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J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3315-3322 3315 Time-resolved Microwave Conductivity Part 1.-TiO, Photoreactivity and Size Quantization Scot T. Martin, Hartmut Herrmann,? Wonyong Choi and Michael R. Hoffmann" W. M. Keck Laboratories, California lnsiituie of Technology, Pasadena, CA 91125,USA ~ ~ ~~ Charge-carrier recombination dynamics after laser excitation are investigated by time-resolved microwave con- ductivity (TRMC) measurements of quantum-sized (Q-) TiO, , Fell'-doped Q-TiO, , ZnO and CdS, and several commercial bulk-sized TiO, samples. After pulsed laser excitation of charge carriers, holes that escape recom- bination react with sorbed trans-decalin within ns while the measured conductivity signal is due to conduction- band electrons remaining in the semiconductor lattice.The charge-carrier recombination lifetime and the interfacial electron-transfer rate constant that are derived from the TRMC measurements correlate with the CW photo-oxidation quantum efficiency obtained for aqueous chloroform in the presence of TiO,. The quantum efficiencies are 0.4% for Q-TiO, , 1.6% for Degussa P25, and 2.0% for Fell'-doped Q-TiO, . The lower quantum efficiencies for Q-TiO, are consistent with the relative interfacial electron-transfer rates observed by TRMC for Q-TiO, and Degussa P25. The increased quantum efficiencies of Fell'-doped Q-TiO, and the observed TRMC decays are consistent with a mechanism involving fast trapping of valence-band holes as Fe'" and inhibition of c harg e-o rde r recom bi nation.When the crystallite dimension of a semiconductor particle falls below a critical radius of ca. 10 nm, the charge carriers appear to behave quantum-mechanically as a simple particle in a box.'-6 As a result of this confinement, the bandgap increases and the band edges shift to yield larger redox potentials. The solvent reorganization Gibbs energy for charge transfer to a substrate, however, remains unchanged. The increased driving force and the unchanged solvent reor- ganization Gibbs energy in size-quantized systems are expected to increase the rate constant of charge transfer in the normal Marcus regi~n.~-~ Thus, the use of size-quantized semiconductor TiO, particles may result in increased pho- toefficiencies for systems in which the rate-limiting step is charge transfer.'o~ll One such system is the oxidation of many common organic pollutants in the presence of TiO, irradiated with bandgap illumination.' ,-' The use of size-quantized semiconductors to increase pho- have been studied previously.P25 is a commercial form of TiO, that generally has a higher photoreactivity than other available forms of TiO, . The variable photoefficiencies of the different forms of TiO, are related to their fundamental charge-carrier dynamics. In order to verify this relationship, we investigate the charge-carrier dynamics of Q-TiO, and P25 by TRMC meas~rements.~~-~~In a typical TRMC experiment, separat- ed charge carriers, which are generated by a laser pulse, lead to a perturbation of the initial microwave ab~orbance.~~.~~ The temporal decay of the conductivity signal (i.e.microwave absorbance) reflects the lifetime of the photogenerated car- riers.The technique has only recently been expanded to semi- conductor parti~les,~~-~~*~~.~~for which conventional techniques (e.9. photoconductivity) are often not possible owing to the necessity for electrode contacts. Efforts have been made here to provide experimental details of the TRMC toefficiencies is supported by several st~dies.'~*'~-'~technique, including the development of However, in other work, size-quantized semiconductors have been found to be less photoactive than their bulk-phase counterparts.' '9,' In the latter cases, surface speciation and surface defect density appear to control phot~reactivity.~'-~~ The positive effects of increased overpotentials (ie.difference between E,, and Etedox)on quantum yields can be offset by unfavourable surface speciation and surface defects due to the method of preparation of size-quantized semiconductor par- ticles. In the present study, the photodegradation of several chlo- rinated compounds in the presence of Q-TiO, are used as control reactions to study the size-quantization effect on pho- toreactivity. The stoichiometry of the reactions is as follows : C,H,Cl, + (X + 7)02 xCO, + zH' + zC1-+ (';')H20-(1) hv. Ti02 The photodegradations of chlor~form,~~pentachlorophe-n01,25926 and 4-chloropheno12 7-31 with Degussa P25 (ie.a bulk-phase TiO, consisting of 80% anatase and 20% rutile) Present address: Institut fur Physikalische und Theoretische Chemie der Universitat GH Essen, Germany. a novel sample holder that is usable in conjunction with conventional lasers. Experimenta1 Preparation Q-TiO, was prepared by the controlled hydrolysis of titanium(1v) tetraisopr~poxide.~' 5 ml Ti[OCH(CH,),], (Aldrich, 97%) dissolved in 100 ml isopropyl alcohol was added dropwise (90-120 min) with vigorous stirring to 900 ml doubly distilled water (2 "C) adjusted to pH 1.5 with HNO, . The transparent colloid can be stored for over one year in a cold room (4"C) without coagulation. To obtain a powdered sample, 150 ml of the colloidal solution was evaporated (35 "C) using a Rotavapor (model R110).The resulting film was dried with an N, stream to yield a white powder. Fe"'- doped Q-TiO, was prepared by a similar procedure in the presence of Fe(NO,), to give an atomic doping level of 1%. Full incorporation of Fe'" into the lattice has been shown by other worker^.^^.^^ The bulk-phase semiconductors used were ZnO (Baker), a-Fe203 (hematite, Fisher), CdS (Alfa), and TiO, (Degussa P25 and Sachtleben Chemie S7, S13, S17, S18, S21 and S24). We used two separate batches of Degussa P25, which were obtained in 1988 (P25-A) and 1993 (P25-B). Characterization Particle sizes were determined by a Philips EM 430 transmis- sion electron microscope (TEM) at 300 kV. Samples for TEM were prepared by placing a drop onto a copper mesh sub- strate covered with a carbon film, followed by removal of the excess liquid with a piece of thin filter paper and drying for 30 s under a tungsten lamp.A representative TEM micro- graph is shown in Fig. 1. The sizes of the Q-particles ranged from 2 to 4 nm with a lattice spacing of 3.6 & 0.1 A. This spacing was in good agreement with the anatase (101) phase lattice spacing of 3.51 X-Ray diffraction (XRD) analysis was carried out with powdered samples on a Scintag PAD5 model DMC-008 using 35 kV, 20 mA Cu-Kcr (1.54 A) radi-ation. The diffraction pattern of the Q-sized TiO, was also characteristic of anatase. The observed line broadening due to the presence of small crystallites was analysed by the Scherrer equation46 and showed that the particles were 3-4 nm in diameter.Degussa P25 has been characterized previ~usly.~'.~~ 30 nm crystallites composed of 80% anatase and 20% rutile aggre- gate to form particles with an average diameter of 1 pm. For the TRMC experiments, P25 (1.44 g 1-') was suspended in HNO, (pH 1.5) and rotary evaporated to a dry powder. The characterization of the TiO, samples obtained from Sachtle- ben Chemie is shown in Table l.,' Irradiation Steady-state photolyses were carried out in a slurry reactor to determine the initial rate constants for the degradation of chloroform (Baker), dichloroacetic acid (Spectrum Chemical Manufacturing, Inc.), carbon tetrachloride (Baker), pentach- lorophenol (Aldrich), and 4-chlorophenol (Aldrich).Irradia- tions were performed with a 1000 W Xe arc lamp (Spindler and Hoyer). The IR component of the incident light was J. CHEM. SOC. FARADAY TRANS., 1994, VOI,. 90 Table 1 TRMC results and photoreactivity data of commercial samples of TiO, relative charge-carrier interfacial electron- surface area quantum" concentration transfer rate sample /m2 g-' efficiency (at 100 ns)/mV constant/ms-' s7 85 0.25% 0.60 0.12 0.06S13 90 0.24% 1.28 S17 380 0.07% 0.12 n/a 0.16S18 230 0.26% 0.52 0.34s21 280 0.44% 0.52 S24 30 0.04% 0.16 n/a S25-B 50 0.39% 0.92 0.13 [HCCI,] = 3.2 mmol 1-' [TiO,] = 0.5 g I-', I = 214 pin 1-' min-', ,i= 320 _+ 5 nm, air equilibrated, pH 4-6. removed by a 10 cm water filter. Depending upon the experi- ment, wavelengths were selected with an interference filter (Oriel, A = 320 f5 nm), a longpass filter (Oriel, A > 320 nm), or a bandpass filter (Corning 7-60-1, 320 < A -= 380 nm).Light intensity was adjusted with neutral density filters. The chemical actinometer Aberchrome 540 { [(E)-ct-(2,5-dimethyl-3-furyl)ethylidene]-3-isopropylidenesuccinicanhydride) was used to determine the incident light intensity, which was found to vary between 100 and 200 pmol 1-' min-' with the interference filter in place and to be 1000 pmol-' 1-' min-' with the longpass filter." Aqueous suspensions (35 ml) of the chlorinated compounds and TiO, (1.0 g 1-') were prepared and the pH was adjusted by the addition of HNO,. Initial degradation rates were determined by the total C1- release after 1 h illumination in the case of HCCI, (63 mmol l-'), DCA (4.8 mmol 1-') and CCl, (5.1 mmol 1-I) and by HPLC (Hewlett Packard Series I1 1090 Liquid Chromatograph) analysis for PCP (60 pmol 1-') and 4-CP (100 pmol 1-').Chloride concentrations were determined with an Orion chloride-selective electrode (model no. 9617B). TRMC Measurements A schematic diagram of the microwave conductivity appar- atus is shown in Fig. 2. A Gunn diode microwave source (100 mW, 38.3 GHz, MACOM Inc.), a PIN diode microwave detector, a Comlinear, model CLC206AI amplifier, a TEK 2440 digitizing oscilloscope, and an HP 432A power meter were coupled into the TRMC unit. The source, detector, and amplifier were enclosed in aluminium Faraday cages.The waveguide system was R-band WR-28 (0.711 cm x 0.356 cm). A Lambda Physik excimer laser (LPX 120) was used for a 308 nm, 50 ns pulse excitation source. Layered metal mesh rror sample11/ holderDower meter iiy 10 nm Fig. 2 Schematic diagram of the apparatus for TRMC measure-Fig. 1 TEM micrograph of quantum-sized TiO, ments J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 interposed between the laser and the sample was used to control incident light intensity, which was 4.5 mJ per pulse otherwise stated. In a typical experiment, between 32 and 256 conductivity decays were averaged to improve the signal-to- noise ratio. The digitized data were transferred to a computer for storage and data analysis.The data were collected on four timescales (200 ns div.-', 2 ps div.-', 200 ps div.-' and 10 ms div.- '). The transients were reproducible within 5% error. The sample holder, which was designed especially for this series of experiments, is shown in Fig. 3. The top plate (i.e. short) has seven slits cut orthogonal to the propagating microwave mode so that the laser light can enter the wave- guide while the microwaves are reflected back into the wave- guide. The sample was prepared as a thick paste supported in trans-decalin (Aldrich). The paste was moulded into a Teflon holder with an illuminated surface area of 13.9 mm2. The Teflon holder was fitted into the waveguide at a distance of 1.65 mm from the short plate. The holder was positioned by the use of an aluminium block inserted from the rear of the waveguide upon which the holder was pressed.The principles of the TRMC experiment have been dis- cussed previo~sly.~~*~~.~' Microwaves from the source pass through the sample and impact on the PIN diode detector, which then transforms the incident microwave power into a voltage for input to the oscilloscope. The absorbed micro- wave power is directly proportional to the conductivity of the sample, 0, for low-conductivity samples.32 The proportion- ality constant, A, is specific to the geometry of the apparatus and the sample, and it is determined by calibration. The change in absorbed microwave power, AP, due to a change in conductivity, An, caused by carrier excitation is given by eqn.(1) where P is the initial microwave power level. -AAa(t)P The microwave power is transformed into a voltage, I/, by the PIN diode detector as P = V.For perturbations below 3%, the response is linear, as shown in eqn. (2).32 AP P The proportionality constant, n, is typically between 1 and 2 and is found by calibration. Substitution of eqn. (1) and eqn. (2) results in eqn. (3). AV(t) = (?)A&) (3) The temporal behaviour of the voltage observed at the digitizer and the conductivity of the sample are thus propor- tional. top view: side view: 7 slit short: inside waveguide Fig. 3 Sample holder for powdered semiconductor Principles of TRMC When free charge carriers couple to the electric field of the microwaves, absorption occurs.39 The strength of the inter- action is expressed in terms of mobilities.The mobilities of a free electron in He gas (1 m2 V-' s -I), of free carriers in Si (0.2), of a hopping electron in an organic compound of ions in solution (lO-'-lO-*), and of dipole moments (lo-' to lo-',) reflect the relative coupling efi- ciencies at microwave frequencie~.~~ The interpretation of TRMC measurements of polar molecules in non-polar sol- vents and of free carriers in silicon are generally under- stood;32,33-s1-53 however, the interpretation of the conductivity decays of semiconductor particles with low carrier mobilities m2 V-' s-') has not been addressed previously. In high-mobility semiconductors such as GaAs (8900 m2 V-' s-'), Si (1950), ZnO (380) or CdS (390), the microwave absorption can be attributed to free car- rier~.~~.~~Furthermore, for these semiconductors, with the exception of ZnO, the electron mobility is several times larger than the hole mobility so that the observed conductivity decay is attributable entirely to free electrons.For low- mobility semiconductors, such as TiO, (10-4-10-s m2 V-' s-1)55 and a-Fe203 (10-s),s6 the microwave absorbance may not be due exclusively to free carriers. The mobility of a shallow trap may be similar to that of a hopping electron in an organic compound (10-3-10-4 m2 V-' s-I); therefore, in these cases free and shallowly trapped carriers may simulta- neously contribute to the conductivity decay of TiO, .Exam-ples of shallowly trapped carriers in TiO, at 25°C include small polarons5' and electrons on Ti"' sites.40 Results Initial degradation rates of chloroform as a function of the concentration of chloroform and the type of TiO, are illus- trated in Fig. 4. At the solubility limit for aqueous chloro- form (63 mmol 1-'), the quantum yield for chloride release is 0.4% for Q-TiO, and 1.6% for P25-A. Doping of the Q-TiO, with Fe"' at 1.0 atom% shows an increase in the quantum yield to 2.0%. Initial degradation quantum efficiencies of chloroform, dichloroacetic acid (DCA), carbon tetrachloride, pentachlorophenol (PCP) and 4-chlorophenol (4-CP) are shown in Table 2 for Q-TiO, and P25-A. The quantum efi- ciencies obtained with Q-TiO, are less than those obtained with P25-A.The observed microwave conductivity decays of ZnO, CdS, P25 and Al,03 are shown in Fig. 5. Since the signal strengths and timescales of decay vary over several orders of 2.5 r 1k 1E 2.0:I2 1.5 nn0*5* V.W 0 10 20 30 40 50 60 [HCCl,]/mmol I-' Fig. 4 Degradation rate of chloroform as a function of concentra- tion. A, P25-A: .,Q-TiO, ; a, Fe"' doped (1%) Q-TiO,. Condi- tions: pH 2.8 (HNO,), [TiO,] = 0.5 g l-', I = 110 pmol 1-' min-' (A= 320 5 nm), air equilibrated. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Quantum efficiencies of Q-TiO, and Degussa P25-A towards photomineralization of chlorinated compounds substrate concentration/mmol 1-' measurement QQTiOy ("/.I QPZS-A ("/.I ~ ~~ chloroform" 63 0.4 1.6 dichloroacetic acid" 4.8 1.1 23.9 carbon tetrachloride"*b 5.1 0.4 3.1 pentachlorophenol" 60 0.3 0.4 Cchlorophenol' 100 0.4 1.8 " 100 pein 1-' min-', 310 < A/nm < 330, pH 2.5-3 (HNO,), [TiO,] = 1 g 1-'.* 0.1 moll-' MeOH. lo00 vein 1-' min-', 1 > 320 nm, pH 3 (HNO,), [TiO,] = 1 g 1-' I I I I 1 I I I I lo-' lo4 10" lo-* time/s Fig. 5 Double log plot of time-resolved microwave conductivity decay of several powdered semiconductors supported in trans-decalin. A,ZnO; +,CdS; ., Al,O,.P25-A; ., magnitude, the data are represented in log-log form. Alumina (a-Al,03), which has a bandgap of 9.5 eV, serves as a blank for the apparatus and indicates the minimum detection limit.35 For direct comparison, NaCl shows a similar decay profile.The relative rates of decay can be characterized by the half-life signal decay. For ZnO, the first half-life is 6.6 ps and the second is 164 ps. For CdS, the first half-life is 60 ns, the second is 770 ns, and the third is 9.2 ps. For P25-A, the first half-life is 1.1 ps and the second is 1.7 ms. Several investi- gators have reported that the half-lives for conductivity decay in CdS and TiO, increase as the decay time increases; our observations are consistent with these previous reports.30*40*41For example, Schindler and Kunst reported a T',, for P25 for 2.0 ps for the conductivity observed at 100 ~s.~OThis value agrees reasonably with our value of 1.1 ps for P25-A. The initial degradation rates of chloroform as a function of light intensity are shown in Fig. 6.Two linear regions are 6.0r. 13.5 Y 5.0 3.0 7 & Y.-E 2.55 4.0 C 5 3.0 2.0 :gE, F9 1.5 E n Gm 2.0 *3 0 1.0 5 0 E 1.0 0.5 I 1 I 0.00.0 0 500 1000 1500 light intensity/pmol I-' min -' Fig. 6 Degradation rate of chloroform as a function of light inten- sity. Conditions: pH 11 (NaOH), [P25-A] = 0.5 g I-', 320 < A/ nm < 380, air equilibrated. observed, and the cross-over occurs at 150 pmol 1-' min-'. The quantum efficiencies vary from 2.5% for I = 50 pmol 1-min-' to 0.3%for Z = 1490 pmol 1-' min-'. In Fig. 7, the time-resolved conductivities at 100 ns of P25-A prepared in HNO, and trans-decalin are plotted as a function of the inci- dent laser pulse energy.Two linear regions are identified with an apparent cross-over at a pulse energy of 6 mJ. Representative conductivity decays for TRMC measure- ments of S7, S13, S17, S18, S21, S24 and P25-B are shown in Fig. 8. For most samples, the interfacial electron-transfer rate constants reported in Tables 1 and 3 are calculated by single- exponential fits of the conductivity data. However, in the case of S17 and S24, the signal-to-noise ratio is too low to facili- tate an exponential fit of the data. Discrepancies between similar samples in Tables 1 and 3 are due to differences in the pre-amplifier, the incident pulse energy (2.5 mJ in Table 1), CW illumination intensity, and batches of P25. A contour plot of the quantum efficiencies reported in Table 1 as a function of the recombination lifetime and the interfacial electron-transfer rate constant is shown in Fig.9(a). The arrows on the data points for S17 and S24 indicate that the interfacial electron-transfer rate constant is unknown. The charge-carrier recombination lifetimes (uide infra) are the charge-carrier concentrations (at 100 ns) report- Table 3 TRMC results for Q-TiO, , Fe"'-doped Q-TiO, , P25-A (No NO,-) and P25-A (NO,-) re1 at ive charge-carrier concentration interfacial electron-transfer sample (at 100 ns)/mV rate constant/ms-' Q-TiO, 3.1 0.052 P25-A (NO NO3-)P25-A (NO,-) Fe"'doped Q-Ti0 , 4.0 2.9 1.o 0.078 0.056 0.065 4.0 > E..2.0 0.-v) 0.00.0 10.0 20.0 30.0 40.0 pulse energy/mJ Fig.7 Effect of the incident laser pulse energy on the initial (100 ns) conductivity of P25-A prepared in HNO, and trans-decalin J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 3319 A 3.0-I1 > -2.5 IDC cx,.-UJ 2.0 1.5 0:O 0.'5 110 1.k 210 2:5 time/ps 1.411 B 1.3 > 1.2E2 1.1 ea3 1.0 0.9 0.8 0.7 0 20 40 60 time/ms Fig. 8 Representative conductivity decays of Sachtleben and Degussa TiO, powders. A, (a)S13, (b)P25-B,(c) S18 and (6)S17,2 ps div.-'. (b) Exponential fit of conductivity decay for P13,k = 0.056 ms-', 10 ms div.-'. n 1.40 4-.-C 1.05 0.70 0.35 0.00 .oo 0.10 0.20 0.30 0.400. k/rns-' s21 +-+---+ S24 0.0 I I I I 0.00 0.05 0.10 0.15 0.20 [carrier]k/mV ms-' ed in Tables 1 and 3.The contours drawn fit the function C = xy. A plot of C us. the product of x and y results in a straight line, as shown in Fig. 9(b) where C is the quantum efficiency and xy is the contour value, i.e. the product of the recombination lifetime and the interfacial electron-transfer rate constant, k. The data points for S17 and S24 are calcu- lated using k = 0.15 ms-' and the extrema of the arrows are calculated for k = 0.03 ms-' and k = 0.30 ms-'. Fig. 9(b) demonstrates more clearly than Fig. 9(a) the correlations between photoreactivity and charge-carrier dynamics. To investigate the effect of HNO, ,P25-A was prepared by rotary evaporation from an acidic slurry (HNO,, pH 1.5) and was then supported in a trans-decalin paste for the con- ductivity measurements. In a separate preparation, P25-A was directly prepared as a paste in trans-decalin.The conduc- tivity decays are shown in Fig. 10. It is apparent that the temporal behaviour of the coated and uncoated samples is similar. In addition, the conductivity decays of Q-TiO, and Fe*-doped Q-TiOz are overlaid in Fig. 10. The TRMC mea- surements for these samples are summarized in Table 3. Discussion In previous papers, we proposed that the larger over-potentials in Q-TiO, us. bulk-phase TiO, should lead to higher quantum yields."*' 138 H owever, the data in Fig. 4 clearly show that Q-TiO, is less photoreactive than P25-A. In similar experiments carried out with substrates represent- ing a variety of postulated mechanistic pathways, including direct hole attack on HCCl, and DCA,24 hydroxyl radical attack of 4-CP and PCP,27-31.59 and electron transfer to CCl, ,'0*11-42*60 the quantum efficiencies (cf.Table 1) AI 4.0 3.0> E2 2.0 .-0 I II I * 1.0t I 0.0 I I I 1 I-1.0 ' -4.00 0 400 800 1200 1600 time/ns I I I 1I la5 t h R"1 > 1.0 E L C ea 'G 0.5 0.0 -20 0 20 40 60 80 ti me/ms Fig. 10 Conductivity decays of P25-A (c), +; Q-TiO, (b),A and Fe"'-doped Q-TiO, (4,Fig. 9 (a) Contour plot of quantum efficiency as a function of ..Samples prepared by rotary evaporation recombination lifetime (seeexplanation in main text) and interfacial electron-transfer rate constant.(b)Linear transformation of contour from HNO, (pH 1.5)and supported in trans-decalin except for P25-A-no NO,-(a),0,which was prepared without HNO, . A, 200 plot. ns div.-' timebase. B, 10 ms div.-timebase. obtained for Q-TiO, appear to be consistently lower than those obtained for bulk-phase P25. These data suggest several possibilities. On the one hand, the lower photoreacti- vity of Q-TiO, may be due to an increased rate of charge- carrier recombination. On the other hand, there may be substantial differences in the interfacial charge-transfer rates between Q-TiO, and P25-A. The differences between Q-Ti02 and P25-A are at least partially understood by considering the preparation methods that may result in more defect sites (trapping sites) and faster recombination rates in Q-TiO, .P25-A TiO, is prepared in a high-temperature flame reactor and is thus expected to have fewer defe~ts~~?~~ whereas Q-TiO, is prepared by sol-gel techniques at much lower temperatures. Different surface morphologies (e.g. hydroxylation density) may also be expected and the interfacial charge-transfer rates may be con- trolled by the relative formation of surface complexes on Q-TiO, as compared to P25-A. The relationship between the charge-carrier recombination rate and quantum efficiency can be expressed as follows: ratetransfer +transfer = (4) ratetransfer + raterecornbination If the charge-carrier recombination rate in Q-TiO, increases due to defects, then the quantum efficiency of interfacial charge transfer decreases and Q-Ti02 should be less photo- reactive than P25.The observed conductivity decays in our TRMC experi- ments should yield the recombination rates of photogener- ated free charge carriers. However, the TRMC response is influenced by other deactivation pathways such as interfacial charge transfer to adsorbed species. In these cases, the con- ductivity decays may be due to both recombination and charge transfer, and the data must be deconvolut-ed,34-38*40941In the absence of recombination or charge-carrier localization during the laser pulse, the conductivity measured after the laser pulse should be proportional to the number of free charge carriers: An = qN(pe-+ ph+) (5) where N is the number of adsorbed photons (if equal reflec- tivities and short penetration depths are assumed among the samples).The data of Fig. 5 are consistent with the prediction of eqn. (5) that the strength of the conductivity signal increases with the mobility of the charge carriers (1pznO= 380 cm2 V-' CpCdS=30034 and 1pTioz= 135).From eqn. (3) we see that AV = (qNVA/n)(pe-ph+) (6) Eqn. (6)predicts a linear correlation between the post-pulse conductivity, AV, and the sum of the mobilities with a corre- sponding slope of (qNV/n)A.Using this relationship, our spe- cific apparatus constant, A, can be evaluated. A linear fit applied to the data in Fig. 5 yields a slope of 0.28 mV/(cm2 v-1 s-l), r2= 0.81 and A = 1.1 x lo3 S-'.Since r2 < 1, there is a high probability that localization (i.e. recombi-nation or interfacial charge transfer) has taken place during the laser pulse. Thus, we consider the above A value to be a lower estimate. Photoelectrochemical Mechanisms We propose the following four processes for charge-carrier recombination :5976 l-13~ charge-carrier generation J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 direct and indirect charge-carrier trapping charge-carrier recombination interfacial charge transfer kvm ecb-+ 0 ___* 0-(VIII) kx hv,+ + R -R+ where eT- is a trapped electron, hT+ is a trapped hole, T,-is an empty electron trap, Th, is an empty hole trap, 0 is an electron acceptor (oxidant) and R is an electron donor (reductant).At present, we believe that the electron is trapped in a surface Ti"' ~ite~~v~~.~~while the hole is trapped in a surface hydroxy gro~p.