples (50 ␮L) were analyzed via GC-MS (30 m capillary column DB1, QP 5000, Shimadzu Corp.) for benzene and chlorobenzene, as well as via GC-TCD (2 m packed columns, molecular sieve and Tenax® , GC14B, Shimadzu Corp.) for the gases hydrogen, oxygen, nitrogen and carbon oxides. The analyses were calibrated with external standards.

3. Results and discussion The reference reaction for testing the three liquid reductants was the hydrodechlorination of chlorobenzene with molecular hydrogen (pure H2 in the headspace volume over the catalyst suspension). It yields a fast and complete conPd Ph–H + version to benzene according to Ph–Cl + H2 − → − + H + Cl . The reaction follows a first order kinetics (0.5–0.01 mM Ph–Cl), irrespective of the decreasing pH value and the increasing chloride concentration during the course of the reaction. This makes it possible to characterise the specific catalyst activity [16] by only one parameter: APd (L g−1 min−1 ) according to Eq. (1) or, if the half-life τ 1/2 is not available from the experimental data, from Eq. (2): APd =

Vwater mPd τ1/2

(1)

APd =

Vwater ln(ct1 /ct2 ) mPd (t2 − t1 ) ln 2

(2)

where Vwater is the volume of the aqueous suspension, mPd the mass of Pd in the applied catalyst, τ 1/2 the half-life of the chlorobenzene, t1 and t2 are two arbitrarily chosen sampling times, ct 1 and ct 2 the corresponding chlorobenzene concentrations. For a better illustration: a specific Pd activity APd = 1 L g−1 min−1 reflects that the contaminant concentration (e.g. chlorobenzene) can be reduced with 1 g Pd in 1 L water within 7 min by 99.2% (1 − 0.57 = 0.992). Therefore, the specific Pd activity is a description of the catalyst efficiency. Usually, the conversion of chlorobenzene and the formation of benzene follow identical kinetics (cbenzene + cchlorobenzene = c0,chlorobenzene ). Therefore, catalyst activities may be calculated as desired, either from the chlorobenzene conversion or from the benzene formation rates. The second approach is beneficial for low degrees of conversion. Typical catalyst activities found were in the range APd = 100 to 200 L g−1 min−1 , depending on the catalyst sample and preparation. This means that 50 mg L−1

F.-D. Kopinke et al. / Applied Catalysis A: General 271 (2004) 119–128

H* HCOOH + Pd

Pd…H* + CO2 R-Cl Pd + R-H + HCl

3.1. Formic acid O=C-H

O=C

-

=

O + R-Cl

O + H-R + Cl

-

-

The Pd-catalyzed hydrodechlorination of chlorobenzene in the presence of formic acid follows approximately a first order kinetics with respect to chlorobenzene up to ≥95% conversion under acidic conditions (c0,PhCl = 0.05 mM, c0,HCOOH = 5.5 mM, c0,catalyst = 250 mg L−1 , pH0 ≈ 3.2). Benzene and chloride (100 ± 5% yield) were the only detectable reaction products. The catalyst activity was found to be in the range of APd = 450 ± 150 L g−1 min−1 (average of three runs). This is even higher than with molecular hydrogen. For comparison of the two reductants, it must be taken into account that the molar concentration of HCOOH was six times higher than that of dissolved H2 . Although the combination of HCOOH with Pd is well known from the literature as a powerful reductant [25–37], it is the first time that the appropriateness of HCOOH as a H-donor in aqueous solution for this reaction has been demonstrated and quantitatively compared with H2 . Anwer et al. [38] described the application of ammonium formate for the Pd-catalyzed dechlorination of some chloroaromatics. However, the reaction was carried out with high reactant concentrations in organic solvents. Their method was designed to safely dispose laboratory wastes rather than to clean contaminated groundwater. During the course of hydrodechlorination runs, only traces of H2 were observed in the headspace. Obviously, this is due to the rapid consumption of molecular hydrogen in the hydrodechlorination reaction. Gaseous H2 appears only at very low residual concentrations of chlorobenzene (<5 ␮M). The initial hydrodechlorination rate (in mol benzene per s) is higher than the HCOOH decomposition rate in the absence of chlorobenzene (in mol CO2 per s) by 1–2 orders of magnitude, in contrast to the ratio of the reactant concentrations (e.g. 0.45 mM chlorobenzene versus 5.5 mM HCOOH). Obviously, the hydrodechlorination stimulates the formic acid decomposition. We see two mechanistic explanations for this finding (cf. Fig. 1): (i) hydrogen saturation of the Pd clusters inhibits the formic acid decomposition. This is consistent with the desorption of H2 from the Pd surface as the rate-determining step. The rapid consumption of chemisorbed hydrogen by the hydrodechlorination facilitates further decomposition of HCOOH. (ii) The hydrodechlorination does not depend on chemisorbed hydrogen, but can utilise directly the hydride hydrogen from chemisorbed formate. If the formate decomposition is the rate-determining step, this hydride transfer increases the overall decomposition rate. The slow auto-decomposition of formic acid may be an advantage. In principle, it offers the potential for temporallyand spatially-controlled generation of H-equivalents: their rate of generation is controlled by the demand.