~~-~~*~~ Based on this mechanism we can write an equation for the change in microwave conductivity as follows : For ZnO, peT-(d[eT-]/dl) and phT+(d[hT+]/dt) can be omitted from eqn. (7) owing to the high mobility of the free charge- carriers. However, in the case of TiO,, these terms must be included since shallowly trapped and free electrons have com- parable mobilities. Even though the observed microwave conductivity signal is due to a mixture of species and mobi- lities, we believe that the observed conductivity as shown in Fig. 8 and 10 can be assigned primarily to electrons In order to explain the timescales for the microwave con- ductivity decays shown in Fig.8 and 10, a mechanism that includes interfacial charge transfer must be invoked because recombination is complete in CQ. 100 ns in the absence of interfacial charge In order for the kinetics of the charge transfer to compete with recombination processes internal to the TiO, particle, interfacial charge transfer of at least one carrier should occur within several ns.68 Therefore, only one charge-carrier type (i.e. holes or electrons) should be present in the particle after 100 ns. Because hole transfer often takes place within ns while electron transfer takes place over the timescale of ns to ms,42964*69 we conclude that elec- trons give rise to the measured TRMC conductivity.The overall quantum efficiency for interfacial charge trans- fer is determined by two critical processes: the competition between carrier recombination and trapping (ps to ns) fol- lowed by the competition between carrier recombination and interfacial charge transfer (pto ms). The measurement of the J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 conductivity at 100 ns gives information on fast recombi- nation while the measurement at longer timescales yields information on interfacial charge transfer. The fall-off in the remaining charge carriers at 100 ns with increasing injection level is shown in Fig. 7. The apparent discontinuity at 6 mJ suggests that a higher-order channel is opened with fewer residual charge carriers at 100 ns.These results are consistent with the inverse relationship observed between quantum effi- ciency for CHC1, oxidation and light intensity shown in Fig. 6.'6924 Based on these results, we believe that the recombi- nation lifetime of charge carriers is inversely proportional to the conductivity at 100 ns. An increase in either the recombination lifetime of charge carriers or the interfacial electron-transfer rate constant is expected to result in higher quantum efficiencies for CW pho- tolysis. The samples S7-P25 are observed to follow this relationship in Fig. 9(a). The linear transformation of the contour plot in Fig. 9(b) makes the correlation more appar- ent. Fig. 9(a) suggests S21 owes its high photoreactivity to a fast interfacial electron-transfer rate constant whereas P25 has a high photoreactivity due to slow recombination.Bickley et al. have suggested that the anatase-rutile structure of P25 promotes charge-pair separation and inhibits recombi- nation.47 The different recombination lifetimes and interfacial electron-transfer rate constants may be due to the different methods of preparation of the samples that result in different crystal defect structures and surface morphologies. A conduction band electron is thermodynamically capable of reducing H', NO,-, 0, and oxidized trans-decalin rad- icals (T+)while a valence band hole is sufficiently powerful to oxidize trans-decalin. The signal strength observed in Fig. 1qa) for P25-A in the presence of HNO, is reduced relative to the uncoated P25-A.This relative change may be due to a fast reduction of H+ or NO,-in the first monolayer or to an inhibition of hole transfer which results in greater charge- carrier recombination. However, the time dependence of the conductivity decay appears to be unchanged in the presence of HNO, . In contrast, the fast timescale recombination observed for Q-TiO, and P25-A (NO,-) appears to be similar as shown in Table 2. This observation suggests that a similar number of defects are present in each material. In general, the time-dependent changes summarized in Fig. 10 indicate that electrons undergo interfacial charge transfer more slowly for Q-TiO, (k = 0.052 ms-') than for P25-A (k = 0.065 ms-').This result is consistent with the observed lower steady-state quantum yields (Table 1). Fe"'-Doped TiO, EPR and transient absorption studies have shown that Fe"' doping in colloidal TiO, acts by trapping holes as Fe" within several ns of excitation and that slow recombination takes place by tunnelling from electrons trapped at surface Ti"' sites to holes trapped at bulk-phase FeIV The Fe"' dopant thus acts to inhibit charge-carrier recombination. These processes can be written in terms of reactions (XII) and (XIII), respectively : kxrr Fe"' + hf-Fe" kxrri Fe" + Ti"' -TiO,-Fe"' (XIII) The oxidation of the sorbed electron donor is written as : kxiv Fe" + R -Ti0,-Fe"' + R' A flat microwave signal in Fig. 10 is consistent with the inhi- bition of carrier recombination within the iron-doped sample.However, the conductivity signal at 100 ns is weaker for the Fe"'-doped sample than for the undoped sample. The mobil- ity of the electrons in Fe"'-doped Q-TiO, may be signifi- cantly reduced because Fe"' is present at a concentration of 2.9 x lo2' ~m-~.For example, the mobility of an electron in silicon drops from 1100 cm2 V-' s-' to 100 cm2 V-' s-l in a similary doped lattice." The net effect is a lower initial microwave conductivity for Fe"'-doped Q-TiO, . In summary, we believe the hole. which is trapped at an FetV site, is transferred to an adsorbed substrate on a sub- millisecond timescale while the interfacial electron transfer occurs on the ms scale [eqn.(VIII) and (IX)]. As a result, we predict higher steady-state quantum efficiencies for Fe"'-doped Q-Ti0,. This prediction is supported by results shown in Fig. 4. Conclu,'cions After pulse laser excitation of charge carriers, holes which escape band-gap recombination are transferred to the sorbed electron donor trans-decalin within ns. The TRMC conduc- tivity signals are due to electrons remaining in the semicon- ductor lattice after hole transfer. The resultant interfacial electron transfer takes place over ms and appears to be faster for P25 than for Q-TiO,. The slower electron-transfer rates observed for Q-TiO, are consistent with the lower steady- state quantum yields. Fe"' doped into the Qi-TiO, matrix serves to trap holes as Fe'" and thus reduces charge-carrier recombination, which in turn results in increased quantum efficiencies.The correlations observed between quantum effi- ciencies and charge-carrier dynamics emphasize the impor- tance of the interfacial charge-transfer rate constant and the charge-carrier recombination lifetime as contributing factors to TiO, photoreactivity. We are indebted to Prof. Nathan S. Lewis for the loan of microwave components, to Prof. Geoffrey A. Blake for the use of the excimer laser and to Dr. Detlef W. Bahnemann for providing the Sachtleben Chemie samples. We are grateful to ARPA and ONR {NAV 5 HFMN N0001492J1901) for financial support. S.M. is supported by a National Defense Science and Engineering Graduate Fellowship.H.H. wishes to thank NATO/DAAD for financing a research vist at the California Institute of Technology. Nicole Peill, Dr. Amy Hoffman and Dr. Andreas Termin provided valuable support and stimulating discussion. References 1 L. Brus, Appl. Phys. A,, 1991,53,465. 2 H. Weller, Adc. Mater., 1993, 5, 88. 3 M. Gratzel, Nature (London), 1991,349,740. 4 H. Weller, Angew. Chem., 1993, 32, 41. 5 A. T. Henglein, Curr. Chem., 1988, 143, 113. 6 P. V. Kamat, Chem. Rev., 1993,93, 267. 7 R. A. Marcus and N. Sutin, Biochim. Biophys. Acta, 1985, 811, 265. 8 R. A. Marcus, J. Phys. Chem., 1990,94, 1050. 9 N. S. Lewis, Annu. Rev. Phys., 1991,42, 543. 10 A. J. Hoffman, G. Mills, H. Yee and M. R. Hoffmann, J. Phys. Chem., 1992,96, 5546. 11 A.J. Hoffman, H. Yee, G. Mills and M. R. Hoffmann, J. Phys. Chem., 1992,96,5540. 12 Photocatalysis and Environment : Trends and Applications, ed. M. Schiavello, Kluwer, Dordrecht, 1988. 13 Homogeneous and Heterogeneous Photocatalysis, ed. E. Pelizzetti and N. Serpone, Reidel, Dordrecht, 1986. 14 Photoelectrochemistry, Photocatalysis and Photoreactors, ed. M. Schiavello, Reidel, Dordrecht, 1985. 15 Photocatalysis: Fundamentals and Applications. ed. N. Serpone and E. Pelizzetti, Wiley, New York, 1989. 3322 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 16 17 18 19 20 21 A. J. Hoffman, E. R. Carraway and M. R. Hoffmann, Enuiron. Sci. Technol., 1994,28, 776. M. Anpo, T. Shima, S. Kodama and Y. Kubokawa, J. Phys. Chem., 1987,91,4305.J. M. Nedeljkovic, M. T. Nenadovic, 0.I. Micic and A. J. Nozik, J. Phys. Chem., 1986,90, 12. Y. Nosaka, N. Ohta and H. Miyama, J. Phys. Chem., 1990, 94, 3752. P. Giuseppe, C. H. Langford, J. Vichova and A. Vleck, J. Photo-chem. Photobiol. A: Chem., 1993,7567. W. Lee, Y-M. Gao, K. Dwight and A. Wold, Mater. Res. Bull., 1992,27,685. 43 44 45 46 47 48 49 50 J. Moser, M. Gratzel and R. Gallay, Helu. Chim. Acta, 1987, 70, 1596. D. W. Bahnemann, Zsr. J. Chem., 1993,33,115. Powder Diffraction File, Sets 22-22; JCPDS: Swarthmore, 1980; vol. PDlS-22iRB, pp. 21-1272. B. D. Cullity, Elements of X-Ray Diffraction, Addison-Wesley, Reading, 2nd edn., 1978, p. 102. R. I. Bickley, T. Gonzalez-Carreno, J. S. Lees, L. Palmisano and R. J. D. Tilley, J. Solid State Chem., 1991,92, 178.Degussa Technical Bulletin, No. 56, 1990. D. W. Bahnemann, personal communication. H. G. Heller and J. R. Langan, J. Chem. SOC., Perkin Trans. 2, 22 23 S. Nishimoto, B. Ohtani, H. Kajiwara and T. Kagiya, J. Chem. SOC., Faraday Trans. I, 1985,81,61. B. C. Faust, M. R. Hoffmann and D. W. Bahnemann, J. Phys. Chem., 1989,93,6371. 51 52 1981,341. P. P. Infelta, M. P. de Haas and J. M. Warman, Radiat. Phys. Chem., 1977,10,353. R. W. Fessenden, P. M. Carton, H. Shimamori and J. C. 24 25 C. Kormann, D. W. Bahnemann and M. R. 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Tabata, CRC Press, Boca Raton, FL, 1991, ch. 6.J. M. Warman, M. P. de Haas, S. W. F. M. van Hovel1 tot West- 55 56 57 58 59 60 61 62 H. 0. Finklea, in Semiconductor Electrodes, ed. H. 0.Finklea, Elsevier, New York, 1988, p. 52. M. Anderman and J. H. Kennedy, in Semiconductor Electrodes, ed. H. 0.Finklea, Elsevier, New York, 1988, p. 153. N. F. Mott and E. A. Davis, in Electronic Processes in Non- Crystalline Materials, Clarendon Press, Oxford, 1971, pp. 117ff. D. W. Bahnemann, C. Kormann and M. R. Hoffmann, J. Phys. Chem., 1987,91,3789. N. Serpone, D. Lawless, R. Terzian and D. Meisel, in Electro-chemistry in Colloids and Dispersions, ed. R. A. Mackay and J. Texter, VCH, New York, 1992, ch. 30, pp. 399-416. M. Prairie, L. R. Evans, B. M. Stange and S. L. Martinez, Environ. Sci. Technol., 1993,27, 1776. C. S.Turchi and D. F. Ollis, J. Catal., 1990, 122, 178. G. Rothenberger, J. Moser, M. Gratzel, N. Serpone and D. K. Sharma, J. Am. Chem. SOC., 1985,107, 8054. erflier, J. J. M. Binsma and Z. I. Kolar, J. Phys. Chem., 1989, 93, 63 C. Boxall and G. H. Kelsall, J. Chem. SOC., Faraday Trans., 5895. 1991,87,3547. 35 J. M. Warman, M. P. de Haas, P. Pichat, T. P. M. Koster, E. A. van der Zouwen-Assink, A. Mackor and R. Cooper, Radiat. Phys. Chem., 1991,37,433. 64 65 66 U. Kolle, J. Moser and M. Gratzel, Znorg. Chem., 1985,24, 2253. R. F. Howe and M. Gratzel, J. Phys. Chem., 1985,89,4495. D. Lawless, N. Serpone and D. Meisel, J. Phys. Chem., 1991,95, 36 R. W. Fessenden and P. V. Kamat, Chem. Phys. Lett., 1986,123, 233. 37 J. M. Warman, M. P. de Haas, M. Gratzel and P. P. Infelta, Nature (London), 1984,310,306. 38 J. M. Warman, M. P. de Haas, P. Pichat and N. Serpone, J. Phys. Chem., 1991,95,8858. 39 S. Ramo, J. R. Whinnery and T. van Duzer, Fields and Waves in Communication Electronics, Wiley, New York, 1984. 40 K. M. Schindler and M. Kunst, J. Phys. Chem., 1990,94,8222. 41 J. M. Warman, M. P. de Haas and H. M. Wentinck, Radiat. 67 68 69 70 71 5166. 0.I. Micic, Y. Zhang, K. R. Cromack, A. D. Trifunac and M. C. Thurnauer, J. Phys. Chem., 1993,97,7277. A. Henglein, Ber. Bunsenges. Phys. Chem., 1982,86,241. D. Bahnemann, A. Henglein and L. Spanhel, Faraday Discuss. Chem. SOC., 1984,78, 151. M. Gratzel and R. F. Howe, J. Phys. Chem., 1990,94,2566. R. F. Pierret, Semiconductor Fundamentals, Addison-Wesley, New York, 2nd edn., 1989, vol. 1. 42 Phys. Chem., 1989,34,581. D. Bahnemann, A. Henglein, J. Lilie and L. Spanhel, J. Phys. Chem., 1984,88,709. Paper 4/02296C; Received 18th April, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003315

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3323-3330 Time-resolved Microwave Conductivity Part 2.-Quantum-sized TiO, and the Effect of Adsorbates and Light Intensity on Charge-carrier Dynamics Scot T. Martin, Hartmut Herrmannt and Michael R. Hoffmann' W. M. Keck Laboratories, California Institute of Technology, Pasadena, CA 91 125,USA Charge-carrier recombination dynamics after a pulsed laser excitation are investigated by time-resolved micro- wave conductivity (TRMC) for quantum-sized (Q-) TiO, and P25, a bulk-phase TiO, . Adsorbed scavengers such as HNO, , HCI, HCIO, , isopropyl alcohol, trans-decalin, tetranitromethane, and methyl viologen dichloride result in different charge-carrier recombination dynamics for Q-TiO, and P25. The differences include a current doub- ling with isopropyl alcohol for which electron injection into Q-TiO, is much slower than into P25 and relaxation of the selection rules of an indirect-bandgap semiconductor due to size quantization.However, the faster inter- facial charge transfer predicted for Q-TiO, due to a 0.2 eV gain in redox overpotentials is not observed. The effect of light intensity is also investigated. Above a critical injection level, fast recombination channels are opened, which may be a major factor resulting in the dependence of the steady-state photolysis quantum yields on /-'',. The fast recombination channels are opened at lower injection levels for P25 than for Q-TiO,, and a model incorporating the heterogeneity of surface-hole traps is presented.When the radius of a colloidal semiconductor particle falls below the exciton radius (1-10 nm), size-quantization effects appear.'-6 Colloidal TiO, can be prepared in the size-quantized regime by sol-gel synthesis methods7-' Over the last several years, we have been investigating the feasibility of using size-quantized TiO, as a strategy to improve the quantum efficiencies of redox transformations. lo-', In partic- ular, ultra-bandgap illumination (A < 388 nm) of a suspen- sion of TiO, particles results in the stoichiometric oxidation of many chlorinated hydrocarbons to CO, and HCl.'5-20 A generalized photoelectrochemical mechanism involving the major processes can be written as follow^:^ '-', photogeneration: TiO, + hv -+e-+ h+ carrier trapping : >Ti4+-02-H + e-t)>Ti3+-02-H >Ti4+ -02-H + h+ + >Ti4+-0'-H recombination : e-+ >Ti4+ -0O-H + >Ti4+-02-H substrate oxidation : >Ti4+-0'-H + R -+ >Ti4+-02-H + R" substrate hydroxylation : >Ti4+-0'-H + R --r >Ti4+ + +'ROZ-H >Ti4+ + H+02-H-+ >Ti4+-02-H + H+ electron deactivation : >Ti3+-02-H + 0,+ >Ti4+-02-H + 0,' 0,'-+ H+ -+ HO,' Critical steps of the photoelectrochemical mechanism of eqn.(1)-(9) have been investigated by the transient absorp- tion spectra obtained following laser flash photolysis of trans- parent colloidal TiO, .24-31 Most investigators have assumed that conclusions reached regarding colloidal TiO, are applic- able to larger bulk-phase TiO, particles.In the present study, t Present address : Institut fur Physikalische und Theoretische Chemie der Universitat, GH Essen, Germany. the basis of this assumption is investigated by using TRMC measurements, which do not discriminate by crystallite size,10,32-39 to compare the photoelectrochemical mecha- nisms of quantum-sized TiO, and bulk-phase TiO, (Degussa p25).l7--20,40-63 We have recently reported that the relative steady-state quantum efficiencies obtained with these two dif- ferent forms of TiO, depend upon the specific reaction mechanisms involving the electron donors (e.g. direct hole transfer or hydroxyl radical attack).lo Flash photolysis-TRMC experiments are used to compare the charge-carrier dynamics of Q-TiO, and P25 TiO, as functions of the acid used in the sol-gel synthesis, of the charge-carrier scavengers present at the particle interfaces and of light intensity.Experimental Preparation and Characterization Powders of Q-sized TiO, particles were synthesized by the controlled hydrolysis of titanium@) tetraisopropoxide in the presence of HCl, HNO,, or HClO, according to procedures described previously. 7~1 The sol-gel suspension from HClO, was neutralized with NaOH to pH 3.0 in order to obtain a powder. A powder containing methyl viologen (MV) dichlo- ride hydrate (Aldrich) was obtained by rotary evaporation of a resuspension of Ti0,-HNO, mixed with MV. Dry powders of P25 were obtained by the rotary evaporation of suspen- sions (1.44 g 1-', pH 1.5) in HNO,, HCl, HClO, or MV.P25 and Q-TiO, were supported with several drops of trans-decalin (T) except for the preparations containing isopropyl alcohol (ISP) and tetranitromethane (TM), which were pre- pared at the time of measurement by adding several drops to the powders. The Q-TiO, particles were characterized as described pre- viously.' high-resolution transmission electron microscopy (HRTEM) lattice spacing images and X-ray diffraction (XRD) Scherrer line broadening showed the size distribution of the particles was between 2 and 4 nm. The electron diffraction (ED) and XRD patterns of the TiO, particles were character- istic of anatase. A blue shift of the absorption onset to 345 nm (3.6 eV) was observed. The absorption coefficient at 308 nm was 2.9 x lo4 cm-'.Degussa P25 particles, which were composed of 30 nm crystallites consisting of 80% anatase and 20% rutile, aggregated to form 1 pm particle^.^'.^' J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 TRMC Measurements The TRMC apparatus is described in the previous paper." In a typical TRMC experiment separated charge-carriers, which are generated by a laser pulse, lead to a perturbation of the initial microwave absorbance. The temporal decay of the conductivity signal (i.e. microwave absorbance) reflects the lifetime of the photogenerated carriers and is given by the following equation I II I I 1 I I 2800 3200 3600 field/G Fig. 1 EPR spectra obtained at 77 K for sol-gel preparations of size-quantized TiO, in (a) HNO, ,(6)HCl and (c) HClO, where a is the conductivity, eT- is a trapped electron, hT+ is a trapped hole, ecb- is a conduction-band electron, and hvbt is a valence-band hole.The hole terms in eqn. (I) are negli- EPR Measurements gible on the timescales of Fig. 1-4 (later) due to fast hole An E-line Century X-band Spectrometer (Varian, Palo Alto, trapping at immobile surface hydroxy group^.'^ In this case, CA) was used to record the first-derivative EPR spectra at the conductivity is due to the two-electron terms in eqn. (12), 77 K. CuSO, was used as a calibration standard. The TiO, and the conductivity decay corresponds to the recombination samples were prepared as gels (ca. 10 g 1-I) for the EPR [eqn. (4)] or emission of electrons to substrate [eqn.(10)-analysis. (13~ I I I I 112.0 I I I I I 1 8.01 10.0 A > 8.0 Ef 6.0 & 4.0.-v, 2.0 0.0 -2.0 ' I I I I I -2.0' I I I I I -400 0 400 800 1200 1600 -2.0 0.0 2.0 4.0 6.0 8.0 ti me/ns time/ps 4.0I I I I I I 2.0 3.0 1.5 >E 2.0 1.0 ---. cp 2-& 1.0 2 0.5.-fn .-0) 0.0 * 0.0 -0.5' 1-1 .O1-20 0 20 40 60 80I I I I-200 0 200 400 600 800 ti me/ps time/ms Fig. 2 Effect of (a) HNO, ,(6) HCl and (c) HClO, on the conductivity decays of Degussa P25 supported in T. A, 200 ns div- ' timebase; B, 2 ps div-'timebase; C, 200 ps div-'timebase; D, 10 ms div- 'timebase. Table 1 Relative initial carrier concentrations and conductivity decay half-lives relative initial carrier concentration half-life adsorbate P25 Q-TiO, P25 Q-TiO, HC10, 11.2 2.1 750 ns 25 ps 1.6 ms 1.6 ps 800 ps 11 ms HNO, 3.6 3.0 8.0 ps 4.3 ms 2.8 ms 6.3 ms HCl 1.4 1.7 35 ps 8.7 ms 2.8 ps 110 ps 8.3 ms T 4.3 3.0 7.9 ps 5.9 ms 2.8 ms ISP 6.8 1.1 3.5 ps 650 ps 8.7 ms 19 ms TM 4.3 n/a 230 ns 2.2 ps 51 ps n/aMV 5.0 0.9 95 ns 1.1 ps 13 ms J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3325 3.0 2.5 --2.5 2.0 2.0 -0.5 I -0.5'I I I I 1 I I 1 -400 0 400 800 1200 1600 -2.0 0.0 2.0 4.0 6.0 8.0 t ime/ns ti me/p 1.5 I I I I 4 > 1.o 1.51 I E1-0.5 C .-0 a 0.0 (4 I I-0.5 L I II I I -200 0 200 400 600 800 time/ps IFig. 3 Effect of (a)HNO,, (b) HC1 and (c) HC10, on the conductivity ns div-timebase. B, 2 ps div-' timebase.C, 200 ps div-' timebase; D, 10 ms div-' timebase. interfacial electron transfer from the conduction-band: The recombination channel [eqn. (4)] is exhausted for t S-10 ns because all the holes have escaped the particle [eqn. (5) substrate reduction : e-+ Ox +0'-(10) and (6)]." The conductivity decay is then due solely to inter-facial electron transfer to the adsorbed substrate [eqn. (10)-short circuit : e-+ Red'+ -+Red (11) (13)], and the microwave conductivity at 100 ns is due to the interfacial electron transfer mediated by a surface trap: fraction of charge pairs that have not recombined [eqn. (4)] or transferred immobile sites [eqn. (3) and (5)-(13)].'* substrate reduction : >Ti3+-02-H + Ox -+ >Ti4+-02-H + Ox'-(12) Results The EPR spectra of Q-TiO, prepared in HNO, , HClO,, orshort-circuit : HCl are shown in Fig.1. The effects of adsorbed HNO,, >Ti3+-02-H + Red" >Ti4+-02-H + Red (13) HCIO,, and HCl on the conductivity decays of P25 and Q-7.0 7.0 6.0 6.0 5.0 5.0 > 4.05 3.0 & 2.0 4.0 E 3.0 > 1-2 2.0 3 1.0 .g 1.0 0.0 0.0 -1 .o' -400 I 0 I 400 I 800 I 1 1200 1600 -1 -2.0.o I0.0 I2.0 I4.0 I6.0 8.0 time/ns time/p 3.0 2.0 2.5 1.5 2.0 >E 1.5 1----. z 1.0 0.5ul v)Z 0.5 .-0.00.0 I I I I I-0.5 -1.0 -0.5 ' -200 0 200 400 600 800 -20 0 20 40 60 80 time/ps time/ms Fig. 4 Effect of (a) ISP, (6) T, (c) TM and (6)MV on the conductivity decays of Degussa P25. A, 200 ns div-' timebase; B, 2 ps div-' timebase; C, 200 ps div-' timebase; D, 10 ms div-' timebase.J. CHEM. SOC. FARADAY TRANS., 1994,VOL. 90 3.0"3:;j 2.5 2.5 2.0 2.0 >E 1.52 1.5 1-2 1.0 .-g 1.0 0.5 fA 0:5 0.0 0.0 1 -0.5' II I I I I I I I-0.5 I -400 0 400 800 1200 1600 -2.0 0.0 2.0 4.0 6.0 8.0 ti rne/ns ti rne/ps 2.5 -1.51I I I I I I I I 1 2.0 -(4 C 1.o >5 0.5 .G0, 0.0 0.0 I I I I I I I 1-0.5 ' I -0.5 (d) -200 0 200 400 600 800 -20 0 20 40 60 80 tirne/ps ti me/ms Fig. 5 Effect of (a) ISP, (b) T and (d) MV (supported in T) on the conductivity decays of Q-TiO, prepared in HNO, .A, 200 ns div-' timebase; B, 2 ps div-' timebase; C, 200 ps div-' timebase; D, 10ms div-' timebase. 3.0-I 2.5--> 2.0 E 100%21.5 -C c El,'5 1.0-.cr, 2.0 -0.5 0.0-4 2* 0% 1.4% I I I I I I I-0.5' I 0.0-5.0 0.0 5.0 10.0 15.0 20.0 25.0 30.0 35.0 0.0 10.0 20.0 30.0 40.0 ti me/p pulse energy/mJ Fig.6 Effect of the incident laser pulse energy (100%= 4.9 mJ Fig. 8 Effect of the incident laser pulse energy on the initial (0,100 pulse-') on the conductivity decays of Degussa P25 supported in ns) and the residual (m, 15 ps) conductivity of Q-TiO, prepared in trans-decalin on the 2 ps div-' timebase NaClO, and T 0.0 0.0 10.0 20.0 30.0 40.0I pulse energy/mJ pulse energy/mJ Fig. 7 Effect of the incident laser pulse energy on the initial (0,100 Fig. 9 Effect of the incident laser pulse energy on the initial (0,100 ns) and the residual (m, 15 ps) conductivity of Degussa P25 prepared ns) and the residual (m, 15 ps) conductivity of Degussa P25 prepared in NaClO, and T in HNO, and T J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 TiO, are shown in Fig. 2 and 3 while the corresponding con- ductivity decays of P25 and Q-TiO, in the presence of T, ISP, TM and MV2+ are shown in Fig. 4 and 5. The relative initial charge-carrier concentrations (Table l), which are determined at 100 ns, are proportional to the conductivity signal strengths observed in Fig. 2(a)-S(a). The time-resolved conductivity signals of P25 supported in T were found to be a function of the incident pulse energy, as shown in Fig. 6. Similar results were obtained for different adsorbate systems (i.e. ClO,-, and Cl- in T), as shown in Fig.(7)-(9) in which the initial charge-carrier concentrations at 100 ns and 15 ps are plotted as functions of the laser pulse energy (i.e. the charge-carrier injection level). For P25 pre-pared in NaClO, and T (Fig. 7) the carrier concentrations at 100 ns and 15 ps appear to have a lower rate of increase above 4 mJ for a similar increase in injection level. In the presence of NaClO, and T (Fig. 8), the critical injection level for Q-TiO, shifts to ca. 10 mJ, and a second critical level at 25 mJ is also apparent. Samples of P25 prepared in HNO, and T (Fig. 9) appear to have critical injection levels at 5 mJ and lower saturation levels than similar samples prepared in NaClO, (Fig. 7). Discussion Char ge-carrier recombination The fate of the majority of photoexcited charge-carrier pairs in undoped ultra-small particulate semiconductor particles is rapid recombination2' since the instantaneous charge-carrier concentrations from the laser pulse are of the order of 1021 At these concentrations, the higher energy states of the energy bands are populated and symmetry-allowed (i.e.direct) bandgap recombination in TiO, is facilitated. Fast recombination also occurs between free holes and trapped electrons [eqn. (4)].1'*24 Both these channels are exhausted within the time-resolution of our experiment. When the residual carrier-level concentration falls to 10l8 ~m-~, the Fermi level is moved out of degeneracy and the residual con- ductivity is detected by our measurement system.A carrier concentration of ca. 10l8 cm-3 should be representative of the steady-state concentrations of accumulated charge car- riers in CW photolysis experiment^.^, Fast charge-pair recombination rates are further enhanced in size-quantized semiconductor particles owing to the mixing of states that relaxes the selection rules for an indirect This latter effect of size quantization is clearly illustrated in the data summarized in Table 1. For all sorbates (electron acceptors and electron donors) the concen- tration of charge-carriers is lower in Q-TiO, than in P25 after 100 ns. These experiments suggest that indirect-bandgap semiconductors in the Q-sized domain are less photoefficient redox catalysts owing to inherently faster rates of charge-carrier recombination.This photophysical effect offsets the predicted gains associated with higher thermodynamic driving forces for interfacial electron transfer. Surface Structures The EPR spectra shown in Fig. 1 suggest the formation of radical species on the surface of the TiO, particles. The species may include >Ti"'-Cl, >Ti'"-ONO, , >Tirv-C1' and >TilV-ONO,'. However, the radicals present in the sample constitute only 1% of the atoms, as indicated by a comparison of the integrated intensity of the EPR signals to the CuSO, standard. The observed radicals suggest that the acid anions interact with the surface of the particle and that the bulk of the anions speciate as >TiW-Cl, >TitV-0NO, and >Tirv-OC1O,.Fluctuating energy Level Model The observed differences as a function of added acids (Fig. 2 and 3) suggest that photogenerated charge carriers undergo interfacial charge transfer to the acid anions on the surface of the TiO, particles. The kinetics of the interfacial charge transfer can be understood in terms of the fluctuating-energy model proposed by Gerischer. 70 The one-electron oxidation potentials of the acid adsorbates under standard conditions are as follow^:^' C1-*--) C1' + e-; E" = -2.6 V (14) NO,--NO,' + e-; E" = -2.3 V (15) OH-*--)'OH + e-; E" = -1.9 V (16) Although the one-electron oxidation potential for perchlorate is not known, we assume that the value lies between 2.0 and 3.0 eV based upon comparison with eqn.(14) and (15). The oxidation potential of the surface-bound hydroxide anion [ie. eqn. (3)] can be estimated also from eqn. (17); it has been reported to be -1.5 V." Thus, the oxidation of the adsorbed acids can take place only from direct valence-band hole transfer since the E:,.. (*OH-acid anion) value <O (i.e. AGO 9 0). The corresponding two-electron reduction potentials for nitrate and perchlorate are as follows:72 NO3-+ 2e-+ 3H+++HN02+ H,O; E" = +0.9 V (17) C10,-+ 2e-+ 2H+++ClO,-+ H,O; E" = +1.2 V (18) Within f0.2 V, a conduction-band electron has a reduction potential of -0.3 V [eqn. (10) and (ll)] and a valence-band hole has a reduction potential of 2.9 V [eqn. (3)J at pH 0. In these calculations, we used the constraints that the conduction-band edge for TiO, is 0.1-0.2 V negative of the flatband potential and that the bandgap of anatase is 3.2 eV.The flatband potential has been determined to be -0.127 and -0.2 V7, us. NHE at pH 0 for colloidal TiO,. The size- quantization effects in Q-TiO, increase the bandgap by 0.4 eV.' This increase corresponds to a shift in the reduction potential of the conduction-band electron to -0.5 V and of the valence-band hole to +3.1 eV for Q-TiO, at pH 0. Since Degussa P25 is predominantly anata~e,~' its band edges are at -0.3 and 2.9 V at pH 0. The apparent reduction potentials of trapped charge car-riers are difficult to estimate. In the case of a trapped elec- tron, Ti"' is 0.2-0.5 eV negative of the conduction-band edge at zero surface charge, E&, of r~tile.~,-'~ Using EEb = -0.3 V us.NHE and pH,,, = 5.8,77 the reduction potential of a trapped electron [eqn. (8), (12) and (13)] is ca. -0.25 V us. NHE. Because the trapped electron is localized, the reduction potential is the same for both Q- and P25 TiO, . The trap is 50 mV below E,, of P25, in agreement with other sugges- tion~,~~and is 250 mV below E,, of Q-TiO,. The reduction potential of a trapped hole [eqn. (6) and (7)] is +1.5 V, as discussed for eqn. (16). The acid anions (Cl-, NO3- and C104-) are chemisorbed onto the surface of TiO, by ligand substitution of surface hydroxy groups and by electrostatic attraction at pH < 6.8. The one-electron oxidation potentials [eqn.(14)-( 16)] indi- cate that the acid anions at the TiO, surface are thermody- namically less capable of hole capture than surface hydroxide but may still serve as good hole traps [eqn. (3)] owing to the high overlap between the energy levels of the acid anions and the valence-band hole for 0.5 < A/eV < 1.0 where A is the reorganization en erg^.^',^^ Although OH- and C1- cannot be further reduced, NO, -and ClO,- undergo multi-electron reductions [eqn. (17)and (18)]and thus may serve as effective electron traps. Inorganic Donors and Acceptors Interfacial charge transfers to the acid anion adsorbates occur on the ps and short ns timescales and compete with interfacial recombination processes. This effect is clearly seen in Table 1, in which the initial charge-carrier concentrations appear to be unique for each adsorbate.In the case of HCl, the chemisorbed C1- anion directly scavenges valence-band holes [eqn. (19)] within 10 ns in direct competition with surface OH- [eqn. (3)]:26 >Ti4+-C1-+ h+ + >Ti4+-C1' (19) After formation, the surface chloride radical may open a second channel for charge-carrier recombination as follows : e-+ >Ti4+-C1* + >Ti4+-C1-(20) The relatively low carrier concentrations observed at 100 ns indicate that the cross-section for electron capture by a surface chloride radical [eqn. (20)]is greater than that for a surface hydroxy group [eqn. (4)J This observation could be explained, in part, by the relative differences in electronega- tivity for OH- and C1-.Since nitrate and perchlorate serve as both potential electron scavengers and hole scavengers, three scenarios are possible to explain their effects: (a) >Ti'V-ONO, and > Ti'V-OC1O, could act as fast hole scavengers; (b)they could act primarily as fast electron scavengers or (c) they could act equally with respect to the trapping of both charge types. In case (a)fast hole trapping exhausts the reservoir of electron- accepting species and the conductivity signal is proportional to the concentration of electrons. Case (b)is the opposite of case (a), i.e. the conductivity signal is due to free holes. In case (c) electrons and holes are equally separated on the surface of the semiconductor particle while the charge carriers remain- ing within the particle recombine rapidly and thus no residual conductivity is observed at 100 ns.As shown in Fig. 2, case (c) does not appear to occur. Thus, preferential trap- ping of one carrier type as in (a) or (b)appears to take place. Although processes (a) and (b)are not directly distinguishable in our experiment, we will attempt to show that interfacial electron transfer does not occur significantly within the time- resolution of the experiment (tide infra). In this case, the con- ductivity signal should be due to electrons as in (a). The relatively high electron-charge-carrier concentrations shown in Fig. 3(a)for P25 in the presence of HClO, suggest that the cross-sections for hole capture by >Ti'V-OC1O, or >TirV-0H are comparable while >Ti'V-OC103' has a lower cross-section for electron capture than >Ti"-OH'.In the case of Q-TiO,, the electron concentration in the pres- ence of HClO, shows that the cross-section for electron capture by >Ti'V-CIO,' increases relative to P25. The lower charge-carrier concentrations observed in the presence of HCl for P25 as compared with Q-TiO, indicate that valence-band hole-capture rates of >TiIV-C1 [eqn. (19)]are greater for P25 than for Q-TiO,. In this case, more carriers are short-circuited by conduction-band electron transfer to >Ti'V-CI' [eqn. (20)]for P25 than for Q-TiO,, and thus the residual charge-carrier concentration at 100 ns is greater for Q-TiO, than for P25 (Table 1).Organic Donors and Acceptors TM and MV2+ are known to serve readily as electron scav- engers in TD, system^',^^,^^ while T, a saturated bicyclic hydrocarbon, and ISP serve as surface hole scavenger^.^.^^*^^ J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 The measured charge-carrier concentrations at 100 ns in P25 are comparable in the presence of MV2+, TNM, ISP and T. If fast interfacial electron transfer were taking place to TM and MV2+, then the carrier concentrations at 100 ns would be expected to be reduced. The first half-lives for the charge carriers in the presence of MV2+, TNM, ISP and T are 95 ns, 230 ns, 3.5 ps and 7.9 ps, respectively. Thus, we conclude that, even in the presence of fast electron acceptors, interfacial elec- tron transfer takes place primarily at timescales 2 100 ns.In the presence of electron acceptors (TNM and MV2+), the conductivity decays flatten on the ms timescale, as shown in Fig. 4(b) and (c) and Fig. 5(b) and (c); this effect suggests that electron transfer does not take place on this timescale. Electron transfer to TNM and MVZ+ takes place during the first 500 ns, as shown in Fig 4(a) and 5(a).After that time, the surface substrate is exhausted and further electrons do not escape the particle until fresh substrate diffuses to the surface, as shown in Fig. 4(d) and 5(4.''The increasing charge-carrier concentrations seen in Fig. 5(a)and (b)suggest that back electron injection from ISP into Q-TiO, takes place. The current-doubling effect occurs when a one-electron oxidant such as ISP produces a radical that can inject an electron into the conduction band of TiO, as follows:7 h+PH PH CH,-C-CH, -CH,-C-CH,-H+IH 0-CH,-C-CH, II -e-,H+ This effect is absent in Fig.4(a) and (b)for P25.However, the highest initial conductivity signal for P25 is observed when ISP is the electron donor, which supports the argument that back electron injection takes place much more quickly than in Q-TiO,. A slower rate for Q-TiO, may be due to the 0.2eV shift to the negative of the conduction-band edge of Q-TiO, , which should slow electron injection. Slower elec- tron injection rates into Q-TiO, may also arise from a drop in the electron-transfer probability due to a higher surface density of hydroxy groups that sterically hinder surface Ti", which is the likely site for electron injection into the conduc- tion band.The sol-gel preparation method is expected to result in higher surface densities of hydroxy groups for Q- TiO, than for P25,which is prepared by flame hydrolysis. Light Intensity Effects The observed decay rates from 100 ns to 35 ps (Fig. 6)follow apparent first-order kinetics ;this observation implies that the charge carriers act independently of one another at t d 35 ps. However, the charge-carrier concentrations at 100 ns are not directly proportional to the initial charge-carrier injection levels because fast recombination is a second-order process. The observed carrier concentrations at 100 ns give us a measure of the extent of charge-carrier recombination at t < 100ns.The charge-carrier injection level can be determined from the number of photons injected into the penetration depth of Q-TiO, at 308 nm: I (~m-~) I (mJ) a (cm-') = 2.3 x 1017hv (eV) A (mm2) (23) where I (cm-,) is the injection level, I (mJ) is the pulse energy, a is the absorption coefficient, A is the cross-sectional J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 area, and hv is the photon energy. In our flash photolysis system, a 4.5 mJ pulse results in ca. 10,' carriers ern-,. The slopes of the lines in Fig. 7-9 can be defined as response factors. The response factor is the relative change in the carrier concentration at a time, t, produced by a change in the injection level.In the absence of higher order (i.e. fast) recombination processes, the response factor should be high while the response factor is zero at saturation. The injection level at which the response factor changes is defined as the cross-ouer threshold. These thresholds (Table 2) are the injec- tion levels at which new higher order deactivation channels are opened and they appear as discontinuities in the lines in Fig. 7-9. The response factors in Table 2 decrease as the injection level increases, which indicates that higher injection levels result in larger quantum yields for fast recombination. In a typical CW experiment, higher light intensities are expected to raise the quantum yields for fast recombination and thus to reduce the net quantum yields for substrate oxidation and reduction.This conclusion provides evidence for our hypoth- esis that the quantum efficiencies for oxidation and reduction in CW photolysis fall as the square-root of absorbed light intensity (acc I-'I2) owing to rapid recombination at higher light intensitie~.'~ The apparent discontinuities at the cross-over thresholds in Fig. 7-9 may be explained in terms of a mechanism involving two chemically distinct hole-trapping sites, S, and S,, in TiO, that have hole-trapping cross-sections CT, and 0, .21 S, is most likely OH,' formed on a surface TiIV site and S, is OH,' formed on a surface p-0 site. We propose that crl $ 6, due to a potentially unfavourable charge balance on p-0.The concentrations of S, and S, following an injection level of [So] are subject to the following constraints: where [S,] and [S,] are the concentration of occupied hole traps. At the time of injection, [S,](O) + [S2](0) = [So]. Due to stoichiometric annihilation, [C](t) = [S,](t) + [S,](t) for t -g qhtwhere qhtis the time for interfacial hole transfer and [C](t) is the concentration of electrons. The time-dependent decay of [S,] and [S,] may be model-led as follows : where kl, ,are the cross-sections for charge-carrier recom- bination [eqn. (4)] and the response factors, m,,,, are exp[ -kl, ,(lo0 ns)]. A plot of the conductivity signal [C](lOo ns) us. the injection level, [So], under the conditions set out in eqn.(11) results in two straight lines with a discontinuity at the cross-over threshold where [So] = S,. The higher surface area and surface hydroxylation of Q-TiO, results in cross-over thresholds at an injection level five to ten times higher than P25, as shown in Table 1. The effect of HNO, and HCIO, on the surface states of Degussa P25 is shown by the difference in cross-over thresholds and response factors. Saturation of the response factor for HNO, indicates that k, is fast. Q-TiO, has a second cross-over threshold, which indicates a third possible site for trapped holes. Conclusions We have shown that the charge-carrier dynamics of Q-TiO, and P25 respond differently with respect to added electron donors and acceptors and to absorbed light intensity. These effects suggest significant differences in the photoelec-trochemical mechanisms for Q-TiO, and P25.Anions such as C1-affect the charge-carrier recombination processes by introducing surface states through specific adsorption as follows: >Ti-OH + C1-+ Hf >Ti-Cl + H,O (27) Hence, the photoreactivity of Q-TiO, could be enhanced by deactivation of fast surface recombination processes. Wide bandgap quantum-sized semiconductors such as ZnO and TiO, should be moving further into the Marcus inverted with respect to the net driving force for oxidation. However, when quantum efficiencies are limited by the rate of reduction, size quantization has been shown to yield high quantum yields for H,O, production on ZnO" or alkene reduction on Ti0,.82 In the case of pollutant oxida- tion on TiO,, quantum efficiencies may be limited primarily by the effective rate of interfacial reducti~n.~~*~~ However, the mixing of states in the size-quantization regime appears to enhance direct electron-hole pair recombination in indirect-bandgap semiconductors (i.e.TiO,). In the size-quantization regime of TiO, ,faster electron-hole pair recom- bination may offset the increased driving force for interfacial reduction. We are indebted to Prof. Nathan S. Lewis for the loan of microwave components and to Prof. Geoffrey A. Blake for the use of the excimer laser. We are grateful to ARPA and ONR {NAV 5 HFMN N0001492J1901) for financial support. S.M. is supported by a National Defence Science and Engineering Graduate Fellowship.H.H. wishes to thank NATO/DAAD for financing a research visit at the California Institute of Technology. Wonyong Choi, Dr. Amy Hoffman, Nicole Peill and Dr. Andreas Termin provided valuable support and stimulating discussion. Table 2 injection level cross-over threshold /lo-,' cm-, system 100 ns 15 ps P25 in NaClO, < 3.6 < 3.6 and T > 3.6 > 3.6 sol-gel Q-TiO, in NaClO, and T -= 14 <25 <8.6 <25 >25 > 25 P25 in HNO, < 6.0 < 6.0 and T >6.0 >6.0 Injection level cross-over thresholds and response factors Ainjection level 100 ns 15 ps 2.4 0.7 0.6 0.2 0.24 0.13 0.12 0.05 0.03 0.003 0.40 0.23 0.03 0.01 response factor Acarrier concentration 3330 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 References 44 C. Y. Hsiao, C. L. Lee and D. F. Ollis, J. Catal., 1983,82,418. 45 M. Barbeni, E. Pramauro, E. Pelizzetti, E. Brogarello, M. 1 D. W. Bahnemann, Zsr. J. Chem., 1993,33,115. Gratzel and N. Serpone, Nouu. J. 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ISSN:0956-5000    

DOI:10.1039/FT9949003323

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3331-3333 X-Ray Investigations on Liquid-crystalline Copolysiloxanes for Second-harmonic Generation Erik Wischerhoff and Rudolph Zentel lnstitut fur Organische Chemie , Jd-Becher- Weg 18-20, Johannes-GutenbergUniversitat, 0-55099 Mainz,Germany Ha rtmut Fischer*t H. H. Wills Physics Laboratory, Royal Fort, TyndalI Avenue, Bristol, UK BS8 ITL Ferroelectric liquid-crystalline copolymers of chiral mesogens and non-chiral comonomers have been described and their mesophase properties have been investigated by X-ray diffraction. Tilted smectic mesophases are obtained for a chromophore content of up to 50 mol%. Liquid-crystalline materials exhibiting a chiral smectic C* phase and some other tilted smectic phases with chirality are known to have a helical structure.However, this helix can be unwound by surface effects and electric and magnetic fields. Thus a macroscopic polar structure with ferroelectric proper- ties is obtained.'*2 There is a wide variety of ferroelectric liquid-crystalline compounds, some having low molar mass and others which are p~lymeric.~-~ These materials have attracted interest for display applications and for their piezoelectric properties, in the case of crosslinked systems.' The polar structure in the helix-unwound state can also be useful for second-order non-linear optical effects like second- harmonic generation (SHG), since these effects rely on a non- centrosymmetric structure in the bulk. Compared with poled polymers, which are often used for this purpose, chiral smectic C* materials have the advantage of a stable polar structure.In contrast to crystals, polymeric liquid crystals have the advantage of being liquid-like, so they can be pro- cessed into any required shape. The potential of chiral smectic C* compounds for non- linear optics was recognised several years ago,* but the first experiments produced disappointing results. Although they showed that SHG was possible, the eficiency was rather poor (d values of about 0.01 pm V-'). This problem was caused by the chemical constitution of common liquid-crystalline materials. In order to obtain second-order non-linear suscep- tibilities in organic materials, long electron-donor-acceptor n systems are needed on the molecular level.However, in most cases liquid crystals do not possess strong donors and accep- tors fixed to their aromatic core in the direction of their polar axis, which is lateral to the long axis of the molecules. It was our attempt to combine the good polar order of the chiral smectic C* phase with a suitable molecular structure for SHG by preparing copolymers of mesogens with a strong tendency to exhibit an S,* phase and chromophores with a donor-acceptor .n system lateral to their polar axis (NLO chromophores). The synthetic route to these materials and first non-linear optical (NLO) experiments are reported else- where.g The intensity of the second-haimonic signal in the sz phase is roughly ten times as high as in the isotropic phase, clearly demonstrating the influence of the polar structure on this effect.We present a detailed X-ray investigation of the phase properties of a selected copolymer system with regard to their dependence on the amount of NLO chromophore incorpor- t Present address: Dept. of Chem. Eng., Polymer Chemistry and Technology, Technische Universiteit Eindhoven, PO Box 5 13, 5600 MB, Eindhoven, The Netherlands. ated. The influence of the NLO chromophores on the phase behaviour of a material is of outstanding importance to its application. Second-harmonic generation can be expected only if the S,* phase is not disturbed by the lateral substit- uents of the chromophores. Materials The copolysiloxanes described here have the general formula shown below and were prepared according to ref.9. All materials are statistical copolymers. (X+ Y):Z=1:2,7:la-lf;(X+ Y):Z= 1:1:2a The compositions and molar masses of the materials are given in Table 1. Table 1 Compositions and molar masses polymer X Y molar mass (GPC) ~~~~~~ ~ ~ ~ la 100 0 27 400 lb 70 30 28 O00 lc 50 50 28 220 Id 30 70 28 800 If 20 80 28 750 2a 70 30 21 700 ~~~ ~~~ The GPC values represent the peak maximum of the GPC curve. Polystyrene standards were used as reference. Structural Investigations X-Ray studies were carried out using an Elliot GX 21 rotating anode generator with a copper target equipped with a flat graphite crystal monochromator and a collimator comprising two pairs of crossed slits.The detection device was a Siemens X-1000 area detector coupled with a PC (GADDS software) for collecting, analysing and storing X-ray images. Tem- perature control of the samples was maintained with a LINKAM hot stage. Samples of the copolysiloxanes were sealed in glass capillaries for temperature-dependent studies. Oriented samples were prepared by pulling fibres out of the anisotropic melt. Fig. 1 shows the X-ray diffractograms of the copolymers lb and Id. All the X-ray diffraction patterns show very similar reflections, as described by Poths and Zentel," and may be interpreted in a similar way. In the X-ray pictures of oriented samples the occurrence of microphase separation' '9' between layers of the polysiloxane main chain and the meso- genic side group can be observed clearly. The reflections cor- responding to the mesogen layers are marked ML and the reflections corresponding to the polysiloxane main chain are marked SOL in Fig.1. In both X-ray patterns, an orienta- tion of the mesogen layer is clearly visible, however, the poly- siloxane layer at 7.5 8, remains unoriented. Thus, a phase separation as found by Diele et d." and Sutherland et a1.12 must occur. X-Ray diffraction studies on copolysiloxanes with different amounts of dimethyl siloxane units in the main chain show similar results, as described by Poths and Zentel." When studying ferroelectric LC polysiloxanes without chromophores they have found an increase of the layer thickness of the smectic layers of about 4 8, per inserted dimethyl siloxane unit.This value is in good agreement with Diele's studies and theoretical predictions deduced from atomic dimensions (3.78 A). In our investigation, similar values are observed. The phase structure is not changed with a change in the ratio of functional units to dimethyl siloxane units. This is also an indication of microphase separation. Since the mesogen layers are well separated from the poly- siloxane layers, no change in the phase behaviour is to be expected. The microphase separation instead favours the for- mation of smectic layers. The change from tilted to orthogonal smectic phases can be observed in the X-ray diffractograms by the change of the angle between small-angle reflections with chromophore con- tents.For 2 70 mol% of chromophore, the intensity maxima of the two reflections are 90" apart, indicating orthogonal smectic mesophases. Table 2 gives the mesophases and structural data of the investigated copolymers. There is a significant change at 80 mol% of chromophore. Materials with lower amounts of chromophore show a layer structure and higher ordered smectic phases below their Sz or S, phase. The material with the highest chromophore content only has a smectic A phase; the layer spacing can only be interpreted by assuming an ML ML Fig. 1 X-ray diffractions of the coploymers Ib and Id J. CHEM. SOC. FARADAY TRANS., 1994, VOL,. 90 Table 2 Mesophases and structural data mesogen polymer phase layer spacing/A distance/l( 48.4 4.67 49 4.60 50 4.65 49 4.60 49 4.65 49 4.60 55 4.67 48.4 4.53 55 4.60 71.6 4.53 38 4.63 interdigitated arrangement of the mesogens, which may be induced by charge transfer interactions or strong dipole- dipole interactions of the chromophore groups.The investigations by X-ray reveal smectic mesophases in all copolymer systems up to 80 mol% of chromophore. This finding is remarkable since lateral substituents fixed to the rigid core of liquid crystal molecules are known to hinder strongly the formation of LC phases. In this case we have 80% of the moieties laterally substituted on both sides of the molecule, but a broad mesophase is still present.Tilted mesophases fulfilling the symmetry requirements for a polar structure can be observed up to 50 mol% of chromo- phore. However, approaching very high chromophore con- tents, only smectic A phases are found. In Table 2, the layer spacings of the different materials at different temperatures are listed. The results from X-ray investigations are fully consistent with texture observations under the polarisation microscope and with ferroelectric switching experiments. As a result, the dependence of the mesophase sequence on the chromophore content can be described by the phase diagram shown in Fig. 2. Conclusions The X-ray investigations performed reveal the influence of laterally substituted NLO chromophores on the phase sequence of ferroelectric liquid-crystalline polymers.Tilted smectic mesophases, which are essential for a polar structure needed for SHG, are obtained for a chromophore content of up to 50 mol%. At higher contents of laterally substituted chromophore, the tilted phase disappears, but a broad S, phase remains present. E.W. and R.Z. wish to thank the Deutsche Forschungsge- meinschaft for financial support. I160 ,,,],,,, I,,I I ,,,I,,, ,I,,,, 140 I 120 100Yh' 80 :: 20 I 0 I0 20 30 40 so 60 70 80 mol% chromophore Fig. 2 Phase diagram of the copolymer system J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3333 References 8 A. Taguchi, Y.Ouchi, H. Takezoe and A. Fukuda, Jpn. J. Appl.Phys., 1989,28, L997. 1 R. B. Meyer, L. Liebert, L. Strzelecki and P. Keller, J. Phys., 9 E. Wischerhoff, R. Zentel, M. Redmond, H. Coles and 1975,36,69. 0.Mondain-Monval, Macromol. Chem. Phys., 1994,195, 1593. 2 N. A. Clark and S. T. Lagerwall, Appl. Phys. Lett., 1980,36,899. 10 H. Poths and R. Zentel, Liq. Cryst., 1994,16, 749. 3 Polymers for Advanced Technologies 3, ed. M. Lewin, M. Jaffe, 11 S. Diele, S. Oelsner, F. Kuschel, B. Hisgen, H. Ringsdorf and J. H. Wendorff and E. Tsuchida, Wiley, Chichester, 1992, vol. 5, R. Zentel, Makromol. Chem., 1987,188,1993. pp, 195-276. 12 H. H. Sutherland, Z. Ali-Adib, B. Gasgous and G. Nestor, Mol. 4 H. Poths, A. SchonfeldFR. Zentel, F. Kremer and K. Siemens- Cryst. Liq. Cryst., 1988, 155,327. meyer, Adv. Muter., 1992,4,351. 5 M. Dumon, H. T. Nguyen, M. Mauzac, C. Destrade and H. Gasparoux, Liq. Cryst., 1991, 10,475. 6 J. Naciri, S. Pfeiffer and R. Shashidar, Liq. Cryst., 1991,10, 585. 7 H. Kapitza, R. Zentel, H. Poths, S-U. Vallerien, F. Kremer and Paper 4/03514C; Received 13th June, 1994 E. W. Fischer, Makromol. Chem. Rapid Commun., 1990,11,593.