…

of catalyst (250 ␮g L−1 Pd) results in a half-life of about 20–40 min, which can be measured conveniently with the applied methods.

121

Pd

Pd

Fig. 1. Possible mechanisms of Pd-catalyzed hydrodechlorination in the presence of formic acid: (A) radical mechanism; (B) hydride mechanism.

In order to be technically feasible, the reaction system has to fulfil a number of requirements. It is known from the literature that in the catalytic nitrate reduction, various anions (in particular, chloride) compete with formate for certain Pd surface sites, giving rise to an inhibition of the reaction [26]. To examine whether chloride affects the hydrodechlorination reaction in a similar manner [13,14,39–43], sodium chloride (500 mg L−1 Cl− ) was added to the catalyst suspension. This caused a prolongation of the induction period before the hydrodechlorination started (e.g. from 10 to 40 min under neutral conditions, see below), but only a minor decrease of the reaction rate (by less than a factor of 2). The results of Hähnlein [26] for nitrate reduction, i.e. an activity loss of a bimetallic Pd/Sn catalyst by a factor of 8 due to the addition of 50 mg L−1 Cl− , were not confirmed in the present study for the hydrodechlorination reaction. Low chloride concentrations (≤20 mg L−1 ) are tolerated by the catalyst, as can be seen from the first order kinetics with a final chloride concentration of 0.45 mM (16 mg L−1 ). Fig. 2 presents concentration curves from a multi-step batch experiment, which illustrates some characteristic findings in this complex reaction system. First, the decomposition kinetics of formic acid in the absence of chlorobenzene was monitored. The concave curvature of the H2 formation indicates the auto-inhibition of the decomposition reaction (about 2% HCOOH conversion after 1.5 h). After 1.5 h chlorobenzene (c0,PhCl = 0.5 mM) was injected into the reaction mixture. The hydrodechlorination started instantaneously (APd ≈ 150 L g−1 min−1 ): chlorobenzene was converted to benzene. The molecular hydrogen which had accumulated in the headspace (up to 1 vol.%) was rapidly consumed, after which the hydrodechlorination rate decreased. Finally, the rate dropped towards zero (only 98% chlorobenzene conversion after 100 h), although gaseous hydrogen (0.2 kPa) was still present. After purging the catalyst suspension with nitrogen (for 30 min) and injecting fresh chlorobenzene (0.5 mM), the hydrodechlorination restarted. However, the catalyst activity was lower by a factor of 30 compared with the initial rate. The various inhibition and re-activation phenomena are not yet completely understood. One type of catalyst deacti-

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ci c0, PhCl

cH2, Headspace [vol.-%] 1.0

1

Ph-Cl

0.5

H2

0.5

Ph-Cl H2

Ph-H

Ph-H

≈

5

t [h]

+ 0.5 mM Ph-Cl

10

80

100

120

+ 0.5 mM Ph-Cl Purging with N2

+ 5 mM HCOOH

Fig. 2. Decomposition of formic acid and hydrodechlorination of chlorobenzene on Pd/Al2 O3 (ccatalyst = 10 mg L−1 , c0,HCOOH = 5.5 mM, c0,PhCl = 0.5 mM, pH ≈ 3.2, 250 mL headspace volume per liter water; dashed lines indicate extrapolated, not measured values).

vation seems to be caused by volatile compounds, because it can be reversed by extensive purging the suspension with an inert gas. CO would be a possible candidate for catalyst deactivation. However, control experiments with traces of added CO were not successful in removing it from the deactivated Pd surface simply by nitrogen purging (see below). CO was not detected in the headspace of any batch experiment (detection limit 50 ppmv corresponding to 0.01% of the applied HCOOH). The lack of detection does not, however, prove the absence of CO, because it can be expected to be strongly adsorbed to the Pd surface. The resulting catalyst deactivation would probably prevent its further generation. From the mechanistic point of view, it is interesting to examine the decomposition of formic acid in the absence of H-consuming compounds. Taking into account the pH-dependent speciation of HCOOH, at least four different decomposition pathways are possible: the dehydration and the dehydrogenation of formic acid and formate (Fig. 3).

B1 CO(g) + H2O(l)

A1 HCOOH(diss.)

Dehydration CO(g) + OH-

Dehydrogenation HCOO-

B2

CO2(g) + H2(g)

H2O

HCO3- + H2(g)

A2

Fig. 3. Reaction scheme for the decomposition of formic acid in diluted aqueous solution at various pH values.