ISSN:0956-5000    

DOI:10.1039/FT9949003331

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3335-3339 Thermal Behaviour and Physico-chemical Characterization of Synthetic and Natural Iron Hydroxyphosphates Dominique Rouzies, Jean Varloud and Jean-Marc M. M. Millet lnstitut de Recherche sur la Catalyse, CNRS, Associe 4 l'universite Claude-Bernard, Lyon I, 2 Avenue A. Einstein, F-69626 Villeurbanne Cedex, France The thermal behaviour of five hydroxyphosphates, lipscornbite [Fe,(PO,),(OH),], barbosalite [Fe,(PO,),(OH),], giniite [Fe,(PO,),(OH), * 2H,0], rockbridgeite Fe,(PO,),(OH),] and Fe,(PO,),(OH), , has been studied. Ther- mogravimetric analysis revealed the hydroxyphosphates to have distinct behaviours : some underwent dehy- droxylation between 623 and 893 K like other phosphates [lipscombite, barbosalite, Fe,(PO,),(OH),], while rockbridgeite and giniite differed.Rockbridgeite, which appeared to contain only iron(iii) ions, underwent a first partial dehydroxylation at only 438K, while giniite totally dehydroxyated only at between 623and 823K. Compari-son of the results obtained by thermogravimetric analysis with those obtained by other techniques, including chemical analysis, X-ray diffraction and Mossbauer spectroscopy, allowed us to propose the existence of a new solid solution for the giniite compound and to confirm that previously observed for rockbridgeite. The solid solution of the former corresponds to a variation in composition due to dehydroxylation-hydration : Fe~'~,Fe~'~,(PO,),(OH),~,(2+x)H,O with 0 <x < 1. The latter corresponds to a variation in the composition due to dehydroxylation-oxidation : Fe~'~,Fe~'~,(P0,)3(0H),_,0,with 0 <x < 1.In both cases, Mossbauer spectros- copy indicates that Fe2+ or Mn2+ substituents may be localized on the different crystallographic sites of the structures. A recent study' has shown that several hydroxyphosphates are potential oxidative dehydrogenation catalysts. For cata- lytic reactions conducted between 573 and 723 K, it appeared of interest to study the thermal behaviour of these com-pounds in this temperature range. This paper reports the thermal and thermogravimetric analyses conducted in argon of four synthetic and one natural hydroxyphosphate: the two polymorphic forms of Fe,(PO,),(OH), , (i) lipscombite, (ii) barbosalite, (iii) Fe,(PO,),(OH), , (iv) Fe,(P0,)4(0H)2 -2H20 called giniite and (v) Fe,(PO,),(OH),, a natural sample called rockbridgeite.Although several published studies have dealt with these hydroxyphosphates, which are well known mineral compounds, none of them have reported in detail their thermal behaviour in a neutral atmosphere. Experimental Sample Preparation Fe,(PO,),(OH), ,lipscornbite and barbosalite samples were prepared as described previou~ly.~.~ They were obtained by hydrothermal synthesis using vivianite [Fe,(PO,), * 8H20] and an amorphous iron(@ phosphate gel as precursors. The hydrothermal syntheses were conducted at 473 K for 5 h. The samples obtained were filtered, washed with water and dried at 373 K overnight.The giniite sample has been prepared by hydrothermal synthesis using a mixture of 5 g of synthetic vivianite Fe,(PO,), -8H,O along with 5 g of phosphoric acid and 1 ml of distilled water per g of solid. The hydrothermal treatment was conducted at 383 K for 120 h. The solid obtained was washed with water and air-dried at 373 K. To our knowledge, this is a novel method for giniite synthesis. The rockbridgeite was kindly supplied by the Museum National &Histoire Naturelle (Paris, France). The chemical analysis of this mineral showed that it contained a relatively large amount of manganese (Table 1). Sample Characterization The chemical composition of the solids was determined by atomic absorption. X-Ray diffraction analysis was performed on the samples using a Siemens D500 diffractometer and Cu-Ka radiation.Differential thermal analyses (DTA) and thermogravimetric (TG) analyses were performed simulta- neously using a SETARAM TGA-DTA 92 thermobalance coupled to a Balzers 420 QMC mass spectrometer. 30-60 mg of the samples were placed in a platinum crucible suspended from one arm of the balance. The analyses were conducted at atmospheric pressure under a deoxygenated and dehydrated argon flow (1 1 h-') with a heating rate of 5 K min-' and a temperature limit of 1023 K. Mass spectrometry was used to verify the nature of the eventual gaseous reaction products, namely 02,H, , H,O, CO and CO,; the precision of the thermogravimetric analyses was ca. 2%. Mossbauer spectra were recorded at room temperature, using a 2 GBqS7Co-Rh source and a conventional constant acceleration spectrometer, operated in triangular mode.The samples were diluted in A120, in order to avoid a too high Mossbauer absorption, and pressed into pellets. The isomer shifts (6) were given with respect to a-Fe and were calculated, as the quadrupole splittings (A) and the linewidth (W), with a precision of ca. 0.02mm s-'. Results Physico-chemical Characterization The main characteristics of the solids studied are presented in Table 2. The P :Fe ratios were calculated from chemical analyses and the values of Fe3+ :(Fe3++ Fe2+) were deter- mined from Mossbauer spectroscopy. The values obtained Table 1 Chemical analysis of the natural rockbridgeite sample ~~ element ~~ ~ wt.% Fe 30.2 P 13.0 Mn 5.5 A1 0.2 Ca 0.1 Mg 0.06 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Main characteristics of the hydroxyphosphates studied ;theoretical values are given in parentheses compound P :Fe Fe3+:(Fe3++ Fe2+Y Fey'Fe"(PO,&(OH), barbosalite 0.69 (0.66) 0.66 (0.66) Fe~1Fe11(P0,)2(OH)2 lipscombite 0.69 (0.66) 0.65 (0.66) Fe!?(PO,),(OH), 0.76 (0.75) 1.00 (1.00) Fey'Fe"(PO,),(OH), .2H20 giniite 0.80 (0.80) 0.65 (0.80) Fe~'Fe"(PO,),(OH), rockbridgeite 0.64 (0.60)b 0.87 (0.80)b a Determined by Mossbauer spectroscopy. In the case of rockbridgeite P :(Fe + Mn) and Fe3+ :(Fe3++ Mn2++ Fe2+) are given. for the P :Fe ratios were in good agreement with the theo- retical ratios.The calculated values for the Fe3+: (Fe3++ Fe2+) ratios corresponded to the theoretical ones except for the giniite and rockbridgeite samples which both indicated an excess of iron(@ cations. The X-ray dif- fraction spectra all revealed pure phases.' Tbermogravimetric Analyses Fe~'Fen(PO,),(OH), ,Barbosalite and Lipscombite The weight loss of the samples occurred at 715, 802 and 823 K for lipscombite and at 743, 793 and 868 K for barbosalite [Fig. l(a) and (b)] (Table 3). The mass spectrometric analyses performed simultaneously allowed the mass losses to be attributed to the dehydration of the solids. The mass losses of the lipscombite and barbosalite are 0.97 and 1.05 mol, respec- tively, per mol of solid.These values are very close to the theroetical values, corresponding to a total content of water, of 1.00. In both cases the dehydration process began at ca. 623 K, the first weight losses were the most important and corresponded approximately to half of the water contained in the compounds. The other half corresponded to the two final smaller water departures and seemed to be equal. Results from DTA analyses have not been presented in all cases. They confirmed the weight losses characterized by poorly defined endothermic peaks. After heating to 923 K, the X-ray diffraction patterns of the samples were totally modified, showing a complete transformation of the solids after dehy- dration. Fe',"(PO,),(OH), The thermogravimetric analysis of this compound showed three weight losses at 802, 818 and 849 K [Fig.l(c)] (Table 3). These weight losses, which were attributed to water depar- tures, corresponded to 1.54 mol of water per mol of solid. This value was again very close to the theoretical value of 1.50. The first two water departures, which were very close in temperature, corresponded in total to ca. 1mol of water. Giniite, Fe',"FeU(PO,),(OH), 2H20 Thermogravimetric analysis of this compound showed three water departures at 473, 658 and 773 K [Fig. 2(a)].The first weight loss, which was the most important, corresponded to + Q,C 2-0.0 --0.1-:-I-2.0 400 500 600 700 800 900 T/K Fig. 1 DTG and TG curves of the two polymorphic forms of Fe','1Fe1'(P0,)2(OH)2,namely (a) barbosalite and (b)lipscombite and (4 Fe,(PO,),(OH) 3 Table 3 compound Fe~1Fe11(P0,)2(OH)2 Fe~'Fe"(PO,),(OH),.2H20 Fe~'Fe"(PO,),(OH), temperature/K barbosalite 743 793 868 lipscombite 715 802 843~~ 802 818 849 giniite 573 633 763 rockbridgeite 434 658 823 Details of the DTA and TG analyses of the hydroxyphosphates loss of water /mol of H20 mol-' ~~ 0.54 0.24 0.27 0.48 0.24 0.24 1.05 0.49 2.60 0.60 0.20 2.20 total loss of water" /mol of H20 mol-' 1.05 (1.00) 0.97 (1.00) 1.54 (1.50) 3.20 (3.00) 2.40 (2.50) a Theoretical values are given in italic.J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 2.6 mol of water per mol of giniite and should correspond to the hydration water (Table 3).The two other water depar- tures, which took place between 623 and 827 K, should corre- spond to the departure of the hydroxy groups. These weight losses correspond to 0.6 mol of water per mol of giniite. This value was lower than that expected from the stoichiometry, whilst that corresponding to the first loss was higher. Rockbridgeite Fe~'Fe''(PO,),(OH), The thermogravimetric analysis of the rockbridgeite sample showed three water departures [Fig. 2(b)]. The first occurred at 434 K and the other two at much higher temperatures, 658 and 823 K (Table 3). Of these last two, the first was sharper and more significant than the second. The first departure cor- responded to 0.20 mol of water per mol of rockbridgeite and the others to 2.20 mol of water per mol of rockbridgeite, which gave a total of 2.40 per mol of rockbridgeite.Miissbauer Spectroscopic Study The Mossbauer spectroscopic analyses of the barbosalite, lip- scombite and Fe,(PO,),(OH), samples, have been p~blished.~.~The results obtained were in good agreement with both the results of the thermogravimetric and chemical analyses. The Mossbauer spectrum of the giniite sample, which is presented in Fig. 3, shows four doublets (Table 4). These doublets, two of which correspond to iron(II1) cations and the other two to iron(@ cations, could not be directly attributed since the detailed crystallographic structure of the compound is not known. However, it was observed that the Fe3+:(Fe2++ Fe3+) ratio was equal to 0.65 instead of 0.80 as deduced from the stoichiometry. Characterization of the sample of rockbridgeite by Moss- bauer spectroscopy showed that the sample contained only iron(II1) cations (Table 4).The spectrum presented in Fig. 4, shows three doublets. These doublets could correspond to the three crystallographic sites known' to exist in the structure of rock bridgei t e { [Fe;'] [Fey] [Fe"] (PO,),( 0H) 1. The unit framework of the structure of rockbridgeite consists of a cluster of three face sharing octahedra Fe(2)-Fe( 1)-Fe(2). These clusters and other FeO, octahedra [Fe(3)] are con- nected to each other by corner sharing (Fig. 5). The crystallo- graphic sites [Fe(l), Fe(2), Fe(3)] have a relative ratio equal to 20 :40 :40.From bibliographic data and by considering the relative areas of the doublets, the doublet with the quad- 0.0 --0.0 --1 .o -4.2 7 -2.0 -.-c -4.4-3.0-E\ --. --v) -0.0--0.00 3 z-1 .o --4.05 3 tu -c -2.0 -4.10 300 500 700 900 TIK Fig. 2 DTG and TG curves of the (a)giniite and (b) rockbridgeite sample -2 -1 0 1 2 velocity/mm s-' Fig. 3 Experimental Mossbauer spectrum of the giniite sample, recorded at 295 K. Solid lines are derived from least-squares fits. rupolar splitting equal to 0.66 mm s-could be attributed to Fe3+ in the cluster [Fe(2) sites] and that with the smaller quadrupolar splitting (0.38 mm s-') to Fe3+ in the Fe(3) sites. The last doublet with the larger quadrupolar splitting (0.99 mm s-') should correspond to the Fe(1) sites.The occupation of these last sites by Fe3+ instead of Fe2+ cations 3 2 v)4-C 2 0 82 veloci ty/mm s-' Fig. 4 Experimental Mossbauer spectrum of the rockbridgeite sample, recorded at 295 K. Solid lines are derived from least-squares fits. Table 4 Mossbauer parameters computed from the spectrum of the synthetic giniite and natural rockbridgeite sample, recorded at 295 K Mossbauer parameters/mm s-' relative compound site 6 W A intensities (YO) giniite Fe3'(l) 0.42 0.26 1.05 42 Fe3+(2) Fe2'(2) 0.44 1.24 0.26 0.24 0.41 2.37 24 11 rockbridgeite Fe2'(3) Fe3'(l) 1.15 0.42 0.25 0.34 2.61 0.99 23 45 Fe3'(2) 0.43 0.29 0.66 33 Fe3'(3) 0.44 0.24 0.38 22 a4 Fig.5 Fe-0 polyhedra connection in the rockbridgeite structure' should give rise to some local distortions responsible for the large quadrupolar splitting observed. This seemed acceptable, but note that a discrepancy between the relative intensities of the observed doublets (22 : 45 : 33) and the theoretical ratio exists. Discussion Lipscombite, barbosalite and Fe,(PO,),(OH), compounds which adopt relatively similar structures, were shown to undergo dehydroxylation in the same temperature range. Two main weight-loss stages could be distinguished. The first corresponded to a sharp DTG peak and the second to a broad and complex DTG peak due to at least two distinct losses of water molecules. If the temperature range of dehy- droxylation of the three basic phosphates cited above (Table 3) is compared with the temperature range of departure of constitutional water of acid phosphates (Table 5), it can be seen that these temperature ranges are approximately the same.Nathan et aL6 showed that natural lipscombite lost its constitutional water between 603 and 843 K and Gleith' reported the stabilization of the lipscombite structure up to 823 K. These results which seem incompatible may be explained by our results. We confirmed the results obtained by Nathan et d6on natural lipscombite as we observed that the weight loss occurs before 843 K. We observed two very distinct weight losses for the synthetic lipscombite, the first one at ca.715 K and the second between 802 and 843 K. An X-ray diffraction study performed on a sample heated to 773 K (i.e.a temperature intermediate between the two water departures) and quenched at this temperature, allowed us to show that the first water departure, which corresponds to the Table 5 Temperatures of water departure for several acidic and condensed phopshates temperature of compound water departure/K FePO, . 2H20 423-413 Fe(H2P04)3 423-473 FeH2P207 823-923 FeH2P,010. 2H20 313-423 613-823" 613-113FeH2P3010 FeH2P3010.5H20 313-413 613-113" Fe3(P0,0H), .4H20 473-503 673-723" Fe3(P04)2(0H)o.8 7 ' 7.13H20 313-413 ' First temperature ranges correspond to the dehydration process. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 departure of half of the water, did not affect greatly the struc- ture of the lipscombite.8 This stabilization of the structure independently reported by Gleith7 may presumably be explained by the formation of OH vacancies up to 803 K: Fe,(PO,),(OH), -+ Fe3(P0,),0,(OH)2 -,,+ xH,O (where x 6 0.5).There are differences between our results and those of Nathan et al.6 for natural samples. First, the natural samples appear to lose water at a lower temperature since the dehy- dration begins at 603 K, and four poorly defined peaks at 628, 683, 783 and 833 K are also observed. Secondly, we did not observe the oxidation of the iron(I1) ion in lipscombite in an inert atmosphere. This oxidation observed at ca. 473 K, as in air, is proposed to take place without the addition of oxygen, but by removing the H ions from the OH groups, thus forming H, .The Mossbauer data obtained for the giniite sample showed that the solid is in a more reduced form than would be expected from the stoi~hiometry.~ However, the chemical analysis of the sample showed the expected P : Fe ratio and DTG analysis showed two water departures which should correspond to the hydration and hydroxylation water. The first water departure corresponded to 2.6 mol of water per mol of giniite, whereas the second corresponds to 0.6 mol per mol of giniite. This last result, together with the result from Mossbauer spectroscopy, allowed us to propose that the syn- thesized sample giniite corresponds to a reduced sample with the formula Fe\~3Fe:'.7(P0,),(OH)l.3-2.6H20.Therefore, it is postulated that the giniite can correspond to a solid solu-tion in which reduction of the iron is possible and is accom- panied by the substitution of hydroxy groups by water molecules: Fe~~,Fe~-,(PO,),(OH),-, -2 + xH,O. This oxido-reduction process is the fifth process known to occur to iron hydroxyphosphates: type 1 corresponds to a variation in composition with formation of iron vacancies; this is the case for lipscombite, as has been postulated by Gleith:7 Fe',"Fe"(PO,),(OH), -+ Fe:", ,,Fey- ,,(PO,),(OH), ; where 0 < x < 1 Type 2 corresponds to a variation in composition due to dehydroxylation and oxidation. This is the case for rock- bridgeite:'' Fe',"Fe"(PO,),(OH), -+ Fey:,Feft-x(P04)2(OH)2 -,ox; where 0 < x < 1 Type 3 is a combination of types 1 and 2; it has been shown to take place for Fe~'(PO,),(OH), :4 FeT(P0,)3(0H)3 + Fef'- 3,Fe:,(P04)3(OH)3 -3,03,; where 0 < x < 1 /?Fe'"Fe'~PO,)O; when x = 1 Type 4 corresponds to an auto-oxidation by decomposition of crystal water.This process has been shown, by Hanzel et d.," to proceed for vivianite upon heating under vacuum: Fe~(P0,),~8H20-+Fe~~,Fe~(P0,),(0H),~(8-x)H,O + H, Type 5 corresponds to a variation in composition due to dehydration and oxidation. It has also been shown to occur for vivianite upon heating in the presence of oxygen:12*13 Fe','(PO,), . 8H,O + 1/20, -+ Fe~-,Fe~'(PO,),(OH), * (8 -x)H,O + H,O The existence of the giniite solid solution is related to this last process.A solid solution has previously been reported J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 which contained only pure iron(@ giniite.14 The latter was believed to vary in content of iron cations and hydroxy groups as well as hydration water molecules from Fe?~2(P04)4(OH)i.s6 '2.75H2O to Fe?!00(P04)4(OH),.,,O,. 4.60H20. It is, however, improbable that a solid solution with such a variation in composition could really exist. Note that the temperature corresponding to the departure of the hydration water molecules is exceptionally high for giniite compared with other phosphates. As shown in Table 5, the temperature of dehydration of either condensed, acidic or basic phosphates is generally between 373 and 473 K.In contrast, the departure of the hydroxy groups begins at rather low temperatures compared with other hydroxy-phosphates. This phenomenon may be linked to the fact that after the departure of the hydration water molecules the structure of giniite is weakened by the empty channels and collapses immediately, resulting in its dehydroxylation. Our results indicate that the structure of the giniite presents three crystallographic sites in a 40 :40 : 20 ratio which resembles rockbridgeite. The two most numerous sites [Fe( 1)and Fe(2)] would be occupied by Fe"' and the less common one Fe(3), by Fe". With this hypothesis, the results obtained by Moss- bauer spectroscopy could be interpreted as a partial oxida- tion of one of the two more numerous sites [Fe(2)].The doub- let characterized by 6 = 0.42 mm s-' and A = 1.05 mm s-' would correspond to Fe" occupying site Fe(l), that with 6 = 1.15 mm s-' and A = 2.61 mm s-' to Fe" occupying site Fe(3) and those with 6 = 0.44 mm s-' and A = 0.41 mm s-', and 6 = 1.24 mm s-l and A = 2.37 mm s-' would corre- spond, respectively, to Fe"' and Fe" occupying site Fe(2). The formula of giniite obtained from DTG and chemical analyses could then be rewritten as: [Fe~'][Fe',l,Fe&] [Fe"](PO,),(OH),,, * 2.6H20. Such a formula leads to a ratio between the four doublets equal to 40: 26: 14: 20, which is not far from what is observed 42 :24 : 11 :23 (Table 4). The Mossbauer analysis of the rockbridgeite sample showed that it contained only iron(rI1) cations.This is pos- sible since it has been shown that the rockbridgeite corre- sponds to a solid solution in which the oxidation level of the solid could change with the variation in the hydroxy group content: Fe~~,Fe~-,(P04),(OH),-x0,with 0 < x < 1." However, the observed water loss, equal to 2.4 mol per mol of rockbridgeite, is much higher than that of 2.0, weighted for a totally oxidized sample. At this point the presence of a large amount of Mn2+, shown by chemical analysis (Table l), has to be considered. Mn2+ is known to substitute for Fe2+ in the natural hydroxyphosphates. The stoichiometry of our compound has been determined from chemical analysis results (Table 1) as Fe~!2Mnb'.,(PO,),(OH)4.800.2. In this 3339 case the sample would then have a water content correspond- ing to 2.4 mol of water per mol of solid which is in better agreement with our TGA experimental results.Normally the Mn2+ ions should substitute the Fe2+ ions in the structure and thus occupy the less numerous sites. This was not shown in the Mossbauer spectroscopy results. The three sites observed for the iron(Ir1) cations give a relative ratio which did not agree with the substitution of Fe2+ by Mn2+ which would lead to a 48 :48 :4 ratio. The discrepancy between the relative intensities of the observed doublets (45 : 33 : 22) and the relative ratio of the crystallographic sites (40:40 : 20) can only be explained if the Mn2+ ions occupy the Fe(3) sites of the structure normally occupied by Fe3+ ions.In such a case, the ratio calculated from Mossbauer spectroscopy and chemical analysis would be 39 : 42 : 19, which is approx- imately the theoretical ratio. The Fe2+ site Fe(1) would then be occupied by Fe3+ cations and the Fe3+ site Fe(3) by Mn2+ cations: [Fe~1][Fe~f2Mn~,8][Fe111](P04)3(oH)4~800~2. This interpretation does not fully agree with the results of the crystal structure refinement which points to the occupation of the Fe(3) site only by trivalent cations.' However, note that a possible substitution of the iron cations in this site by Mn2+, +CaZ or Fe2 + cations has already been proposed. ' Finally one may suggest that the first water depature, which occurs at a lower temperature and corresponds to 0.20 mol of water per mol of rockbridgeite, could correspond to the formation of water from hydroxy groups and neighbouring oxygen ions already present in the structure.References 1 D. Rouzies, J. M. M. Millet, D. Siew Hew Sam and J. C. VCd-rine, Appl. Catal., submitted. 2 D. Rouzies and J. M. M. Millet, Hyperfine Interact., 1993, 77, 11. 3 D. Rouzies, Ph.D. Thesis, Lyon, 1992. 4 M. Ijjaali, M. Malaman and C. Glietzer, Eur. J. Solid State lnorg. Chem., 1989,26,73. 5 P. B. Moore, Am. Mineral., 1970, 55, 135. 6 Y. Nathan, G. Panczer and S. Gross, Thermochim. Acta, 1988, 135, 259. 7 M. Gleith, Am. Mineral., 1953, 38, 612. 8 J. M. M. Millet, to be published. 9 P. Keller, N. Jb Miner. Abh., 1980, H12,561. 10 M. L. Lindberg, Am. Mineral., 1949,34, 541. 11 D. Hanzel, W. Meisel, D. Hanzel and P. Gutlich, Solid State Commun., 1990,76, 307. 12 R. F. Vochten, E. de Grave and G. Snoops, N. Jhar. Miner. Abh., 1979, 137, 208. 13 J. L. Dormann and J. F. Poullen, Bull. Miner., 1980, 103, 633. 14 J. L. Jambor and J. E. Dutriziac, N. Jb. Miner. Abh., 1988, 159, 51. 15 G. Amthauer and G. R. Rossman, Phys. Chem. Miner., 1984, 11, 37. Paper 4/03 19 1A ;Received 3 1 st May, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003335

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3341-3346 3341 Applications of EPR to a Study of the Hydrogenation of Ethene and Benzene over a Supported Pd Catalyst: Detection of Free Radicals on a Catalyst Surface? Albert F. Carley, Hywel A. Edwards, Brynmor Mile, M. Wyn Roberts and Christopher C. Rowlands School of Chemistry and Applied Chemistry, University of Wales College of Cardiff, P.O.Box 912, Cardiff, UK CFI 3TB Fred E. Hancock and S. David Jackson lCl Katalco, R & T Group, Billingham, Cleveland, UK TS23 ILB EPR has not been used extensively in the field of catalysis despite it being the most sensitive technique avail- able for detecting free radicals and paramagnetic metal ions, which are intermediates or participants in many catalytic processes.In this paper, we describe its application to the detection of paramagnetic intermediates on the surface of a palladium catalyst during heterolytic hydrogenation reactions. EPR, in conjunction with spin trapping, is shown to provide a convenient, simple method for detecting hydrogen adatoms generated by the dissociative chemisorption of hydrogen on alumina-supported palladium catalysts (t0.04% Pd) at room tem- perature. By using D, we have also been able to demonstrate directly the occurrence of hydrogen/deuterium spillover onto the alumina surface by H/D atom transfer to surface hydroxy groups. Alkyl and aromatic free radical intermediates formed by reaction of the H’ adatoms with alkenes and benzene have also been observed by EPR at a catalyst surface for the first time.The hydrogen atoms reacting with the spin traps, ethene and benzene, are not strongly chemisorbed hydrogen (Pd-H) but those weakly adsorbed H-adatoms in equilibrium with H,(g). Despite it being a valuable technique for detecting, identify- ing and monitoring free radicals and paramagnetic ions, elec- tron paramagnetic resonance (EPR) spectroscopy has been rather neglected in the field of catalysis. Nevertheless, there have been important contributions by Lunsford,’ Kevan and Naryana,2 Kasai and Bi~hop,~ Rhodes and Webster4 and others on zeolite catalysts, the incisive results of Che and Tench on metal oxides,> the studies of complex aluminas and silicas6 and the clear demonstration of free radical par- ticipation on metal oxide surfaces by Lunsford and co-workers using matrix-isolation methods.’ CH2FH2H There have been relatively few studies by EPR of real cata- lyst systems since they can be complex, often giving poorly &I CH2-CH2 resolved anisotropic spectral transitions which are difficult to assign.Our understanding of catalytic events at the mecha- A/ \&Inistic molecular level is still at a ‘primitive’ state compared with that of gas and liquid phase reactions. This is in spite of the application of powerful new surface science techniques. CH2-CH2-CH2-CH2Supported-palladium catalysts with very low Pd loadings CH3-CH3(g) (<OM% w/w Pd) are used industrially for the removal of hethyne from ethene streams by the selective hydrogenation of the ethyne to less than 5 ppm before passage of the ethene to Scheme 1 polymerisation plants.Gases such as CO are often added to used avoided level crossing muon spin resonance (ALC-pSR) increase selectivity. Besides the hydrogenation reactions, oli- to detect and measure the diffusion characteristics of ethylgomerisation also occurs to yield a complex mixture of and cyclohexadienyl radicals on silica and alumina sur-hydrocarbons ( 2C4), commonly termed ‘green oil ’, which face~,~2-14 the situation in general is clearly unsatisfactory deactivates the catalyst. The mechanism of reactions shown in Scheme 1 is based on that in the recent review by Bos and Westerterp.* Surface hydrogen atoms, ethylidene, vinyl, ethyl and higher alkenyl and alkyl radicals are implicated as inter- mediates.Elegant labelling and spectroscopic techniquesg support this postulated mechanism from which rate expres- 1 sions have been derived and kinetic parameters extracted. lo spin trapHowever, no surface participants other than the ethylidene Uintermediate, which is a spectator species,” have been char- acterised spectroscopically. Although there is the recent 1 impressive work of Roduner and co-workers, who have rlspin addud ~~~~ ~ t This paper was presented at the 27th International ESR con- ference at the University of Wales, Cardiff, 21st-25th March, 1994. Scheme 2 H2or H2+ C2H4 --T----sinter degassedspin trapsolution in flask F \ catalyst Fig. 1 Schematic of apparatus for exposure of the supported Pd catalyst to gas mixtures, followed by treatment with spin trap solu- tions and transfer to EPR tubes and hence this appeared to be an area where EPR could be usefully deployed.The present work illustrates how free atoms and free radicals present on a palladium catalyst (Scheme 2) can be detected by the method of spin trapping.” Our experiments are based on the pioneering, but neglected, work of Howe and co-workers16 in their studies of hydrogen atoms on the surface of zinc oxide using spin traps. Experimental The apparatus used is shown in Fig. 1. The crushed ‘egg shell’ commercial catalyst (ca. 0.04% w/w Pd on either an alumina or calcium aluminate support with total surface areas of 55 m2 g- ’ and 10 m2 g- ’,respectively, and Pd par- ticles of 1.5-3.0 nm dispersed near the primary support parti- cle surface) were pretreated in situ at 400°C for 2 h under vacuum.By opening tap A they were exposed to atmospheres of pure hydrogen or hydrogen and ethene mixtures at 25°C for up to 10 min. The degassed spin trap solutions [O.Ol mol dm-N-tert-butyl-a-phenylnitrone (PBN) or 2,4,6-tri-tert-butylnitrosobenzene (BNB) in degassed CH,Cl,], which were isolated in flask F from the catalyst during these oper- ations, were than poured onto, and stirred with, the catalysts for prescribed periods of time before being tipped into the quartz EPR sample tube through the coarse glass sinter. After closing tap B, the sample was inserted into the EPR cavity of a JEOL JES-RE2X spectrometer operating at X band with 100 kHz modulation and detection at this fre- quency.The spectra were collected and averaged over 10 multiple scans. In some experiments, the catalyst was evac- uated to mbar for ca. 30 min to remove all but strongly chemisorbed hydrogen adatoms before exposure to the spin trap solution under helium gas. The catalyst was prepared by established methods at ICI Katalco.” H2(> 99.99%), D2(> 99.8%) and ethene (~99.98%) were supplied by BOC Special Gases; benzene, CH2Cl,( >99Y0) and the spin traps were supplied by Aldrich. Results and Discussion Hydrogen Dissociation and Spillover Fig. 2(a) shows the EPR spectrum of a PBN solution after admixing with the hydrogenated catalyst for 10 min.The iso- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1 rnT I-+ I Fig. 2 (a) EPR spectrum at 25°C of a PBN solution in CH2C12 (0.01 mol dm-3 PBN) after mixing with a Pd catalyst treated with H, (1 x lo5 Pa) for 10 min at 25 “C; (b) simulation of spectrum using parameters given in the text tropic spectrum is clearly a triplet of triplets and is accurately simulated [Fig. 2(b)] by the hyperfine parameters a&) = 1.53 mT with one nitrogen nucleus and 42) = 0.82 mT, with two equivalent hydrogen nuclei. The spectrum can be assigned unequivocally to the hydrogen atom spin adduct, A (Scheme 3)’ by comparison with well established spectra and hyperfine parameters for the monohydrogen atom adduct of PBN.16*18 The present spectra are superior to those previously reported in that M,(N) = -1, M,(2H) = +1 and M,(N) = 0, MA2H) = -1 together with the corresponding high-field transitions are clearly resolved here compared with the unre- ’ilsolved doublets previously observed.’ No EPR tran-sitions were detected when the undoped support was similarly treated.0 *. 0 spin addud A Scheme 3 There can be little doubt that the PBN has been able to scavenge surface hydrogen atoms to give the adduct A (Scheme 3). The rate constants for radical reactions in solu- tion with PBN are of the order of 1 x 10’ dm3 mol-” s-’ and clearly large enough in the case of adsorbed hydrogen atoms to ensure that they can be removed efficiently from the metal surface.We sought to confirm our conclusions by substituting pure deuterium for hydrogen in an identical experiment, expecting the triplet of triplets from the hydrogen atom adduct to be replaced by a triplet of doublets of triplets with a,,(l) = 0.12 mT from the monodeuteriated, monohydrogenated spin adduct. Fig. 3(a)shows that these were indeed the major tran- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1 mT Pd-Pd -0-Al-O-Al-O- Al-f O-AI-O-Al-O-Pd-Pd+-d dl 0 Al 0 Al 0 A1 0 Pd-Pd -0-AI-0-Al-0-AllO-AI-0-Al-0-Pd-Pd11 b 0 11< Al 0 Al 0 Al 0 Scheme 4 1 mT+ 1 mT I--+ Fig. 3 (a) EPR spectrum of a PBN solution after mixing with the Pd catalyst exposed to D, for 10 min at 25°C; H denotes lines from the dihydro PBN spin adduct, all other transitions arise from the monohydro-monodeuterio spin adduct; (b) simulation of a 15 :85 composite spectrum from the dihydro and monohydro-mono PBN spin adduct mixture.(c) spectrum of a PBN solution after mixing with a Pd catalyst treated with D,O before conditioning at 400°C and then exposure to pure D, for 10 min at 25 "C. sitions observed (ca. 85%) but, surprisingly, despite the high isotopic purity of the D, used (>99.8%), there were still sig- nificant transitions from the dihydrogen adduct [labelled in Fig. 3(a)by HI, indicating directly that hydrogen atoms were generated by adsorbing D, on a supported palladium surface. The simulation in Fig. 3(b)is that for a 15 :85 HD composite adduct spectrum. The most likely source of hydrogen is the surface hydroxy groups of the alumina support (usually about five OH groups per 100 A,).This was proven by repla- cing the surface hydroxy groups with deuterioxy groups during multiple treatment of the catalyst with D20 at 80°C (1 :100 catalyst to D,O ratio). This D,O-exchanged catalyst was pretreated at 400°C and then exposed to D, gas and PBN solution as already described. The EPR spectrum, Fig. 3(c), shows that, as expected, the dihydrogen signals have been considerably reduced. We believe these results to be direct evidence for hydrogen spillover, where hydrogen atoms formed on metal particles spill over onto the surface of the support, which does not, by itself dissociatively chemisorb molecular hydrogen. The We now consider the nature of the hydrogen/deuterium atoms that react with the spin trap to give the hydrogen/ deuterium spin adduct.Two possibilities exist: (a) the hydro- gen atoms are those strongly chemisorbed to the palladium; (b) they are mobile hydrogen atoms only weakly adsorbed at the palladium surface. The following experiment was per- formed to distingush between these two alternatives. The H, treated catalyst was evacuated for 30 min at 21 "C and mbar and then exposed to the PBN spin trap solu- tion under helium. Thus, only strongly chemisorbed hydro- gen adatoms were left on the surface to react with the PBN. This procedure resulted in the complete loss of the dihydro- gen spin adduct showing conclusively that (b) is correct, i.e.the active hydrogen is not that strongly chemisorbed at the palladium surface. It is interesting to speculate that these hydrogen adatoms may be those that are active in alkene hydrogenation reactions. Experiments are in progress to check whether the rate of hydrogenation can be correlated with the concentrations of these active, mobile hydrogen ada t oms. The ratio for the relative H and D atom spin adduct con- centrations (15/85 = 0.18) is a measure of the equilibrium constant for the following equilibrium at the palladium and alumina surfaces, assuming the H moieties are weakly bound. ---Pd---D+ H-O-Al---=---Pd---H + D-0-Al---We make the reasonable assumption that the surface H and D atoms are transferred to PBN with equal facility. The dif- ference in zero-point energies calculated by using vpdWH= ~1850 cm-' and v ~ =-3650 cm-' and, assuming2, that vD = vJJ2 is 5.2 kJ mol-' which is higher than the value of 4.3 kJ mol-' estimated from the ratio of spin adduct popu- lations.The difference is close to the experimental error but is in the direction expected for a more loosely bound H adatom reacting with PBN. Formation and Trapping of Alkyl and Cycloalkadienyl Radicals at a Pd Surface Ethene As already discussed and illustrated in Scheme 1, ethyl rad- icals have been suggested as important intermediates in ethene hydrogenation. The involvement of ethyl species in the hydrogenation of ethene was first established by Kemba1123 using deuterium, where evidence for the monodeuteriated species C,H,D(g) was obtained mass-spectrometrically.These ethyl species were assumed to be chemisorbed. More recently, Thomson and Webb24 concluded that only a weakly held C2H4 species was hydrogenated, the more strongly che- occurrence of such spillover is still a topic of controver~y'~*~~ and is difficult to trace experimentally, as illustrated by NMR studies of the Euro Pt/SiO, catalyst. These were first inter- preted to show the occurrence of spillover, and then reinter- preted to demonstrate its absence.21 We envisage the hydrogen spillover of hydrogen atoms to occur by a 'Grotthus type' chain or 'knock-ony effect as illustrated in Scheme 4. misorbed ethene being inactive. Since this paper was sub- mitted, Bowker et a1.,25on the basis of a mass-spectrometric temperature-programmed desorption approach have further discussed the role of weakly adsorbed ethene in catalytic hydrogenation at Rh(ll1) surfaces.Here we report a demons- tration of the participation of ethyl radicals as intermediates using the spin-trapping technique. Fig. 4 EPR spectrum of a PBN solution after mixture with a Pd catalyst which had been exposed at 25°C to a 50: 50 mixture of H,-C2H, for 10 min The pretreated catalyst was exposed at 298 K to a 50 :50 mixture of H, and C2H4 for 10 min before the PBN spin trap was added and the resulting solutions were then manipulated in the way already described. Fig. 4 shows the spectrum of the resulting solution which exhibits no hydrogen atom PBN adduct transitions, indicating that the ethene has reduced the surface hydrogen atom concentration to undetectable levels.The triplet of triplets spectrum of Fig. 2(a) is now replaced by a triplet of doublets with uN = 1.47 mT and u,(l) = 0.23 mT which arise from a radical spin adduct, B (Scheme 5), rather than a hydrogen atom adduct. The nitrogen and hydrogen hyperfine parameters are consistent with X' being an alkyl radical, although an alkoxy radical or alkyl peroxyl spin adduct cannot be excluded because of the similarities of the hyperfine parameters for these three adducts. spin addud B Scheme 5 In order to distinguish between these two possibilities, we used the nitroso spin trap BNB which is much more defini- tive in this regard since a-proton hyperfine parameters from the parent radical can usually be observed.On repeating the experiment with 50 :50 H2-C,H, mixtures, but using BNB solutions, the spectrum, shown in Fig. 5, was observed. The main transitions show unambiguously that X' has a CH; ter- minal group. They can be accurately simulated by the follow- ing parameters; g = 2.007 & 0.0005,a, = 1.31 mT, a,(2H) = 1.81 mT, ad2H) = 0.08 mT which are close to those for the ethyl BNB adduct.26 We have, therefore, conclusive proof of alkyl (probably ethyl) radical formation on the palladium surface, and have observed the occurrence of surface hydro- gen atom addition to ethene to give ethyl or higher C,H;,+ 3 2) radicals.The additional triplet of quartet fea- tures, nsrked R in Fig. 5, arise from a radical addition to the oxygeri md (1: the BNB as indicated by an upward shift of *r.L;-cemz of riie spectrum and, hence, a downward shift in ~-k:a-to g = 2,0037. Such oxygen addition only occurs for ._ ~~ic-~k~~.4. -28 indicating that oligomeric and/or poly- i-w-3: -xii-:dz are being formed and trapped in addition to J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 R 9 Fig. 5 EPR spectrum at 25°C of a BNB solution after mixing with a Pd catalyst exposed to a mixture of 50 :50 H,-C2H, at 1 bar for 10 min at 25 "C. Lines A are from the ethyl spin adduct with addition at the nitrogen end of the nitroso groups and lines R are from a secondary alkyl radical RCHR attaching at the oxygen of the nitroso group. small primary radicals such as ethyl.The quartet of lines indicate three equivalent hydrogens; two are probably the rn-protons of the benzene ring which usually have hyperfine parameters of 0.186 mT(a, = 1.0 mT), while the third is a single proton on the attacking radical, i.e. a secondary radical, R-CH-R. This constitutes good evidence for the occurrence of either an intermolecular hydrogen transfer between an oligomeric radical and an adjacent preformed macromolecule or an intramolecular hydrogen transfer or back biting by a large oligomeric or polymeric primary alkyl radical, Scheme 6. Attack at the oxygen of BNB, indicates a Scheme 6 large secondary or tertiary radical and is consistent with the presence of a large and bulky oligomeric or polymeric radical on the palladium surface.Fig. 6 shows the spectrum when a 80 : 20 H,-C,H4 mixture was used. The decrease in the R components is evidence for this assignment to an oligomeric or polymeric component which should decrease with the Fig. 6 EPR spectrum of a BNB solution after mixing with a Pd catalyst exposed to a 80 :20 H,-C2H, mixture J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 decrease in ethene levels. Further experiments are now in progress to study the reactions of C2D4, CH2CD, and dial- kenes on a range of catalysts. Benzene Because the addition of hydrogen atoms to benzene is less efficient than to alkene~,~’,~~ we increased the benzene con- centration substantially by using it as the solvent for BNB.The procedure was as described but a 0.01 mol dm-3 BNB solution in benzene was mixed with the palladium catalyst after it had been exposed to pure H, for 10 min. The resulting spin adduct spectrum, shown in Fig. 7(a), again shows the complete loss of the hydrogen atom adduct tran- sitions and their replacement by a well resolved major sextet of pentets arising from the following interactions: a&) = 1.22 mT, 43) = 1.22 mT and 444) = 0.07 mT. Fig. 7(b) shows a simulation of the spectrum based on these param- eters. We assign the experimental spectrum to the cyclo- hexadienyl spin adduct (C, Scheme 7) resulting from the spin adduct C Scheme 7 two-step sequence of a surface hydrogen atom reaction with benzene to give the 2,3-cyclohexadien- 1-yl radical which is then scavenged by the BNB to give adduct C.The large sextet proton hyperfine parameters arise from the three protons at the 1 and 2 positions of the cyclohexadienyl moiety and the small proton quintet hyperfine parameters from the two protons at the 4 and 5 positions of the cyclo- hexadienyl and the two at the 3 and 5 positions of the 1,3,5- tri-tert-butylbenzene. This is the first report of the EPR spec-trum of the cyclohexadienyl radical adduct of BNB but the 1 mT w Fig. 7 (a) EPR spectrum of a BNB solution in benzene after addi- tion to a Pd catalyst exposed to H,at 1 mbar for 10 min at room temperature; (b)simulation of spectrum using the parameters given in the text parameters are those expected by comparison with the allyl spin adduct reported by Terabe and Konaka,26 uH (from one CH,) = 1.81 mT; aN(1) = 1.34 mT and a,(2 m-H) = 0.08 mT.No coupling is observed from the central proton of the allyl entity or those of the other CH2 group which were equivalent in the parent allyl radical, i.e. we do not have allyl x-type bonding across the two terminal carbons as observed in many transition-metal allyl complexes. These assignments will be confirmed by using suitably deuteriated benzenes and a combination of H2 and D, experiments. Again, the spin trapping technique has revealed, for the first time, the presence of precursors en route to the complete hydrogenation of benzene to cyclohexane. We are presently conducting a comprehensive study of a range of substrates and catalysts.Note that the samples were ‘as used’ in industrial plants and examined without prior treatment other than the normal activation procedures. Conclusions In the present study it is clear that at palladium surfaces only weakly chemisorbed hydrogen interacts with the spin trap and is therefore observed by EPR. This raises the question as to whether the spin trap intercepts the hydrogen adatom, H(s), just subsequent to the cleavage of the dihydrogen bond but prior to it being chemisorbed or whether the spin trap interacts with the weakly chemisorbed hydrogen adatom, H(a): 9H2(g)*H(s) H(a)-+ With increasing coverage the heat of H(a), a chemisorption will decrease and only under these conditions would we anticipate the reversible step to be significant (see above).It is therefore either the H(s) species or the weakly chemisorbed hydrogen, H(a), that we observe by EPR and not the strongly chemisorbed hydrogen, H(a). An analogous argument may also apply to the hydro- genation of ethene: C2H4(g) C2H4(a); strong chemisorbed and inactive in hydrogenation C2H4(g)-’ C2H%s); weakly adsorbed at high coverage and active C,HX(s) + H(s) --* C,HT(s); two dimensional gas reaction : first step of hydrogenation C2HZ(s)+ H(s)+C,H,(g); desorption of ethane This mechanism is essentially that of Kemball,23 and Thomson and Webb.24 The C2H, spin trap species has been observed unequivocally in the present EPR study.If we accept that the spin trap is specific to radical-type species only, then the present observations support a mechanism involving the participation of ethyl radicals present in a two- dimensional gas-like overlayer composed of weakly adsorbed ethene C,HX(s) and weakly adsorbed H adatoms H(s), in equilibrium with H2(g). Both H(s) and C,H;(s) are observed by EPR. Analogous free radical chemistry is observed in the catalytic hydrogenation of benzene. We are grateful to ICI Katalco for their generous support of this work. H.A.E. thanks the SERC for the award of a CASE studentship and ICI Katalco for their sponsorship. We also thank Professor M. S. Spencer for illuminating discussions and comments.3 346 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 References 17 ICI, UK Pat., 7940086, 1979. 1 J. H. Lunsford, Adv. Catal., 1972,22,265. 2 L. Kevan and M. Naryana, Intrazeolite Chemistry, ed. G. D. Stucky and F. G. Dwyer, ACS Symp. Ser. 218, 1984, p. 283. 3 P. H. Kasai and R. J. Bishop Jr., Zeolite Chemistry and Cataly- sis, ACS Monograph 171, ed. J. A. Rabo, 1977, p. 350. 4 C. J. Rhodes and B. C. Webster, J. Chem. SOC., Faruday Trans., 1993,89,1283. 5 M. Che and A. J. Tench, Adv. Catal., 1983,32, 1. 6 M. Che and E. Giamello, Catalysis Characterisation: Physical Techniques for Solid Materials, ed. B. Imelik and J. C. Vedrine, Plenum Press, New York, 1994, p. 131. 7 D. J. Driscoll, W. Martin, J. X. Wang and J. H. Lunsford, Adsorption and Catalysis on Oxide Surjiaces, ed.M. Che and G. C. Bond, Elsevier, Amsterdam, 1965, p. 403. 8 A. N. R. Bos and K. R. Westerterp, Chem. Eng. Process., 1993, 32, 1. 9 S. Le Viness, V. Nasr, A. H. Weiss, Z. Schay and L. Guczi, J. Mol. Catal., 1984, 25, 131. 10 L. Z. Gra and K. G. Kho, Kinet. Catal., 1988,29,381. 11 F. Zaera and G. A. Somorjai, J. Am. Chem. SOC.,1984,106,2288. 12 I. V. Reid, T. Azuma and E. Roduner, Nature (London), 1990, 345, 328. 13 E. Roduner, Chem. SOC. Rev., 1993,22,337. 14 E. Roduner, Phys. Rev. B, in the press. 15 E. G. Janzen in Free Radicals in Biology, ed. W. A. Pryor, Aca- demic Press. New York. 1980. vol. 4. D. 115. 16 T. Matsuzaki, T. Uda,’A. Kkusak;, G. W. Keulks and R. F. Howe, J. Am. Chem. SOC., 1980,102,7511.18 0. Augusto, P. R. Ortiz de Montellano and A. Quintanilha, Biochem. Biophys. Res. Commun., 1981,101,1324. 19 M. S. Spencer, Appl. Catal., 1989,55, N23. 20 New Aspects of Spillover Eflect in Catalysis, Stud. Surf. Sci. Catal., ed.T. Inui, Elsevier, Amsterdam, 1993, vol. 77. 21 M. A. Chesters, A. Dolan, D. Lennon, D. J. Williamson and K. J. Packer, J. Chem. SOC., Faruday Trans., 1990, 86, 3491; M. A. Chesters, D. Lennon, H. Viner and K. J. Packer, J. Chem. SOC.,Faruday Trans., submitted. 22 L. Lelander and W. H. Saunders Jr., Reaction Rates of Isotopic Molecules, Krieger Publishing Co., Malabar, 1991, 109. 23 C. Kemball, Proc. Chem. SOC., 1954,30,1164. 24 S. J. Thomson and G. Webb, J. Chem. SOC., Chem. Commun., 1976,526. 25 M. Bowker, J. L. Gland, R. W. Joyner, Y. Li, M. M. Slin’ko and R.Whyman, Catal. Lett., 1994,25,293. 26 S. Terabe and R. Konaka, J. Chem. SOC., Perkin Trans. 2, 1972, 369. 27 B. Mile, C. C. Rowlands, P. D. Sillman and M. Fildes, J. Chem. SOC.,Perkin Trans. 2, 1992, 1431. 28 B. Mile, C. C. Rowlands, P. D. Sillman and A. Holmes, J. Chem. SOC.,Perkin Trans. 2, 1993,2141. 29 J. E. Bennett and B. Mile, J. Chem. SOC., Faruday Trans. 1, 1973, 69, 1398. 30 C. H. Bamford and C. F. H. Tipper, Comprehensive Chemical Kinetics, 1969,3,97. Paper 4102235A; Received 14th April, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003341