For evaluation of a broad pH window, the protonation equilibria of carbonic acid (pKA1,true = 3.3, pKA1,apparent = 6.5) must also be taken into account. According to the available thermodynamic data (Table 1), three of the four pathways have strong driving forces. Only the dehydration of formate in alkaline solution (reaction B2) is not thermodynamically favourable. Although the two dehydrogenation reactions are strongly favoured from the thermodynamic point of view, the dehydration may in practice become significant, because trace amounts of CO are sufficient to cause an irreversible poisoning of the catalyst. This was demonstrated in a control experiment by adding small amounts of CO (0.5 ␮mol CO corresponding to 0.005% of the present HCOOH, nCO :nPd ≤ 0.2 mol/mol) to a batch experiment under standard conditions. Both reactions, hydrodechlorination and formic acid decomposition, stopped instantaneously. In a number of control experiments it could be ruled out that CO2 is reduced to CO under the applied conditions (pCO2 = pH2 = 50 kPa). The presence of carbon dioxide did not affect the measured hydrodechlorination rates. Beside the thermodynamic driving forces, the protonation equilibrium between formic acid and formate (HCOOH ↔ HCOO− + H+ , pKA = 3.75) can be decisive for the overall reaction rates. This is to be expected if the two species have different decomposition rates at the Pd surface, and is the kinetic aspect of the reaction mechanism. Table 2 gives some Pd activities for the decomposition of formic acid depending on the pH value of the solution, calculated from the formation rates of H2 and CO2 . The data indicate that formic acid is much more reactive than

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123

Table 1 Thermodynamic data for the decomposition of formic acid and formate in aqueous solution (data from [21] and calculated according to the scheme in Fig. 3) Species, reaction pathways

F G◦ (kJ mol−1 )

F H◦ (kJ mol−1 )

H2 O(l) CO2(g) CO(g) HCOOH(l) Pathway A1a Pathway B1a Pathway A2 Pathway B2

−237.1 −394.4 −137.2 −361.4

−285.8 −393.5 −110.5 −424.7

R G◦ (kJ mol−1 )

R H◦ (kJ mol−1 )

Log Kdecomposition

−28.9 −8.8 −25.9 50.0

31.2 28.4 n.a.b n.a.

5.0 1.5 4.5 −8.7

a The free enthalpy of dissolution of HCOOH in water (HCOOH (l) → HCOOH(diss) ) was estimated from the vapour pressure of the pure HCOOH (p0 298 K = 5.75 kPa) and its Henry’s law constant (log KH [−] = −5.1), according to diss G = RT ln γW = RT ln (KH RT/VW p0 ) = −4.1 kJ mol−1 . b Not available.

Table 2 Dependence of the starting activities of a Pd-catalyst for the decomposition of formic acid and formate on the pH value of the suspension (ccatalyst = 3 g L−1 , corresponding to 15 mg L−1 Pd, c0,HCOOH = 5.5 mM) Starting pH value of the catalyst suspension

pH = 3.0

pH = 5.5

pH = 8 . . . 9

cHCOOH /cHCOO− = 10(pK a−pH) A0,Pd (L g−1 min−1 )

5.6 1.3

0.018 0.017

≤6 × 10−5 0.012

formate in the decomposition reaction. Under acidic conditions, 90 ± 10% of the theoretical yields of the two products were detected after 20 h, i.e. the reaction was almost complete. Under alkaline conditions, the decomposition was slow but stable over more than 50 h, with no indication of catalyst deactivation. These findings show that formate probably has a significant decomposition rate, independent of its protonated counterpart (kHCOO −:kHCOOH ≈ 0.01 at pH = 8 . . . 9). Although hydrodechlorination with formic acid succeeds rapidly and completely under acidic conditions, a pH value close to neutral conditions is more convenient for practical 10.6

pH

9.8

applications. Fig. 4 shows results of an experiment with a stepwise decrease of the pH value. The hydrodechlorination proceeds very slowly under strongly alkaline conditions (APd = 0.02 L g−1 min−1 at pH = 10.6). The reaction rate increased with decreasing pH value (APd = 15 L g−1 min−1 at pH = 8.2). The rate coefficients remained approximately constant through each of the 90 min measuring periods. After adjusting the pH value from 8.2 to 6.9 (injection of 0.1N H2 SO4 ) the kinetics of the reaction changed dramatically. The formal rate coefficient increased in a way similar to an auto-catalytic reaction from APd ≈ 60 L g−1 min−1 after 1 min up to ≥350 L g−1 min−1 after 15 min. During this reaction period the pH value of the suspension increased slightly from 6.9 to 7.3. Thus, the acceleration cannot be attributed to a further acidification of the reaction mixture by the formed HCl, as could be supposed. The phenomenon of an induction period followed by an auto-catalytic acceleration of the hydrodechlorination was observed in a number of similar experiments, when the pH value of the catalyst suspension was adjusted close to neutral conditions (7.0 ± 0.3) before 9.0

8.2

6.9

00

-2-2

≈

-3-3

APd -4-4

0.02

1.0 … 1.5

11

15

[L x g-1 x min-1]

-5-5 0

1

50

≈

ln (C/C0)

-1-1

2

3

4

5

6

350 7

t [h] Fig. 4. Dependence of the hydrodechlorination of chlorobenzene with formic acid on Pd/Al2 O3 on the pH value of the reaction mixture (ccatalyst = 250 mg L−1 , c0,HCOOH = 5.5 mM, c0,PhCl = 0.05 mM).

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injecting the formate stock solution (pH = 7.0). Only the length of the induction period (≤20 min) and the degree of the acceleration varied slightly. The reaction was complete (>99.9% conversion) within less than 60 min in all cases (c0,PhCl = 0.05 mM, cPd = 1.25 mg L−1 ). An alternative interpretation of the observed concentration profiles might be the combination of an induction period with a zero-order reaction kinetics. This would affect an increase of apparent first order rate coefficients (‘auto-acceleration’) during the progress of the reaction, as depicted in Fig. 4. Although it is not the intention of the present study to elucidate reaction mechanisms in detail, the findings may be tentatively interpreted on the basis of the assumptions presented in Fig. 1: hydrodechlorination can proceed via two different reaction mechanisms. Under alkaline conditions, the relatively slow hydride mechanism dominates, whereas under neutral and acidic conditions the fast radical mechanism via H-atoms is favoured. This interpretation is supported by the fact that the catalyst activities are in the same order of magnitude for H2 under neutral conditions as for HCOOH under acidic conditions. The induction period might be due to some kind of catalyst formation, including the pH-dependent surface charge of the carrier material ␥-Al2 O3 . The sensitivity of Pd catalysts to a number of catalyst poisons is well known [11,12,17–20]. One concept to protect the metal clusters is to embed them in polymer membranes, which are permeable for the reactants but impermeable for undesired ionic compounds [12,15,17]. In a number of batch experiments we examined the applicability of formic acid as H-donor for this type of membrane-based catalyst (0.4 wt.% of finely dispersed Pd in a non-porous poly(dimethylsiloxane) membrane). The hydrodechlorination activity was low in acidic solution (APd = 0.05 L g−1 min−1 at pH = 3.0) and even lower in alkaline solution (APd = 0.005 L g−1 min−1 at pH = 10.0). The same membrane-based catalyst, however, is fairly active (APd = 7 L g−1 min−1 ) when H2 is provided, regardless of the pH value of the solution. The observed order of reaction rates (H2  HCOOH > HCOO− ) corresponds with the chemical reaction rates and the expected diffusion rates of the various reductants through the polymer membrane. Clearly, the membrane-based catalyst cannot compete with the very high reaction rates of the suspended carrier catalyst, due to inevitable transport resistances. After palladium, rhodium is the most active metal for hydrogenolysis of C–Cl bonds [14,40]. A significant difference in the reactivity of the two metals lies in the ability of rhodium to catalyze the hydrogenation of aromatics as a parallel reaction to the hydrodechlorination under mild reaction conditions. Thus, Rh produces a mixture of benzene and cyclohexane from chlorobenzene, whereas only traces of cyclohexane are obtained with Pd as catalyst. In the present study we examined a Rh/␥-Al2 O3 -catalyst (0.5 wt.% Rh, from Sigma-Aldrich, Germany) for the HCOOH-driven reduction of chlorobenzene under the same experimental conditions as the Pd-catalyst. Rh was found to be inactive for