出版商:RSC    

年代:1994

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J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3347-3356 FTIR Studies on the Selective Oxidation and Combustion of Light Hydrocarbons at Metal Oxide Surfaces Propane and Propene Oxidation on MgCr,O, Elisabetta Finocchio, Guido Busca and Vincenzo Lorenzelli lstituto di Chimica, Facolta di Ingegneria , Universita, P. le Kennedy, 1-16129 Genova, Italy Ronald J. Willey Department of Chemical Engineering, Northeastern University, Boston, MA 021 15,USA The interaction of propene and propane with the surface of the oxidized spinel MgCr204+x has been studied in the temperature range 300-773 K by FTIR spectroscopy. This solid is reduced reversibly by reaction with these organic compounds in the temperature range 300-673 K, giving rise to stoichiometric MgCr,04 and more oxi- dized organic species that finally produce CO,.Comparison with the results of adsorption and oxidation of C, oxygenates (propan-1-01, propan-2-01, allyl alcohol, propionaldehyde, acetone, acrolein, propionic acid and acrylic acid) as well as of C, and C, oxygenates showed that the predominant oxidation pathways for the two molecules are different. Oxidation of propene occurs predominantly through its previous activation at C, to give strongly adsorbed acrolein and acrylate species. These species later burn. Acetone is the primary oxidation product of propane at the surface, at 423 K. Acetone is later oxidized to formate species (which rapidly decompose) and to acetate species that burn at higher temperatures (573-773 K). The different observed paths are rationalized by taking into account the lower C-H dissociation energy at the methylene group in the case of propane and at the allylic methyl group in the case of propene.The data reported here are consistent with the data available on catalytic alkane oxidation over this and similar catalysts. A comparison is made with the behaviour observed with more selective oxidation catalysts like Mg,(VO,), , V,O,-TiO, and MOO,-TiO, . A mechanism for propene and propane catalytic combustion is proposed. The production of oxygenates via catalytic partial oxidation of hydrocarbons can be achieved with high selectivities using light alkenes as the reactants. Accordingly, the selective oxi- dation of alkenes constitutes a prominent area of industrial petrochemistry. '-' However, several potentially useful oxida- tions of alkenes, e.g.acetaldehyde and acetone production from ethene and propene and the one-step propylene oxide synthesis, still cannot be achieved effeciently in the gas phase. Moreover, efforts are focused today on the use of light alkanes arising from natural gas as raw materials in pet- rochemistry. Nevertheless, the partial oxidation of alkanes is far more difficult than that of alkenes and is at present limited, at the industrial scale, to butane oxidation to maleic anh~dride.~Other processes, like the direct ammoxidation of propane to acrylonitrile' and the oxidative dehydrogenation of propane and butane to the corresponding alkenes6 are now under promising development.On the other hand, the hydrocarbon catalytic combustion is also of great interest at present to limit pollutant emissions of waste In particular, the replacement of noble metal-based catalysts with the less expensive transition-metal oxide-based catalysts has been attempted. In order to develop new, more efficient and more convenient processes based on hydrocarbon catalytic oxidation, detailed mechanisms for these reactions are required. However, the factors affecting the selectivities in hydrocarbon oxidations are still far from clear.2 Moreover, according to Spivey," a general theory of catalytic combustion on metal oxides still does not exist. In our laboratory, FTIR spectroscopy has been recently applied to the study of some oxidation and oxidative dehy- drogenation reactions on selective catalysts.'1-14We recently undertook more systematic studies on selective and unselective oxidation of light hydrocarbons at catalyst sur- faces. As a first step, we report here studies on propene and propane oxidation on MgCr,O, . This material was chosen because it is recognized as a good combustion catalyst for alkanes and alkenes,' 5*1 like chromia and other chr~mites."~'~ MgCr,O, was found to be more active than chromia in methanol oxidation," and is more stable than chromia.' Moreover, it displays significant selectivity in some oxidative dehydrogenation reactions.20 It is a typical p-type semiconductor, and this property is associ-ated with its activity both as an oxidation catalyst and as an oxygen21 and ethanol22 sensor.Finally, it can be prepared as an aerogel and it displays excellent optical properties for IR spectroscopic Experimental The preparation of the MgCr,O, aerogel has been described previ~usly.~~A mixture of Mg acetate and Cr acetylacetonate was dissolved in methanol and hydrolysed with the stoichio- metric amount of water. The resulting gel was dried under supercritical conditions in an autoclave (524 K, 127.5 bar), i.e. above the critical point of methanol (512.6 K, 80.9 bar). The powder was then calcined for 6 h at 973 K. The resulting surface area was 53 m2 g-'. FTIR spectra have been record- ed with a Nicolet Magna 750 instrument, using conventional IR cells connected to an evacuation-gas manipulation appar- atus.The catalyst powder was pressed into self-supporting discs, calcined in air at 723 K for 1 h and outgassed at 723 K for 20 min before use in the interaction experiments. Liquid adsorbates were from Carlo Erba (Milano, Italy), while gases were purchased from SIO (Milano, Italy). The chemicals used in the catalyst preparation were from Aldrich. Results Catalyst Characterization The X-ray diffraction (XRD) pattern of the catalyst shows all of the diffraction peaks typical of the spinel MgCr,04 only (JCPDS table 10-351). The FTIR-FT-FIR spectrum (Fig. 1) also shows the typical features of the spinel MgCr,0424 with J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 wavenumber/cm-' Fig.1 FTIR spectra of MgCr,O,: (a) FTIR spectrum (1200-400 cm-') of a KBr pressed disc; (b) FT-FIR spectrum (500-50 cm-') of the powder deposited on a polyethylene disc; (c) FTIR spectrum of a pressed disc of the pure powder after outgassing at 723 K; (d)after (c) and following an interaction with hydrocarbon in the temperature range 573-773 K the four IR-active modes at 630, 480, 415 and 250 cm-'. However, the IR spectrum, which is also sensitive to non- crystalline and surface species, shows an additional weak complex band in the region 1000-800 cm- ', typically associ- ated with Cr-0 stretchings involving high-valency Cr. Although it seems reasonable to identify these species as involving hexavalent Cr6 + species, according to Klissurski et al." it cannot be excluded that Cr5+ is involved. The above spectral features indeed roughly correspond to the most intense bands of Mg chromate(vI), MgCrO, .26 However, the IR spectra of metal chromate(v) species are characterized by Cr-0 stretching bands (900-800 cm-27) at frequencies only slightly lower than, or superimposed upon, those of the chromate(v1) species.The FTIR spectra of pressed discs of the pure catalyst powder show, as usual, a transmission window above 800 cm-', with maximum transmittance at 1300 cm-' and a pro- gressive decrease in transmission towards higher frequencies due to the wavelength dependence of radiation scattering. Also these samples show the above-cited absorption in the region 1050-800 cm-', clearly structured into one sharp band at 1002 cm-', and components at 975,940,920 and 840 cm-'.A broad absorption is also found near 3500 cm-', due to H-bonded surface hydroxy groups (Fig. 2). The IR spectra of the pure catalyst disc after interaction with hydrocarbons or oxygenated organic compounds in the temperature range 573-773 K are different [Fig. l(6) and Fig. 2(b)].The absorption in the region 1050-800 cm-' is either decreased strongly in intensity or disappears completely, leaving weaker harmonic bands of the skeletal vibrations of MgCr,O,. However, if oxygen or air are admitted into the cell in the same temperature range the absorptions in the 1000-800 cm- 'range reappear quickly. This behaviour agrees with the known p-type semicon- ducting behaviour of MgCr,O, .21~22~28,29In fact, MgCr,O, in oxidizing atmospheres is partly oxidized mainly at the surface" giving rise to MgCr,O,+,, and is associated with the appearance of Cr=O stretching bands of high-valency chromium in the region 1050-800 cm-'.On the other hand, interaction with reducing agents, like hydrocarbons, oxygen- ated compounds and hydr~gen,'~ causes the non-stoichiometric oxygen to be destroyed, producing nearly stoichiometric MgCr,O, . Our data show that redox cycles are relatively fast in the temperature range 573-773 K, i.e. in the range where this 1 4000 3800 3600 3400 3200 3000 wavenu mber/cm -' Fig. 2 FTIR spectra (OH-stretching region) of MgCr,O, disc out- gassed at 723 K: (a) oxidized sample; (b) sample previously reduced in hydrocarbon atmosphere (673 K, 150 Torr) compound is active as an oxidation catalyst' 5,16~19-20 and that the analysis of the IR spectrum in the region 1050-800 cm-' allows the oxidation state of the catalyst to be moni- tored.Reduction also causes the growth of a band near 3750-3650 cm-', due to free surface hydroxy groups [Fig. 2(b)], while bands due to adsorbed hydrocarbon residues may also be present in the region 1800-1200 cm-' [Fig. l(d)]. As discussed recently,', the surface acidity may have a rel- evant role in the catalyst oxidation activity and selectivity. Therefore, we tested the surface acidity of oxdized MgCr,O,+, and of nearly stoichiometric MgCr,O, via adsorption of pyridine as probe (Fig.3). The spectrum of pyridine on the 'reduced' nearly stoichiometric catalyst shows sharp bands at 1607, 1576, 1488, 1443, 1219, 1148, 1068 and 1042 cm- I, associated with pyridine coordinated over Lewis acidic Cr3+ sites.30 The spectrum of pyridine adsorbed on the oxidized sample shows the same pattern but with additional bands. Sharp peaks at 1593 and 1033 cm-I persist after outgassing at room temperature (RT) and are associated either with physisorbed pyridine or with species bonded to OH groups. This species is not so evident on the reduced sample, which might suggest the presence of H- bonding with the hydroxy groups that are responsible, in the 1700 1600 1500 1400 1300 1200 1100 1000' wavenumber/cm-' Fig. 3 FTIR spectra of pyridine adsorbed on oxidized MgCr,O, at room temperature (a), and after successive outgassing for 30 min at room temperature (b), 373 K (c) and 423 K (6);(e) as (b) but on a pre-reduced sample J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 activated oxidized sample, for the broad band near 3500 cm-' [Fig. 2(a)]. Additionally, two other weak bands at 1640 and 1538 cm-' are also evident and persist after outgassing at 373 K; these bands are associated typically with pyridinium cations, produced by protonation of pyridine by hydroxy groups of sufficient Brarnsted acid strength. These data show that the oxidized catalyst MgCr204+, has medium Lewis acidity, similar to that of the reduced one, but is also a weak Brarnsted acid, having active hydroxy groups able to interact with pyridine by H-bonding and by protonation. Behaviour similar to the above was found in the case of oxidized and reduced Cr,O, (using ammonia as the basic probe molecule3') and is ascribed to the greater cova- lency of the Cr"+-0 bond (n= 6 or 5) compared with the Cr3+-0 bonds, resulting in a stronger acidity of the OH bonded to the high-valency Cr than that bonded to Cr3+.The subtraction spectra relative to pyridine adsorbed on the oxidized sample also show a negative sharp band at 1003 cm-'. This band has been assigned above to Cr=O stretch- ing of a chromate species. Its disappearance upon pyridine adsorption provides evidence for the ability of the corre-sponding species to interact with pyridine.It seems reason- able that the band at 1003 cm-' is due to the Cr=O stretching of a surface chromyl group having one coordi- native unsaturation, which can consequently coordinate one pyridine molecule, resulting in a weakening of the Cr=O bond and a decrease of the stretching frequency. Surface chromyl species perturbed by adsorbed pyridine were also observed on oxidized ~hromia.~, Interaction with Propene Gas The FTIR spectrum of the adsorbed species arising from contact of the activated MgCr,04 with propene gas at RT is reported in Fig. qa). Most bands are assigned to molecularly adsorbed propene as deduced from a comparison with the spectra of propene in the gas-phase and adsorbed on other oxide surface^.^^.^^ The sharp band at 1638 cm-' is due to the C-C stretching (1653 cm- 'in the gas phase), while those at 1454, 1440, 1420, 1377, 1178 and 1048 cm-' are associated 3349 that the spectrum is detectable and that the C=C stretching is definitely shifted to lower frequencies show that the inter- action strength is non-negligible and probably implies elec- tron withdrawal from the C=C double bond.The vibrational perturbation is weaker than that previously observed on reduced E-C~,O~, but is similar to that occurring on the unreduced FeCrO, dehydrogenation catalyst,' whose surface is also 'covered' by chromate species. We can suppose that chromate species, more than reduced Cr3 + centres, are involved in propene molecular adsorption. Molecularly adsorbed propene is desorbed by outgassing at RT.However, other small absorptions that cannot be assigned to adsorbed propene are observed in the spectrum. These bands include a definite peak at 1675 cm-', almost certainly due to the C=O stretching of an adsorbed carbonyl com- pound, a broad band near 1600 cm-' [possibly due to the vas(COO) of a carboxylate species, whose symmetric counter- part is masked by the CH deformations] and sharp peaks at 1350 and 1256 cm-'. Heating the sample in propene at 373 K causes the forma- tion of stronger bands certainly due to transformation pro- ducts [Fig. 4(c) and (41.The carbonyl band at 1680 cm- ' is now more evident and exhibits a shoulder at 1647 cm-'. The region 1600-1300 cm- ' presents several peaks [1595, 1562, 1452, 1435, 1424, 1382, 1372 (very weak), 1356 and 1330 cm-'1, which are associated with COO stretchings of differ- ent carboxylate species and CH deformations of several adsorbed species.Weak bands are also observed at 1255 and 1241 cm-', while below 1200 cm-' we find peaks at 1180, 1160 (weak), 1115, 1090 and 1030 cm-', (very strong). The features in the 1200-1000 cm-' region are assigned to C-C stretchings and/or C-C/C-0 coupled stretchings of alco- holate species. The bands below 1300 cm -completely disappear upon further heating at 473 K [Fig. 4(f)],when the components near 1680 and 1640 cm-' also decrease in intensity at differ- ent rates, thus demonstrating that they are not due to the same species. Instead, heating causes a further growth of the absorption in the region 1600-1300 cm-', which finally exhibits peaks at 1596, 1560, 1500, 1435, 1385 and 1355 cm-'.with different CH deformations, as dicussed previo~sly.~~,~~ Above 523 K all bands decrease in intensity and CO, is The vibrational perturbation is small and does not give a precise indication of the adsorption mode. However, the fact 0.40---~~ 0.351 0.30 a 0.25 C;0.20 20 m 0.15 0.10 0.05 0.00 2000 19001800170016001500 14001300 12001100 wavenumber/cm-' Fig. 4 FTIR spectra of the adsorbed species arising from propene adsorption over MgCr,O,+, at room temperature (a), after 30 min at room temperature (b), heating in propene at 373 K (c), and after outgassing at 373 K (4,432 K (e),473 K (f)523 K (9) found in the gas phase.Moreover, Cr=O stretching bands are substantially eroded, providing evidence for the reduction of the oxidized catalyst surface according to the oxidation of the adsorbed organic species. This behaviour indicates that propene oxidation at the MgCr,O, surface gives rise above 373 K to alcoholate and carbonyl species that later transform upon heating to carbox- ylate species and finally burn above 523 K. The experiments described below are aimed at the individual identification of these oxygen-containing species. Interaction with Propane Gas The interaction of the activated oxidized catalyst with 50-150 Torr of pure propane gas has also been investigated in the temperature range 300-773 K.Below 373 K, the interaction does not give any appreciable effect. Above 373 K, instead, this interaction gives rise to new absorption bands (Fig. 5) associated with adsorbed species. Simultaneously, the Cr=O stretching bands at 1000-900 cm-', typical of the oxidized catalyst, decrease in intensity, and corresponding negative bands appear in the subtraction spectra. The interaction at 423 K produces two broad and rather strong bands at 1600 and 1440 cm-' and weaker components at 1380 and 1350 cm-'. Together with these absorptions, a pronounced band split at 1702 and 1692 cm- 'and weaker bands at 1245, 1176 3350 0.30 0.25 al 0.20 C 0.15 0 0.10 0.05 0.00 I ,. , . . . , . . . , . . . , .. . , . . . , . . , . . . I . , , . . . , . . . , . .Y 2200 2000 1800 1600 1400 1200 wavenumber/cm-’ Fig. 5 FTIR spectra of the adsorbed species arising from propane adsorption over MgCr,O, and successive evacuation at 423 K (a), 453 K (b),473 K (c), 523 K (6)and 573 K (e) and 1094 cm-’ are found. The features at 1702, 1692, 1245 and 1094 cm-’ disappear progressively upon further heating in vacuum, while the other bands grow and shift slightly to 1595, 1439, 1385 and 1356 cm-’ after heating to 473 K, but later decrease in intensity and disappear above 573 K. Under these conditions, CO, is detected in the gas phase. In the CH-stretching region one weak band is observable at 2880 cm-’ after interaction in the 423-573 K range. A rough assignment of these bands can be proposed at this stage, and will be refined below.The strongest bands at 1595 and 1400 cm-’ are typical of C02 asymmetric and sym- metric stretchings of carboxylate anions, while the split band at 1702, 1692 cm-’ is clear evidence of the formation of a carbonyl compound (C-0 stretching). The bands at lower frequencies are associated with CH deformations and/or C-C stretchings. The spectra show that propane interacts with the surface in the temperature range 373-573 K and undergoes paritial oxidation with a consequent reduction of the catalyst surface. Oxygenate species are produced, whose exact structure will be investigated by comparison with the spectra of any conceivable oxidation product of propane directly adsorbed onto the catalyst surface.These oxygenated species underwent further oxidation and burned in the tem- perature range 573-773 K, with production of gas-phase co, . On the basis of these observations we can conclude that oxidation of propane and propene produces adsorbed species with clearly different IR spectra (cf.Fig. 4 and 5). This sug- gests that two different oxidation paths operate for the two molecules. Identification of the Adsorbed Oxygenate Species and of the Oxidation Pathways The above data showed that partial oxidation products are formed from propene and propane at the surface of our cata- lyst. The gas-phase partial oxidation of propene is carried out industrially to produce a~rolein~~ acrylic acid.36 Other or products can be obtained in non-negligible yields over oxide catalysts, i.e.acet~ne,~*~~-~’ acetal-pr~pionaldehyde,~**~’ deh~de,~’acetic acid and C6 hydrocarbons (hexa-lS-diene and ben~ene).~~.~’ Propane catalytic oxidation also gives rise to a~rolein,~~ acrylic acid41 and propene,42 but acetone and acetic acid were also mentioned as by-products. To determine the nature of the adsorbed oxygenate species observed upon J. CHEM. SOC. FARADAY TRANS., 1994,VOL. 90 propene and propane oxidation over MgCr,04 , we pro-duced them via adsorption of the corresponding carbonyl compounds, carboxylic acids and alcohols. In Fig. 6 the spectra of the species arising from the adsorp- tion at room temperature of acrolein, acetone, propi-onaldehyde and acetaldehyde are reported.The bands assigned to the molecular adsorbed species are summarized in Table 1. In the case of acetaldehyde, strong bands of acetate species (1560, 1435 and 1350 cm-’) are already observed, showing its very fast oxidation. Bands arising from oxidized species are also evident, but weak in the cases of acrolein and propionaldehyde. In contrast the only ketone studied, acetone, is apparently almost stable under these con- ditions. The spectra of carboxylate species produced by adsorption on MgCr204 of formic, acetic and propionic acids are shown in Fig. 7. The spectrum of formate species [Fig. 7(a)] is char- acterized by bands at 2975,2880, 1602, 1390 and 1360 m-’. The four lower-frequency bands are due to CH stretching, asymmetric COO stretching, CH deformation and symmetric COO stretching, re~pectively.~~ The highest-frequency band is due to the combination v,,(COO) + S(CH).47 The band due to the asymmetric COO stretching actually exhibits shoul- ders on both sides, probably due to the presence of species adsorbed differently or on different sites.0.38 0.36 0.34 0.32 0.30 0.288 0.26 2 0.22 0.24 B 0.20fi rn 0.18 0.16 0.14 0.12 0.10 0.08 0.06 I’~‘I’’~1’”1”’1’”I”’I”’1’”I”’1’’.I’’’ 2200 2000 1800 1600 1400 1200 wavenumber/cm-’ Fig. 6 FTIR spectra of the adsorbed species arising from contact of MgCr,O,+, with acrolein (a), acetone (b), propionaldehyde (c), acetaldehyde (4,all after adsorption and outgassing at RT 2.0 1.8 -1 3200 2800 1700 1500 1300 1100 wavenum ber/cm -’ Fig.7 FTIR spectra of the adsorbed species arising from the adsorption of formic acid (a),acetic acid (b)and propionic acid (c) at RT and after outgassing at RT over MgCr,O, J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3351 Table 1 Position of the IR absorption bands of adsorbed and pure carbonyl compounds acrolein" acetoneb propionaldehyde' acetaldehyded ads. gas ass. ads. gas ass. ads. gas ass. ads. gas ass. E"} y"} .?il) 1754 1708 1724 v(c-0) 1731 v(C=O) v(c--0) 1743 v(C==O)1695 1710 .7.111 1613 1625 v(C-C) 1454 6,(CH,) 1467 1468 1427 1420 S(CH2Y 1435 1435 4&CH,)1422 1410 6,(CH3) 1413 1423 1366 1360 6[CH(0)lf 1369 1363 6,(CH3) 1395 1395 1347 1340 1281 1275 WH-) 1253 1254 -1176 1158 4C-C) ?"} 1215 v,(CCC) 1178 1138 1170 1232 1094 1090 w(CH,) 1096 1100 " Ref.43. Ref. 44.'Ref. 45. Ref. 46. Splittings due to the presence of different sites. The spectrum of acetate species [Fig. 7(b)]is characterized by strong bands at 1560 and 1435 cm-', the latter being more intense and sharper than the former, assigned to the asymmetric and symmetric COO stretchings. At the lower- frequency side a sharp peak at 1350 cm-'is observed, assign- ed to the symmetric CH, deformation, while a very weak band is found near 1055 cm-', probably due to methyl rocking.48 Acetate species are characterized by extremely weak CH stretchings, with a sharp band at 2935 cm-' and broader ones at 3015 and 2990 cm-'.The spectra of propanoate species are, as expected, more complex. In the V(CH) region three sharp and rather strong bands are observed at 2978, 2944 and 2886 cm-'. In the lower-frequency region, the strong and broad bands at 1560 and 1425 cm-' are associated, as always, with the asym- metric and symmetric COO stretchings. Additional sharp but strong peaks are found at 1471, 1380 (shoulder), 1301 and 1080 cm-',which can be assigned to CH, asymmetric and symmetric bending, CH, deformation (probably a rocking mode) and C-C ~tretching.