the hydrodechlorination in the presence of HCOOH as well as for the HCOOH decomposition (pH ≈ 3). The same catalyst was active in the H2 -driven reduction with ARh = 25 L g−1 min−1 , which is comparable to Pd-catalysts in the virgin form (not activated with NaBH4 ). After contact with HCOOH, the Rh-catalyst was completely deactivated even for hydrodechlorination in the presence of H2 . This may possibly be due to CO poisoning. 3.2. Isopropanol Isopropanol was selected as being one of the most reactive representatives of the group of alcohols. The basic reaction is C3 H7 OH + Ph–Cl → C3 H6 O + Ph–H + HCl. The hydrodechlorination of chlorobenzene proceeds in the presence of isopropanol, but with a very low reaction rate. It corresponds to specific catalyst activities of APd = 0.004 L g−1 min−1 at the beginning and 0.001 L g−1 min−1 at the end of a 2 weeks reaction period. (c0,PhCl = 0.45 mM, cisopropanol = 16.6 mM, ccatalyst = 2.5 g L−1 , pH0 = 7). After the 2 weeks, only 15% of the chlorobenzene had been converted to benzene. Neither an increase of the isopropanol concentration (factor of 10) nor a variation of the pH value of the suspension (from 4 to 10) gave rise to an increase of the hydrodechlorination rate. It was speculated that one reason for the low hydrodechlorination rates might be the competition between chlorobenzene and oxygen, traces of which could possibly penetrate into the reaction vessels during long reaction periods. Therefore, the reaction bottles (including the mininert® valves) were handled in plastic containers under helium-purged water. In this way, the leakage rate of the reaction system could be minimized (100 ␮L of air per week). From the O2 :N2 ratio (<0.05) in the helium headspace over the catalyst suspension it could be deduced that trace amounts of O2 entering the apparatus were completely reduced. However, the system of Pd plus isopropanol was not sufficiently active to reduce larger amounts of oxygen (no significant conversion of a 5 mL O2 pulse within 3 days). After 15 days, 20 mL of H2 were injected into the reaction mixture in order to test the remaining catalyst activity. Chlorobenzene and oxygen were both reduced at high rates (APd = 30 and 350 L g−1 min−1 , respectively). This proves that the catalyst was not deactivated, and that isopropanol is not an appropriate reductant under the applied mild conditions. The reason for this is of a kinetic (reaction mechanism) rather than of a thermodynamic nature (i−C3 H7 OH+ 0.5O2 ↔ C3 H6 O + H2 O, R G◦ = −212 kJ mol−1 ). 3.3. Hydrazine The driving force for the decomposition of pure hydrazine into its elements according to N2 H4 → N2 + 2H2 is high (R G◦ = −150 kJ mol−1 ). Nevertheless, we did not observe any spontaneous decomposition of hydrazine in aqueous catalyst suspensions in the absence of oxidants:

F.-D. Kopinke et al. / Applied Catalysis A: General 271 (2004) 119–128

R-Cl Pd…H*

Pd + ½ H2

R-H + HCl fast slow

?

+ R-Cl

- N2

Pd…(NH2)2

Pd + (NH2)2

fast

+ O2 N2 + 2 H2

2 H2O + N2 + [H2]Traces

Fig. 5. Schematic presentation of Pd-catalyzed reactions of hydrazine with chlorinated compounds and oxygen.

molecular hydrogen is not idly released and dissipated. It is remarkable that the injection of oxidants such as COCs or O2 initiates a momentary release of small amounts of gaseous hydrogen (cf. Fig. 7b), which is then rapidly consumed. Apparently, the desorption of hydrogen is possible only under non-stationary conditions, i.e. when the concentration of H-consumers increases steeply. Most of the experimental findings are integrated in the reaction scheme in Fig. 5. When hydrazine (500 mg L−1 , 15.6 mM) was added to the catalyst suspension, its pH value increased up to about 10. Hydrazine is a weak base (N2 H4 +2H+ ↔ N2 H5 + +H+ ↔ N2 H6 2+ with pKA1 = 8.11 and pKA2 = −0.88). Under alkaline conditions it is mainly present as neutral species. The hydrodechlorination of chlorobenzene follows a kinetics close to a zero-order with a catalyst activity of APd = 5±1 L g−1 min−1 at c0,PhCl = 0.45 mM. Although this value

125

is lower with hydrazine than with hydrogen and formic acid, it has a significantly different meaning due to the different reaction orders. This becomes evident especially for low substrate concentrations, e.g. at high degrees of conversion. For achieving 99% of dechlorination, 6.6 half-lives are necessary with H2 or HCOOH (first order kinetics), whereas two half-lives are sufficient with N2 H4 (zero-order kinetics). This may partly compensate for the lower catalyst activity. Typical concentration curves are presented in Fig. 6. Although the disappearance of chlorobenzene and the appearance of benzene in the headspace over the catalyst suspension are not fully synchronous, the characteristic reaction kinetics becomes clearly evident. In mechanistic terms, a reaction order close to zero implies either that chlorobenzene is not involved in the rate-limiting step (possibly the hydrazine decomposition) or that the active centres are saturated with chlorobenzene down to very low concentrations. The catalyst activity remains stable for at least 24 h and several hydrodechlorination runs (five successive injections of chlorobenzene with c0,PhCl = 0.45 mM), but tends to decrease slowly under reaction conditions (by a factor of about 3 after 7 days). Chloride up to 500 mg L−1 does not affect the catalyst activity. Oxygen is reduced by about one order of magnitude faster than chlorobenzene under identical reaction conditions. However, the abundant presence of O2 does not appear to affect the hydrodechlorination rate (see Fig. 7a). After 100 min and 60% chlorobenzene conversion, a pulse of oxygen (5 mL) was added to a running experiment. It can be clearly seen that the two hydrogen consumers, O2

11

Benzene

Ci/C0, PhCl

0,75 0.75

0,5 0.5

Chlorobenzene 0,25 0.25

00 0

0.5

1

1.5

2

2.5

3

t [h] Fig. 6. Hydrodechlorination of chlorobenzene with hydrazine (ccatalyst = 500 mg L−1 , c0,N2 H4 = 15.6 mM, c0,PhCl = 0.45 mM, pH0 = 9.3, APd = 4.5 L g−1 min−1 ).

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11

Benzene

Ci/C0, PhCl

0,75 0.75

0,5 0.5

- - - Injection of 5 ml O2

0,25 0.25

Chlorobenzene

00

0

5

10

15

20

t [h]

(a)

Injection of 5 mL O2 16

N2

C i [Vol.-%]

12

8 H2 (x 200) 4 O2 0 0,0 0

0,5 0.5

1,0 1

(b)

1,5 1.5

2,0 2

2,5 2.5

3,0 3

t [h]

Fig. 7. Hydrodechlorination of chlorobenzene with hydrazine: effect of oxygen (ccatalyst = 500 mg L−1 , c0,N2 H4 = 15.6 mM, c0,PhCl = 0.45 mM, pH0 = 9.3, APd = 5 L g−1 min−1 addition of 5 mL O2 after 100 min, p0,O2 = 10 kPa).

and chlorobenzene, do not compete for the same hydrogen pools. Fig. 7b illustrates the temporary appearance of traces of H2 in response to the injection of O2 . Fig. 8 demonstrates the different reactivities of N2 H4 and H2 for the Pd-catalyzed hydrodechlorination. In neutral and acidic solutions, hydrazine is present in its protonated form as hydrazonium ion. This is much less reactive as a H-donor in Pd-catalyzed reductions than the non-protonated hydrazine. Table 3 lists dependence of the reaction rates for the reduction of chlorobenzene and oxygen on the pH value of the reaction mixture. The data are arranged in chronological order, showing the course of a long-term experiment with the same catalyst sample. After 6 h reaction time, 20 mL of gaseous hydrogen was injected into a running experiment. The catalyst activity rose immediately by a factor of 50 from 3 to 150 L g−1 min−1 . This is a typical value for the fresh catalyst in the presence of H2 . Thus, hydrazine did not cause any short-term catalyst deactivation in alkaline suspension. After 22 h the suspension was purged with Ar (removal of hydrocarbons and H2 )

Table 3 Dependence of the change of the Pd-catalyst activity for the reduction of chlorobenzene and oxygen with hydrazine on the pH value of the reaction mixture in a long-term experiment (duration of the experiment: 20 days, ccatalyst = 500 mg L−1 , c0,N2 H4 = 15 mM, c0,PhCl = 0.45 mM) pH value

10.0 3.0 4.0 3.0 7.8 3.5 7.0 3.0 2.9 + H2 (40 kPa) 8.1 + H2 (20 kPa)

APd (L g−1 min−1 ) for the reactions N2 H4 + 2Ph–Cl → 2Ph–H + 2HCl + N2

N2 H4 + 2O2 → 2H2 O + N2

5.0 0.01 0.1 <0.001 n.d. 0.07 0.005 <0.001 0.03 (· · · 0.01 after 7 days) 0.055

80 n.d. 12.5 1.5 7.5 0.5 3.0 0.2 2.5 (· · · <0.1 after 7 days) 10

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127

Injection of 20 mL H2 1

Chlorobenzene

1

Stripping with Ar

Ci/C0, PhCl

0.7575

0.5 5

0.25

Benzene

0 0

5

10

15

20

25

30

t [h]

Fig. 8. Hydrodechlorination of chlorobenzene with hydrazine: effect of hydrogen (ccatalyst = 250 mg L−1 , c0,N2 H4 = 31 mM, c0,PhCl = 0.45 mM, pH0 = 9.3, APd = 3 L g−1 min−1 , addition of 20 mL H2 after 6 h, c0,H2 ≈ 0.3 mM from p0,H2 = 40 kPa).

and spiked again with 50 mg L−1 of chlorobenzene. The hydrodechlorination restarted with a similar rate as in the first run (APd = 5 L g−1 min−1 ). The pH value was always adjusted for a period long enough to provide adequate kinetic data, e.g. 2 h for the fresh catalyst at pH = 10 and 1 week for the spent catalyst at pH = 3. The reaction bottle was handled in a He-purged plastic container (see above) to minimize undesired air penetration. The O2 -reduction rate was determined by injection of small oxygen pulses. The initial reaction rate dropped down instantaneously from APd = 5.0 L g−1 min−1 at pH = 10 to APd = 0.01 L g−1 min−1 at pH = 3.0, which is a low but still significant value. At this pH value, the fraction of non-protonated hydrazine (cN2 H4 /cN2 H5 + = 10−8.11 /10−3 ≈ 10−5 ) is negligible. Therefore, it seems reasonable to attribute hydrazonium a significant reactivity, which may be estimated to be a factor of 500 lower than that of hydrazine. Of course, a more complex influence of the pH value on the reaction rate cannot be ruled out (e.g. protonation equilibria on the ␥-Al2 O3 catalyst carrier). A continuous loss of catalyst activity was observed for both reactions during the course of the experiment. The oxygen reduction is less sensitive to pH variation and catalyst deactivation than the hydrodechlorination. After 2 weeks of reaction time, H2 was added to the catalyst suspension as a pH-independent reductant. As the data in Table 3 show, the catalyst was still moderately active in the H2 –O2 recombination, but almost completely deactivated for the hydrodechlorination, even under slightly alkaline conditions. All these findings lead to the conclusions not only that hydrazonium is a less reactive H-donor than hydrazine, but also that it acts as a catalyst poison. The catalyst deactivation in the presence of hydrazonium is relatively slow. It can only partly be reversed by shifting the pH value back to alkaline