~' The adsorption and oxidation of C1, C, and C, alcohols has also been investigated; some of these experiments gave valuable information for the identification of propane and propeme oxidation pathways. The spectra of the adsorbed species arising from allyl alcohol adsorption on MgCr,04 are shown in Fig.8. Most features correspond closely with the spectrum of adsorbed acrolein [Fig. qa)], i.e. strong C-0 stretching at 1690 cm-' (moving to lower wavenum- bers with decreasing coverage), weak C-C stretching 1613 cm-',sharp CH deformations at 1427, 1366 and 1281 cm-' (the last weak) and sharp C-C stretching at 1176 cm-'.'2*43 The additional bands at 1640 and 1021 cm-' can be assigned" to C-C and C-0 stretching of allyl alcoholate species. So allyl alcohol is partly oxidatively dehydrogenated to acrolein at RT. Further oxidation of acrolein gives rise to carboxylate species from RT to 423 K, when the aldehyde has almost disappeared.In the complex spectrum of the carbox- ylate species arising from allyl alcohol oxidation, features assignable45 to formate species (1600, 1390 and 1360 cm-') and acetate species (1560, 1435 and 1350 cm-') are certainly present. However, additional components like the sharp shoulder at 1640 cm-' [v(C=C)], the bands near 1500 and 1440 cm-' (COO stretchings) and the bands near 1370 and 1270 cm-'(the last a shoulder, CH deformations) correspond to those reported for acrylate species.' 2*5031 The spectra of the adsorbed species arising from contact of propan-2-01 with the catalyst are shown in Fig. 9. According to previous st~dies,'~.~~ the three bands at 1163, 1128 and ~,(CH,) WH2) -1433 6,(CH,) S[CH(O)] 1395 1395 6[CH(O)]w(CH2) 1346 1352 6,(CH3) t(CH2) 4C-C) 1179 1114 v(C-C), r(CH,) v(C-C) 1102 1107 y(CH3), r(CH,) Coupled modes." ' "''"l~ l~~l~~ ~ ~ ~l 2000190018001700160015001400 1300 1200 1100 wavenumber/cm-' Fig. 8 FTIR spectra of the adsorbed species arising from allyl alcohol adsorption over MgCr,O,_, at RT (a), and after successive outgassing for 30min at RT (b), 373 K (c), 423 K (4,473K (e),523 K (f),573 K (9) 1.1 \,-7-~-.7-#--F.-Y 2000190018001700 16001500 1400 1300 12001100 wavenumber/cm-' Fig. 9 FTIR spectra of the adsorbed species arising from propan-2- 01 adsorption over MgCr,O,+, at RT (a), and after successive out- gassing at 373 K (b), 423 K (c), 473 K (4,523 K (e),573 K (f)and 623 K (9) 3352 1103 cm- ' characterize the 2-propoxy species (coupled C-C and C-0 stretchings) produced by propan-2-01 dissociation, while the broad band at 1284 cm-' characterizes the undis- sociated adsorbed propan-2-01 (COH in-plane deformation).However, the sharp bands at 1710 and 1235 cm-', due to C-0 stretching and C-C-C asymmetric stretching of adsorbed acetone [Fig. 6(b) and Table 11, show that propan- 2-01 is oxidatively dehydrogenated at RT on the catalyst surface. The bands at 1467, 1385 and 1347 cm-' are associ- ated with CH deformations of the three species. After the sample is heated, the 2-propoxy bands disappear near 420 K when the acetone bands rise to their maximum intensity. Later, acetone bands also decrease in intensity and disappear, while the bands at 1580, 1435, 1385 and 1350 cm-' grow. These bands are due to a mixture of formate and acetate species [Fig.7(a) and (b)]. The most intense bands, assigned here to 2-propoxy groups and acetone, are also observed (with small shifts) in the spectrum of the species arising from propene transform- ation (Fig. 4); however, they are in this case very weak. In contrast, the strong bands at 1680, 1640 and 1030 cm-' observed there cannot be assigned to these C, species con- taining oxygen at C,. Finally, the complex spectrum of the carboxylate species arising from propene oxidation is defi- nitely different from that of the species arising from propan- 2-01 oxidation. This allows us to conclude that the pathway involving oxygen insertion at C2, definitely predominant on V20,-Ti0, oxidation catalysts,I2 is of only minor impor- tance for MgCr,O, .However, the most prominent features of the spectra of the adsorbed species arising from propene oxidation (Fig. 4) are definitely similar to those of the species arising from allyl alcohol oxidation (Fig. 8). The spectral region 1200-1000 cm-I is dominated by C-0 or coupled C-C/C-0 stretchings of alcoholate-like species. The spectra obtained after propene interaction at 373 K with the catalyst surface show a prominent sharp band at 1020 cm-', and a complex pattern in the region 1150-1090 cm-'. Comparison with the spectra of the surface species arising from ethanol and propan-1-01 adsorption (Fig.10) allows us to exclude the formation of ethoxy and 1-propoxy species from propene in significant amounts. In fact, both these species are characterized by a main sharp band near 1040 cm-' and other components in the range 1120-1040 cm-'. In contrast, the bands of alkoxy groups arising from propene are those typical of allyl alcoholate species [lo20 cm-', Fig. 7(a)] and 2-propoxy species [1150-1090 cm-', Fig. 6(a)]. Analogously, no evidence is found for adsorbed acetaldehyde and propionaldehyde in the spectra of the species arising from propene adsorption. In conclusion, com- parison of the spectrum of the adsorbed species arising from propene oxidation with those arising from allyl alcohol oxi- dation shows that the same species are formed and dominate the spectra in both cases, although 2-propoxy species and adsorbed acetone can also be found, in small amounts.In the spectra of the species arising from propane oxida- tion (Fig. 5), bands due to acetone (near 1700 and 1240 cm-') can be found. For this reason acetone adsorption and oxida- tion has been studied in detail (Fig. 11). At room tem-perature, adsorbed acetone gives rise to sharp and intense bands, as summarized in Table 1. In the temperature range 423-573 K acetone is converted to carboxylate species. Com- parison of the spectra in Fig. 1 l(d) and (e)with those reported in Fig. 7 shows that the former are due to a mixture of ace- tates and formates. At this temperature ketone species are no longer observed: the typical bands at 1710 and 1233 cm-' have disappeared.Obviously, the spectra of the species arising from acetone and propan-2-01 transformation are very similar (cf Fig. 8 and 11). J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1, l.,..lv1100 wavenumber/cm-' Fig. 10 FTIR spectra of the adsorbed species arising from contact of MgCr,O,+, with ethanol (a) and propan-1-01 (b)at room tem- perature Comparison of the spectra of the adsorbed species arising from oxidation of propane with those due to species arising from adsorption and oxidation of acetone strongly supports the assignment of the carbonyl species formed by propane oxidation as acetone. This is due to the close correspondence of most band positions, as well as of the further oxidation products of both molecules. The absence of the strong bands at 1470, 1300 and 1080 cm-' in the spectra of the adsorbed species arising from propane oxidation on the MgCr,O, surface clearly argues against the presence of propanoate species.On the other 0.20.41 0, c(D nin 2000 1800 1600 1400 1200 wavenumber/cm-' Fig. 11 FTIR spectra of the adsorbed species arising from acetone adsorption at room temperature over MgCr,O, and following evac- uation at RT (a),423 K (b),473 K (c), 523 K (4,573 K (e),623 K (f), 673 K (9) J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 3.41 3.03*21 2.8 1 y 2.6/2.42*2L2.0 11.84 8 ' ' ' I . ' ' 2000 1800 1600 1400 1200 ' 1000 waven u m ber/cm -Fig.12 FTIR spectra of adsorbed species arising from formic acid adsorption over MgCr,O, at 373 K (a)and following decomposition at 473 K (b)and 573 (c) hand, the absence of a sharp v(C=C) band in the region near 1635 cm-' and of a COO asymmetric stretching component near 1500 cm -characterizing adsorbed acrylate species (see above), precludes the presence of acrylate and, consequently, C, carboxylate species are apparently not formed at all. Simi- larly, no trace is found of the C, aldehydes acrolein and pro- pionaldehyde (cf: Fig. 5 and 6). In contrast, the spectra of carboxylate species arising from both propane and acetone oxidation can easily be obtained from a sum of the spectra of acetate and formate species. From this analysis we conclude that carboxylate species are obtained by the oxidative cleav- age of the C(l)-C(2) bond of acetone and propane. Thermal Evolution of Carboxylate Species The study of the species arising from adsorption of C,-C, acids as well as of the corresponding aldehydes and alcohols allowed us to determine the stability range of these adsorbed species on MgCr,O,.For completeness, we also studied the stabilities of the carboxylate species produced by propane oxidation, i.e. formates and acetates. As mentioned above, the IR spectrum of the formate species shows, after adsorption at room temperature, three main bands (1602, 1390 and 1360 cm-'). Above 473 K these bands quickly decrease in intensity (Fig. 12) giving rise to carbon oxides in the gas phase and Q1 C e n8 bands at 1500 and 1350 cm-' assignable to surface carbonate species,26 while there is a very slight decrease in the intensity of the Cr=O stretching bands.This supports the idea that formate decomposes following the reaction : HCOO--+ CO + -OH. This reaction is a simple decompo- sition, and does not imply catalyst reduction. However, some of the CO can be oxidized to CO,, with consequent catalyst reduction. The evolution of acetate species with temperature is shown in Fig. 13. Acetates, characterized by bands at 1560, 1435, 1350 and 1055 cm-' (see above), are substantially stable up to 573 K, while they progressively disappear in the range 573-773 K, giving rise to bands due to carbonate species and to gas-phase carbon oxides.However, their disappearance is accompanied by a strong decrease of the Cr=O stretching bands in the region 1000-800 cm-'. This provides evidence that the disappearance of these species is due to their overox- idation to carbon oxides, with a parallel reduction of the catalyst surface. Our data show that formate species undergo decomposi- tion below 573 K, so they are very unstable at the tem- perature at which propane begins to be oxidized. Acetate species, instead, burn at higher temperatures with a conse- quent reduction of the catalyst; they are almost stable up to above 573 K. Discussion Mechanism of Propene Oxidation at the MgCr,O, Surface The present results show that propene reacts with the surface of oxidized MgCr,O, at room temperature, but this reaction becomes more evident near 373 K when adsorbed acrolein and allyl alcoholate species are observed, although acrylate species (and perhaps other carboxylate species) have already appeared.However, acrolein and allyl alcohol in the tem- perature range 373-473 K are rapidly oxidized, giving rise to acrylate species as the predominant products, which become unstable above 523 K, being burnt. Thus, we have provided evidence for the main oxidation pathway of propene on MgCr,O, (Scheme 1). This pathway implies the breaking of the allylic C-H bond with a subsequent oxygen insertion to give allyl alco- holate species, which are, in fact, observed oia the detection of + C H2= CH-C HZO H It'\ 2000 1800 1600 1400 1200 1000 wavenumber/cm-' Fig.13 FTIR spectra of adsorbed species arising from acetic acid co -co2 adsorption over MgCr,O, at 473 K (a),573 K (b),673 K (c), 773 K Scheme 1 Propane and propene oxidation pathways and their con- (4 nection at the surface of the oxidized MgCr,O, catalyst the typical sharp C-0 stretching band at 1020 cm-'. Further evolution strictly parallels allyl alcohol oxidation. The key step in this mechanism is the breaking of the allylic C-H bond. The active site in our case is identified as a Cr"+=O site (n= 5 or 6), which should be reduced to CF2)+.A likely intermediate in this process is the forma- tion of an allyl species, either anionic or radical, as proposed for the selective oxidation of propene to acrolein.2 However, such an intermediate can be observed only if its formation is faster than its further reaction with surface oxide species to give allyl alcoholate.Anionic allyl species have been unequivocally identified on Zn0,53 and 21-0,, 34 which are not active allylic oxidation catalysts (so the subsequent oxygen insertion should be slow or may not occur at all) and have a strong basic character. On true allylic oxidation cata- lysts such as uranyl antimonate,', acrolein was found upon propene adsorption, but neither the allyl intermedite nor allyl alcoholate were detected spectroscopically. This is prob- ably because anionic or radical-like allyl species are oxidized so rapidly to allyl alcoholate species that they cannot be observed.This implies that the rate-determining step is the first hydrogen abstraction from the methyl group. Also allyl alcoholates can be rapidly dehydrogenated to acrolein at the temperature at which they form, so their detection can be difficult too. This situation differs from that found in the case of toluene oxidation to benzaldehyde on vanadia-titania.' In this case, the rate-determining step is probably the reac- tion of the intermediate benzyl species with surface oxygen to give benzyl alcoholates, which are later rapidly dehydroge- nated. For this reason benzyl species are observed while benzyl alcoholates are not." The reverse is observed here for allyl and allyl alcoholate species. Besides the allylic oxidation pathway at C(l), another propene oxidation pathway is evident on MgCr,O, (although it is definitely of minor importance) with the pro- duction of species oxidized at C(2) (Scheme 1).According to our previous this pathway is associated with the electrophilic attack of weakly Bransted acidic OH groups of the alkene double bond, according to the Markovnikov rule, giving rise to secondary propoxy species that are later oxi- dized to acetone, and finally undergo oxidative breaking of the C(2)-C/3) bond leading to acetate and formate species. This pathway (called oxidative hydration2) was shown to be predominant on vanadia-titania," which is less active as an oxidation catalyst and gives mixtures of acetone and acrolein upon propene ~xidation.'~ Mechanism of Propane Oxidation at the MgCr,O, Surface The present data show that propane gas interacts with the surface of oxidized MgCr,O, starting from above 373 K giving rise to adsorbed oxygen-containing species that can further be oxidized to carbon oxides.Complete burning of hydrocarbon and oxygenated organic compounds by the oxi- dized MgCr,O, catalyst is observed at 773 K. The data pre- sented here also give a relatively simple picture of the propane combustion pathway on this surface. In fact, the first detectable adsorbed oxidation product is acetone, which is later converted into a mixture of formate and acetate species, by oxidative cleavage of the C(l)-C(2) bond. Formate species rapidly decompose while acetate species burn only near 773 K (Scheme 1, left-hand side).Surprisingly, com- pounds oxidized at C( 1) like acrolein and acrylate species are very evident upon popene oxidation under the same condi- tions, but are not found at all with propane. So, the pathways of propane and propene oxidation (Scheme 1, right-hand side) are considerably different. This supports the data reported several years ago by YaoS7 which showed that the J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 oxidation pathways of alkanes and alkenes are different on a-Cr203. The formation of acetone from propane implies propane activation at C(2). This agrees with the lower dissociation energy of C-H at secondary carbons (-CH,- methylene groups, 94 kcal mol-') with respect to C-H at primary carbon atoms (methyls, 99 kcal mol- ,).Accordingly, also the C-H bond dissociation energy at the primary carbon in the case of propene is in agreement with the lower energy for C-H bond cleavage at an allylic methyl group (77 kcal mol-') with respect to that of a vinylic group (105 kcal mol-'). The lower temperature at which propene oxidation to acrolein is observed, with respect to propane oxidation to acetone, is justified by the lower C-H bond dissociation energies of allylic methyls with respect to alkane methylenes. It seems reasonable, although it is not strictly proved here, that cleavage of the first C-H bond of propane is followed by insertion of an oxygen before the second C-H bond at C(2) is broken.This hypothesis is confirmed by the detection of allyl alcoholates from propene, as well as by studies which are now in progress5' concerning the oxidation of other hydrocarbons. Therefore, 2-propoxy species should be the first intermediates in propane oxidation, although they cannot be observed easily because at the temperature at which their formation is sufficiently fast, their further trans- formation to acetone is even faster, making their concentra- tion zero. The active site for C-H bond scission is certainly associ- ated with the high-valency Cr ions responsible for the Cr=O stretching bands in the region 1000-800 cm-'.It is not excluded that the sites active for propane activation do not entirely coincide with those active for propene oxidation.In fact, it is possible that only the most active sites can attack both molecules, the least active ones being able to attack propene only. On the other hand, sites able to break the C-H bonds of the propene methyl groups can cooperate with sites able to interact with the alkene double bond. Acetone is a rather reactive molecule, and this causes its strong adsorption and further oxidation at the C(2)-C(3) bond to formate and acetate species. The last stages in propane combustion correspond to the evolution of acetate and formate species. However, the behaviour of these two species is different. Formates decompose rapidly, giving CO, while acetates need further oxidation to give finally CO, . Comparison with Other Catalytic Systems: Activity-Selectivity in Heterogeneously Catalysed Oxidation of Propene and Propane The data reported above concern surface reactions observed to take place from an oxidized MgCr,O, catalyst surface with propane and propene.However, they do allow a devel-opment of the theory of catalytic hydrocarbon combustion mechanisms on metal oxides. The conditions under which these interactions are observed (373-773 K) agree roughly with those at which this compound was found to act efi- ciently as a combustion catalyst for propene and light alkanes.15*58Our data show that propene combustion can be obtained in the absence of gaseous oxygen, totally at the expense of the Cr=O oxidized sites of the catalyst, which in the 573-773 K range are quickly reoxidized by gas-phase oxygen.This strongly supports the idea that catalytic com- bustion of propene and propane over this catalyst can occur completely with a Mars-Van Krevelen type me~hanism.'~ Our data also show that adsorbed partial oxidation pro- ducts are formed upon propene and propane oxidation at the surface of MgCr,O,, well known to be essentially a non- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 selective combustion catalyst. These adsorbed partial oxida- tion products are intermediates in the formation of carboxylate species, already recognized as intermediates of propene combustion over chromia-based catalysts.' 'v60 The data concerning propene oxidation on this com-bustion catalyst can be discussed in relation to previous data obtained on selective oxidation catalysts.' '-14 Similar surface reaction pathways are observed for selective and non-selective oxidation catalysts.However, important differences can be found in the temperature at which the same steps occur on different catalysts. Propene activation is observed to commence on MgCr,O, at 373 K, while on Mg3(V0,), it starts at near 523 K.'4*61The much lower reactivity of Mg vanadate with respect to Mg chromite with propene explains why the former acts as a selective catalyst for propane oxida- tive dehydrogenation to propene,42 while the latter is anon-selective combustion catalyst for both propane and propene.' Acrolein is also definitely oxidized on MgCr,O, at 373 K, while it is much more stable on the surfaces of V,0S-Ti02 l2 and MOO,-TiO,.62 This could explain why these catalysts are rather selective in the oxidation of propene to a~rolein~~ while MgCr,O, is not. The reaction pathway proposed in Scheme 1 implying propane activation at C(2) on MgCr,O, also agrees with the behaviour reported for another chromite, ZnCr,O,, as a catalyst for n-butane oxidation. In fact, over this catalyst complete butane combustion has been observed at 623 K, while at 493 K butane oxidation gives acetic acid and methyl vinyl ketone [i.e. compounds functionalized at C(2)] with rather high ~electivity.