conditions. After adding a fresh catalyst sample to the suspension from the long-term experiment, the fresh catalyst developed its full hydrodechlorination activity under alkaline conditions. This means, hydrazonium does not produce a kind of dissolved catalyst poison. The deactivation affects only that portion of the catalyst which was in contact with hydrazine under acidic conditions. 3.4. Economic considerations A rough estimation of the prices for reduction equivalents of H2 , HCOOH, and N2 H4 , which of course may strongly depend on the regional conditions and the applied mass fluxes, resulted in a ratio of 35:50:150 /kmol (2H). Hydrogen gas was calculated for large volume tanks (3000 Nm3 H2 at 4.5 MPa) rather than for the frequently used smaller tanks (10 Nm3 at 20 MPa), which have specific costs higher by a factor of about 4. The on-site electrochemical generation of hydrogen is an interesting alternative to storage tanks, especially for small scale production and from a safety point of view. However, the investment costs are significantly higher than for rented storage tanks. The most attractive H2 supply would be a small, low pressure pipeline connected with a larger hydrogen generation and distribution system, such as is available in many chemical factories. In summary, our estimation (which is based on information supplied by several German chemical companies) indicates that gaseous hydrogen and formic acid are comparable with respect to their cost-effectiveness, whereas hydrazine is significantly more expensive. Clearly, storage, handling and dosage of the liquid reagents are technically less demanding. Furthermore, there might be a potential for cost savings with formic acid and hydrazine, if qualities lower than the usually offered purity and concentration standards (85 and 35% concentrated aqueous solutions, respectively) can be accepted.

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4. Conclusions Formic acid, isopropanol and hydrazine were investigated as reductants for the Pd-catalyzed hydrodechlorination of chlorobenzene in aqueous suspension at ambient temperature. The intention was to assess alternatives to molecular hydrogen having high water solubilities. Formic acid was found to be as reactive as H2 under acidic and neutral conditions, but less reactive under alkaline conditions. The observed kinetics imply two pH-controlled reaction mechanisms, which are tentatively interpreted as H-atom and hydride transfer. H-consumers, such as chlorinated compounds and oxygen, strongly stimulate the decomposition of formic acid. Isopropanol is less reactive by about five orders of magnitude. Hydrazine is effective as a H-donor for the hydrodechlorination only under alkaline conditions. However, the reaction is a factor of 30 slower than with H2 . The reaction rate is inversely dependant on the pH value for both formic acid and hydrazine as reductants. The protonated acid is more active than the formate, whereas the non-protonated base is more active than the hydrazonium ion. The common feature of the two reactants is that the non-ionic species are more active than the corresponding ionic species in each case. A possible explanation of this finding could be that the adsorption of ionic species from aqueous solution onto the Pd surface is less favourable than that of their non-ionic counterparts. From the technical and economic point of view, formic acid is a promising substitute for H2 .

Acknowledgements The authors thank the BMBF (German Federal Ministry of Education and Research) for financial support within the SAFIRA project (Remediation Research in Regionally Contaminated Aquifers) and Dr. Detlev Fritsch from the GKSS Research Centre, Geesthacht for preparation of the membrane-based Pd catalyst and helpful discussions.

References [1] R.W. Gillham, S.F. O’Hannesin, Groundwater 32 (1994) 958. [2] T. Bigg, S.J. Judd, Environ. Technol. 21 (2000) 661. [3] F.-D. Kopinke, K. Mackenzie, R. Koehler, A. Georgi, U. Roland, H. Weiß, Chemie-Ingenieur Technik 75 (2003) 329. [4] W. Zhang, J. Nanoparticle Res. 5 (2003) 323. [5] W.W. McNab, R. Ruiz, Chemosphere 37 (1998) 925. [6] C. Schueth, M. Reinhard, Appl. Catal. B 18 (1998) 215.

[7] C. Schueth, S. Disser, F. Schueth, M. Reinhard, Appl. Catal. B 28 (2000) 147. [8] L. Prati, M. Rossi, Appl. Catal. B 23 (1999) 135. [9] B. Coq, F. Figueras, J. Mol. Catal. A 173 (2001) 117. [10] K. Pirkanniemi, M. Sillanpää, Chemosphere 48 (2002) 1047. [11] N.E. Korte, J.L. Zutman, R.M. Schlosser, L. Liang, B. Gu, Q. Fernando, Waste Manage. 20 (2000) 687. [12] K. Mackenzie, R. Koehler, H. Weiß, F.-D. Kopinke, in: G.B. Wickramanayake, A.R. Gavaskar, A.S.C. Chen (Eds.), Chemical Oxidation and Reactive Barriers: Remediation of Chlorinated and Recalcitrant Compounds, Battelle Press, Columbus, Richland, 2000, p. 331. [13] F.J. Urbano, J.M. Marinas, J. Mol. Catal. A 173 (2001) 329. [14] B. Coq, G. Ferrat, F. Figueras, J. Catal. 101 (1986) 434. [15] D. Fritsch, K. Kuhr, K. Mackenzie, F.-D. Kopinke, Catal. Today 82 (2003) 105. [16] F.-D. Kopinke, K. Mackenzie, R. Koehler, Appl. Catal. B 44 (2003) 15. [17] F.-D. Kopinke, R. Koehler, K. Mackenzie, Ch. Schueth, Grundwasser 7 (2002) 140. [18] P. Albers, J. Pietsch, S.F. Parker, J. Mol. Catal. A 173 (2001) 275. [19] J.C. Rodriguez, J. Santamaria, A. Monzon, Appl. Catal. A 165 (1997) 147. [20] N.S. Figoli, P.C.L. Argentiere, J. Mol. Catal. A 122 (1997) 141. [21] D.R. Lide, CRC Handbook of Chemistry and Physics, 75th ed., CRC Press, Boca Raton, 1994. [22] C. Grittini, M. Malcomson, Q. Fernando, N. Korte, Environ. Sci. Technol. 29 (1995) 2898. [23] W. Zhang, C.-B. Wang, H.-L. Lien, Catal. Today 40 (1998) 387. [24] Y. Liu, F. Yang, P.L. Yue, G. Chen, Water Res. 35 (2001) 1887. [25] E. Kempin, Synthese und Charakterisierung katalytisch aktiver Membranen auf Basis von Poly(vinylalkohol), Thesis, Berlin, 2001. [26] M.S. Hähnlein, Entwicklung und Charakterisierung von Edelmetallträgerkatalysatoren und Edelmetallnanosolen zur katalytischen Nitratund Nitritreduktion sowie zur Sorboseoxidation, Thesis, Braunschweig, 1999. [27] J. Yu, J.B. Spencer, Chem. Eur. J. 5 (1999) 2237. [28] R.B. King, N.K. Bhattacharyya, K.D. Wiemers, Environ. Sci. Technol. 30 (1996) 1292. [29] R.B. King, N.K. Bhattacharyya, H.D. Smith, K.D. Wiemers, Environ. Sci. Technol. 32 (1998) 3178. [30] D. Sander, W. Erley, J. Vac. Sci. Technol. A 8 (1990) 3357. [31] G.-Q. Lu, A. Crown, A. Wieckowski, J. Phys. Chem. B 103 (1999) 9700. [32] D.M. Ruthven, R.S. Upadhye, J. Catal. 21 (1971) 39. [33] U. Pruesse, M. Kroeger, K.-D. Vorlop, Chem.-Ing.-Tech. 69 (1997) 87. [34] N. Aas, M. Bowker, J. Phys.: Condens. Mat. 3 (1991) 281. [35] F. Solymosi, I. Kovacs, Surf. Sci. 259 (1991) 95. [36] H. Wiener, Y. Sasson, J. Blum, J. Mol. Catal. 35 (1986) 277. [37] H. Wiener, J. Blum, Y. Sasson, J. Org. Chem. 56 (1991) 4481. [38] M.K. Anwer, D. Sherman, J.G. Roney, A.F. Spatola, J. Org. Chem. 54 (1989) 1284. [39] G. Del Angel, J.L. Benitez, J. Mol. Catal. A 165 (2001) 9. [40] J.L. Benitez, G. Del Angel, React. Kinet. Catal. Lett. 70 (2000) 67. [41] P. Forni, L. Prati, M. Rossi, Appl. Catal. B 14 (1997) 49. [42] M.A. Aramendia, V. Borau, I.M. Garcia, C. Jimenez, F. Lafont, A. Marinas, J.M. Marinas, F.J. Urbano, J. Mol. Catal. A 3517 (2002) 1. [43] M.A. Aramendia, V. Borau, I.M. Garcia, C. Jimenez, F. Lafont, A. Marinas, J.M. Marinas, F.J. Urbano, J. Catal. 187 (1999) 392.