~~ On the other hand, acetic acid is frequently reported as a product of propane oxidation over oxide catalyst^:^' this could be the result of activation at C(2) and agrees with the stability of acetate species to overoxida- tion, evidenced above, allowing desorption of acetic acid.According to Scheme 1, propene is not an intermediate in the main propane oxidation path, so that propene and propane oxidation follow predominantly two different path- ways on the MgCr,O, surface. On the other hand, our data suggest that decomposition of 2-propoxide groups (the first intermediate starting from propane) to propene and an OH group is competitive with its oxidative dehydrogenation giving rise to acetone (Scheme 1). This competition could provide a 'cross-roads' between alkane oxidative dehydroge- nation and oxidation at C(2). On MgCr,O,, the dehydroge- nation of 2-propoxy groups to acetone is much faster than their decomposition to propene, which is consequently not an intermediate for further oxidation.This makes MgCr,O, an oxidation (combustion) catalyst rather than an oxidative dehydrogenation catalyst for propane. This picture finds support in the comparison between the behaviour of MgCr,O, and Mg vanadate catalysts, which are completely opposite. On Mg vanadates, in fact, alcoholate species decompose quickly at very low temperature, giving only traces of oxidized species. Therefore, the equilibrium propene s2-propoxide is displaced left (towards the alkene) on Mg vanadates, in contrast to Mg chromite where it is shifted to the right (towards the alkoxide). This agrees with the behaviour of Mg vanadates as oxidative dehydrogenation catalysts for production of propene from propane.,, Comparison of the behaviour of the different metal oxide catslysts allows us to apply the generalization of the mecha- nism of Scheme 1 to all of them.The behaviour of an oxida- tion catalyst is strongly influenced by its acid-base properties. In fact, the decomposition of the alcoholate species giving rise to the alkene is a purely acid-base reac-tion. According to our data, the equilibrium propene e2-propoxide is displaced left (towards the alkene) on both very basic and very acidic catalysts (in fact, alcohol dehydration is either acid- or base-catalysed66), while it is shifted right (towards the alkoxide) for catalysts with medium Brsnsted acidity.In fact, Mg chromite shows Brransted acidity while Mg vanadate does not.61 The propane oxidative dehydrogenation pathway (through propene) can be followed by allylic oxidation, giving rise to acrolein and/or acrylic acid. Therefore, catalysts with no Brsnsted acidity but strong Lewis acidity (allowing decompo- sition of 2-propoxy groups to propene) and very active allylic activation centres are the best candidates for propane oxida- tion to acrolein. These data support the idea that the main difference between a selective and a non-selective oxidation catalyst is in its ability to overoxidize selected products. This leads us to propose that a combustion catalyst is one that is able to perform selective oxidation, but from which the partial oxidation product is unable to be desorbed without further oxidation.This is supported by two observations : (i) Efficient partial oxidation processes from hydrocarbons are limited to the production of compounds such as acrolein, acrylonitrile, maleic and phthalic anhydrides, which have intrinsic chemical stability owing to their ability to delocalize ionic charges. Compounds whose ionic charges are more localized (like acetone or the non-conjugated aldehydes, with respect to acrolein) are more reactive towards the oxide cata- lyst surface (which always contains electrophilic and nucleo- philic sites) and are desorbed less easily. (ii) Basic dopants frequently increase the selectivity but decrease the activity of partial oxidation catalysts, becase they lower the adsorption strength of the oxygenate compounds (which are always Lewis bases) on the surface (which always contains Lewis acid sites), but this tends to limit the activity for C-H bond scission (which needs Lewis acidity',).Therefore, the main requirement of a partial oxidation catalyst is inactivity towards overoxidation of the desired product, in spite of being a poorly active catalyst. This is the case for V2O5- and Moo3-based catalysts, which are gener- ally much less active than typical combustion catalysts, like, among others, those based on Cr203. This view is only partially in disagreement with that of Haber2*67 which associates selective oxidation with nucleo- philic oxygen species (0'-lattic oxide ions) and combustion with electrophilic oxygen (O,, 0,-and O,,-).Our study confirms that lattice oxygen performs selective oxidation, but suggests that it may also be involved in non-selective com- bustion pathways. This does not preclude the involvement of electrophilic oxygen (together with nucleophilic oxygen) in a different combustion pathway. Conclusions The FTIR study of propene interaction with the surface of oxidized MgCr,O, shows that the alkene is totally oxidized in the temperature range 573-773 K at the expense of surface Cr"+=O (n = 5 or 6) species. However, this occurs through two different 'selective oxidation ' pathways at the surface, one of which involves activation at C(1) via ally1 alcoholate species, and the other activation at C(l) via 2-propoxy species.Combustion seems to involve essentially successive overoxidation of the partial oxidation products. Propane is oxidized at the same surface to acetone, which further transforms into acetates and formates and, finally, into carbon oxides; this occurs in the 573-773 K range. This oxidation pathway is completely different from the main one undergone by propene. The surface reactions observed allow us to propose a detailed reaction pathway for both selective and unselective oxidation of propane and propene. On the basis of the comparison of these data with data available for 3356 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 other catalytic systems we propose this pathway as a general 34 G.Busca, V. Lorenzelli, G. Ramis and V. Sanchez Escribano, mechanism for alkane and alkene catalytic oxidation. Muter. Chem. Phys., 1991,29, 175. 35 R. K. Grasselli and J. D. Burrington, Adu. Catal. Relat. Subj., 1981,30,133.This work has been supported in part by MURST (Rome, 36 K. Waissermel and H. J. Arpe, Industrial Organic Chemistry, Italy) and by NATO (CRG n. 900463.). Verlag Chemie, Weinheim, 1978, p. 256. 37 Y. Moro-Oka, S. Tan and A. Ozaki, J. Catal., 1968,12,291. 38 N. Giordano, N. Meazzo, A. Castellan, J. C. J. Bart and V.References Ragaini, J. Catal., 1977,50, 342. 1 Chemical Processing Handbook, ed. J. J. McKetta, Marcel 39 R. F. Howe and H. Minming, in Proc. 9th Int. Congr. Catal., Dekker, New York, 1993. Calgary, The Chemical Institute of Canada, Ottawa, 1988, p.2 A. Bielanski and J. Haber, Oxygen in Catalysis, Marcel Dekker, 1585. New York, 1991. 40 Y. C. Kim, W. Ueda and Y. Moro-oka, in New Developments in 3 New Developments in Selective Oxidation 11, ed. V. Cortes Corb- Partial Oxidation, ed. G. 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ISSN:0956-5000    

DOI:10.1039/FT9949003347

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3357-3365 Oxidative Coupling of Methane over La,O, Influence of Catalyst Preparation on Surface Properties and Steady and Oscillating Reaction Behaviour V. R. Choudhary" and V. H. Rane Chemical Engineering Division, National Chemical Laboratory, Pune 4 1 1 008, India The catalyst precursor and calcination conditions used in the preparation of La,O, have been found to have a marked influence on the surface properties, basicity and base strength distribution, and catalytic activity and selectivity in the oxidative coupling of methane (OCM) under different reaction conditions. Among the La,O, catalysts, only that prepared from lanthanum acetate by decomposition in N, (but not in 0,)showed symmetric temperature and concentration oscillations in OCM at or below 933 K over a narrow temperature range.The unsteady reaction behaviour is found to be very complex and strongly dependent upon the OCM process param- eters and the catalyst parameters. Furthermore, the results on the La,O, obtained from the acetate indicated a strong possibility of the formation of new catalytic sites active at lower temperatures (below 873 K) during the OCM reaction at higher temperatures. Earlier studies'-7 showed that La203 has high activity and by drying at 393 K for 12 h. The lanthanum acetate (Aldrich) selectivity in OCM to C, hydrocarbons. La203 catalysts and lanthanum nitrate (GR, Loba) were ground with deion- with promoters such as Li,8 LiC1,' Na, , Sr and BaI8 ized water sufficient to form a thick paste and dried at 393 K have been studied as well as La-promoted MgO and for 12 h.The lanthanum hydroxide, lanthanum carbonate (I Ca0.'9-22 Taylor and Schrader6 reported that La203 and 11) were prepared by precipitation from an aqueous solu- obtained from different precursors showed different catalytic tion of ammonium hydroxide, sodium carbonate or ammon- performance in the OCM process. Recently we have observed ium carbonate, respectively, at pH 10-11 at room unsteady reaction behaviour, with periodic fluctuations in temperature. The precipitate was washed with deionized reaction temperature and concentration indicating symmetric water until free from cations and anions, and dried at 393 K oscillations, in OCM (above 823 K but below 973 K) over for 12 h.The dried catalyst mass was decomposed at 873 K La203 obtained from lanthanum acetate by thermal decom- for 6 h in static air, pressed binder-free and crushed to 22-30 position in N,.23We have now investigated the influence of mesh size particles, then calcined at 1023 and 1223 K in a precursor and reaction conditions on the bulk and surface flow of N, or 0, (12000 cm3 g-' h-' ). The calcination con- properties of La,03 and also on the catalytic activity and ditions of the catalysts are given in Table 1. selectivity and unsteady reaction behaviour in OCM under The surface area of the catalysts calcined at 1023 and 1223 different process conditions. K was determined by the single-point BET method by mea- suring the adsorption of nitrogen (30 mol% in He) at liquid- nitrogen temperature, using a Monosorb surface-areaExperimental analyser (Quanta Chrome Corp.).The crystal size and mor- The La,O, catalysts (Table 1) were prepared by thermal phology of the catalysts were studied by scanning electron decomposition of different precursors. The hydrated La,O, microscopy (SEM). The crystal phases were studied by was prepared by treating powdered La203 (Aldrich) with powder X-ray diffraction (XRD). deionized water (2 ml g-') on a water bath for 4 h while The acidity distribution on the catalysts was determined by maintaining a constant water content of the slurry, followed temperature-programmed desorption (TPD) of ammonia Table 1 Catalyst precursors and calcination conditions calcination conditions" surface area cat a1 y st precursor TIK atmosphere* /mz g-' Ia Ib IIa hydrated La203 hydrated La203 lanthanum acetate 1223 1023 1223 3.8 6.3 2.8 IIb lanthanum acetate 1023 4.5 IIC lanthanum acetate 1023 4.4 IIIa lanthanum nitrate 1223 1.7 IIIb lanthanum nitrate 1023 4.4 IVa IVb Va Vb VIa VIb lanthanum carbonate (I) lanthanum carbonate (I) lanthanum carbonate (11) lanthanum carbonate (11) lanthanum hydroxide lanthanum hydroxide 1223 1023 1223 1023 1223 1023 0.4 3.0 2.0 2.1 2.7 22.1 " Before calcination, catalyst precursor was decomposed at 873 K in static air for 6 h.The period of catalyst calcination was 2 h. Space velocity, 1200cm3 g-' h-'. (chemisorbed at 323 K) on the catalyst (0.5 g) from 323 to 1223 K at a linear heating rate of 20°C min-' in a flow of moisture-free helium (20 cm3 min- ') in a quartz reactor.The desorbed ammonia was detected with a thermal conductivity detector and also measured quantitatively by chemical analysis. The basicity and base-strength distribution on the catalysts were determined by measuring the step-wise thermal desorp- tion (STD) of CO, from the catalyst (0.5 g) in a quartz reactor, from 323 to 1173 K in a number of successive tem- perature steps (i.e. 323-423, 423-573, 573-773, 773-973 and 973-1 173 K). After the maximum temperature of the respec- tive step was attained, it was maintained for a period of 30 min to- allow desorption of the CO, adsorbed reversibly on the catalyst at that temperature. The amount of CO, desorbed in each step was determined gravimetrically by absorption in an aqueous barium hydroxide solution.The detailed procedure for measuring the base-strength distribu- tion by the STD of CO, and the estimation of CO, chemi-sorption data from the STD data have been described earlier.7*24 The CO, chemisorption data reported here are presented after subtracting the CO, content of the catalyst, which was determined by measuring the CO, evolved when the catalyst (after pretreatment at the calcination temperature in an He flow for 1 h) was heated from its calcination tem- perature to 1273 K in a flow of pure He for 1 h. Throughout, the chemisorption is considered as the amount of adsorbate retained by the presaturated catalyst after it was swept with pure He or N, for a period of 30 min.The steady/unsteady OCM reaction over the catalysts was carried out in a tubular quartz flow reactor packed with cata- lyst particles (0.1-0.2 g) between quartz wool plugs. The J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 reactor was kept in a vertical tubular furnace. The reaction temperature (controlled by a digital proportional tem-perature controller) was measured by a chromel-alumel thermocouple located in the catalyst bed. The temperature under unstcady conditions was measured as a function of reaction time. The reaction was carried out under following conditions: amount of catalyst, 0.1-0.2 g; feed, pure CH,-0, ; CH4/0,, 3.0-8.0; space velocity, 51 600-103 200 cm3 g- ' h-and temperature, i.e.reactor temperature, 673-1123 K. The product gases after the removal of water by conden- sation at 273 K were analysed by an on-line gas chromato- graph using Porapak-Q and Spherocarb columns. The concentration of 0, in the product stream was recorded con- tinuously by an on-line paramagnetic 0,-analyser (Oxymat I, Fuji Electric). High-purity gases He ( >99.99%), CH, (99.995%),CO, (99.995%),0, ( >99.5%)and NH, (99.99%)were used. Before being used the catalysts were pretreated in situ at their calcination temperature in a flow of He (20 cm3 min- l) for 1 h. Results and Discussion Surface Properties The surface areas of the catalysts are included in Table 1.The surface area of La,03 is strongly influenced by the precursors and the calcination temperature. The SEM micrographs of La,03 (IkVIb) are presented in Fig. 1 and that of La,03 (IIa) and La,03 (IIc) in Fig. 2. A comparison of the SEM photographs of La203 (I&-VIb) H1 prn Fig. 1 SEM micrographs of La,O, (a)Ib, (b)IIb, (c) IIIb, (d) IVb, (e)Vb and cf)VIb J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 (a1 H1 pm Fig. 2 SEM micrographs of La,O, obtained by calcination of La acetate (a) at 1023 K in 0, flow for 2 h and (b) at 1223 K in N, flow for 2 h shows a strong influence of the precursor on the crystal size and morphology. Furthermore, a comparison of the SEM micrographs of La203 (IIa, IIb and IIc) indicates that the crystal size and morphology of La203 (11) is influenced by the temperature and gas atmosphere used in the catalyst calcina- tion.XRD analysis of the La203 (I-VIa) catalysts obtained by the decomposition and calcination of the different catalyst precursors at 1223 K has indicated the presence of only pure La,03 crystalline phase. The La,03 (Ib, IIIb and VIb) cata-lysts were also found to contain only the La203 phase but the La203 (IIb, IVb and Vb) obtained from lanthanum acetate and lanthanum carbonate also showed the presence of minor amounts of La,(C03), and Laz0,C03 (lanthanum oxycarbonate). The XRD data have been presented else- where., The XRD results indicate no lanthanum carbonate or oxy- carbonate to be present in the catalysts calcined at 1223 K.However, these carbonates are not completely decomposed at 1023 K and hence are retained to a small extent in the bulk of the catalysts obtained from the decomposition of lantha- num acetate or carbonates. These observations are quite con- sistent with earlier st~dies.~~~~~~' The COz content of IIb, IVb and Vb was found to be 0.56, 0.75 and 1.35 mmol g-', respectively. The COz content of Ib, IIIb and VIb was found to be negligibly small. Ia-VIa showed no evolution of COz when heated up to 1373 K. The values of basicity (STD and chemisorption of CO,) for these catalysts are reported in this work after subtracting the values of their COz content. Fig. 3 shows the TPD curves of NH, from the catalysts, from their initial surface coverage by ammonia (ei) which cor- responds to the total acidity.It is seem that the acidity (measured in terms of NH, che-misorbed at 373 K) of La203 is strongly influenced by the precursor used in the catalyst preparation. The TPD curves (Fig. 3) show that the acid strength distribution on La203 is also very strongly influenced by the catalyst precursor. The presence of more than one peak in the TPD for all the cata- lysts reveals that there is more than one type of site for NH, 323 523 723 923 1123 1323 TI Fig. 3 TPD of ammonia on (a)Ia, Bi = 1.69 mmol g-'; (b) IIa, Bi = 0.69 mmol g- ;(c) IIIa, Bi = 0.24 mmol g-;(d)IVa, Bi = 0.23 mmol g-'; (e) Va, Bi = 0.18 mmol g-'; (f) VIa, Bi = 0.17 mmol g-I chemisorption. The ammonia chemisorption site (i.e.acid site) on the catalyst is expected to be a surface La3+. Thus, the results indicate that surface La3 + in different coordi- nations (3, 4 and 5 coordination) are present on the catalysts and their relative concentration is strongly influenced by the precursor. It is interesting to note that Ia has the highest acidity and contains only strong (major) and very strong (minor) acid sites compared with the other catalysts that are less acidic and also differ in their total acidity and acid strength dis- tribution. The temperature dependence of CO, chemisorption on the La,03 catalysts calcined at 1023 and 1223 K is shown in Fig. 4. The chemisorption of CO, at higher temperature points to the involvement of stronger basic sites.Hence, the CO, chemisorption us. temperature curves present the type of site energy distribution in which the number of sites are expressed in terms of the amount of CO, chemisorbed as a function of chemisorption temperature. The results (Fig. 4) indicate that the basicity distribution on the catalysts is very broad and is strongly influenced by both the precursor and the calcination temperature. The observed base-strength distribution on the catalysts is expected to be due to the presence of surface 0,-sites (i.e. basic sites) in different coordinations (3, 4 and coordination). The strong basicity is attributed to the low- coordinated 0,-surface sites. The total concentration of basic sites (measured in terms of CO, chemisorbed at 323 K) is decreased on increasing the calcination temperature.This is expected to be mostly due to decrease in the surface area by catalyst sintering. Whereas, the larger decrease in the strong basic sites (measured in terms of C02 chemisorbed at 773 K), relative to that of the total basic sites, on increasing the calcination temperature (Table 2) is mostly due to the removal of crystal defects which results in a decrease in low-coordinated surface 0,- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1 I I I I I I using very high space velocity (103 200 cm3 g-' h-l). The (a1 results showing the influence of temperature on the methane 0.8 -conversion, C, selectivity, C2HJC,H, and CO/CO2product ratios are presented in Fig. 5-7.The following general obser- vations have been made : (1) For all the catalysts the ethene/ethane ratio is increased with increasing temperature. This is consistent with earlier studies.7.2 5.28-3 1 (2) The temperature dependence of the CO/CO, ratio in the products is different for different catalysts. The CO/CO2 ratio is decreased for IIa and IV-VIa but passes through a minimum or maximum for Ia and IIIa, depending upon the CH4/02ratio, with increasing temperature. (3) For IIa and IV-VIa the selectivity is increased with increasing temperature, the increase being very large for IVa. However, for Ia and IIIa the selectivity is decreased for CH4/02 = 4.0 but increased for CH4/02 = 8.0 with increas- ing temperature. (4) In general, the conversion is increased with increasing temperature.The increase is, however, very small for I-IIIa and Va at CH4/02 = 4.0. I I I I (b1 30r1 0 373 473 573 673 773 873 973 T/K Fig. 4 Temperature dependence of chemisorption of CO, on (a)b catalysts and (b)a catalysts. 0,I; 0,11; A, 111; V,IV; A, V and V, VI. The observed changes in the surface acidity and basicity with changes in the preparation conditions are due to changes in the coordination number of surface La3 + and 02-and probably also to the modification of the habits of the microcrystals produced with different crystal surface plane abundances. I 1 I I 973 1023 1073 1123 973 1023 1073 11: Steady OCM T/K T/K The OCM over I-VIa was carried out at 973-1123 K and Fig.5 Temperature dependence of (a)methane conversion and (b) Ia; 0,IIa; A, IIIa; A, IVa; 0,Va and V,VIawith a CH4/02 ratio of 4.0 and 8.0 at atmospheric pressure C, selectivity of 0, Table 2 Surface properties and catalytic activity and selectivity in OCM at 1023 and 1223 K, CHJO, = 4.0 and space velocity 103 200 Cm3 g-'h-' basicity/mmol g -' catalytic properties surface area acidity" CO, content catalyst /m2 g-' /mmol g-' /mmol g-' totalb strong' CH, conversion C, selectivity C, yield C,HJC,H, calcined at 1023 K I 6.3 0.00 0.346 0.227 24.0 55.5 13.3 1.03 I1 4.5 0.56 0.428 0.371 26.5 41.7 9.5 1.15 Ill 4.4 0.00 0.161 0.063 23.7 54.1 12.8 0.85 IV 3.0 0.75 0.602 0.567 22.2 55.2 12.3 1.60 V 2.1 1.35 0.710 0.640 23.7 51.4 12.2 1.oo VI 22.1 0.00 0.643 0.463 25.3 53.0 13.4 0.98 calcined at 1223 K I 3.8 1.69 0.00 0.091 0.056 25.0 57.0 14.3 1.11 I1 6.3 0.69 0.00 0.205 0.040 24.8 51.5 12.8 0.90 111 1.7 0.24 0.00 0.091 0.019 26.2 56.8 14.9 0.85 rv 3.4 0.23 0.00 0.148 0.101 4.2 28.6 1.2 0.14 V 2.0 0.18 0.00 0.126 0.090 27.2 55.6 15.1 0.90 V1 2.7 0.17 0.00 0.126 0.074 25.6 54.7 14.0 0.86 Measured in terms of NH, chemisorbed at 373 K.Measured in terms of CO, chemisorbed at 323 K. 'Measured in terms of CO, chemi- sorbed at 773 K. J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 0.4t /d P 0" 0.8 FO 0.61 I O.10.2 Ol I I 1 I 923 973 1023 1073 1123 T/K Fig. 6 Temperature dependence of ethene/ethane ratio over 0,la; e,IIa; A, IIIa; A,IVa; V, Va and V,VIa.CH,/O, :(a)4 and (b) 8. The above observations reveal a strong influence of cata- lyst precursor on the catalytic activity and selectivity in the OCM process. The results (Fig. 5 and 6)also show that an increase in the CH4/02 ratio from 4.0 to 8.0 causes a decrease in the conver- sion and the ethene/ethane ratio but the selectivity is gener- 1.6-1.2--0.8 0.4-N I I I 1 923 973 1023 1073 1123 T/K Fig. 7 Temperature dependence of the CO/CO, product ratio over 0,Ia; 0,IIa; A, IIIa; A,IVa; V, Va and V,VIa. CHJO,: (a)4 and (b) 8. ally increased. The influence of the CH4/02 ratio is quite similar to that observed for OCM over rare-earth metal oxides,' Mg029 and several other catalyst^.^^.^' The results of OCM over I-VIb at 873-1023 K, CH4/02 ratio of 4.0 and flow of 103 200 cm3g-h-' (at 1 atm) are presented in Table 3.They show that, in OCM over catalysts calcined at 1023 K, the reaction temperature has a strong influence on the conversion and product selectivity. The methane conversion, C, selectivity and ethene/ethane ratio increase but the CO/CO2 ratio decreases with increasing reaction temperature. The increase in the ethene/ethane ratio with decreasing CH4/02 ratio is most probably due to the availability of 0, at higher concentration for the following gas-phase reactions involved in the formation of ethyl radicals and ethene from et hane.32,3 C,H6 + 0, +C,H,' + HO,' (1) C2H5' + 02 --* C2H4 + HO2' (2) C2H6+ HO,' -+ C,Hs' + H,O, (3) H,O, + Z -+ 20H' + Z (4) (where Z is a third body, e.g.water molecule) C2H,j + OH' +C2H5' + H,O (5) Ethane is expected to be formed by gas-phase coupling of methyl radicals. 34 The increase in the ethene/ethane ratio with increasing temperature is expected to be due to the decomposition of ethyl radicals formed in reactions (l), (3) and (5) and thermal cracking of ethane at the higher temperatures CZHs' +C2H4 + H' (6) CzH6 +C,H4 + H2 (7) It may also be due to the increase in the rate of the gas-phase reaction of ethyl radicals to form ethene [reaction (2)] and the oxidative dehydrogenation of ethane on the catalyst surface. The increase in the C, selectivity with increasing tem- perature is expected to be mostly due to a decrease in the formation of carbon oxides from methyl radicals by the fol- lowing gas-phase reaction.34 CH,' + 02+CH,O,' + +CO, CO, (8) The formation of methylperoxy radicals (CH302'), which leads to CO and CO,, is not favoured at higher temperature~~,~~and hence the C, selectivity is expected to increase with the temperature.However, the decrease in the selectivity at the higher temperatures in some cases (Fig. 5) is attributed mostly to the conversion of methyl radicals to CO and CO, on the catalyst surface. To a small extent, it may also be due to combustion of ethane and ethene in the gas phase and/or on the catalyst surface at the higher tem- peratures. Note that unsteady reaction behaviour showing oscil- lations has been observed for IIb in a narrow temperature range 853 < T/K < 933.However, no unsteady (i.e. oscillating) reaction over the other catalysts at 933-1023 K was observed. The unsteady/steady OCM over La,O, IIa, IIb and IIc catalysts at different reaction conditions is discussed later. Comparison of Catalyst Surface Properties, Activity and Selectivity in Steady OCM The catalysts (1-VI), calcined at 1023 and 1223 K are com- pared for their surface properties and catalytic activity and selectivity in the steady OCM at 1023 K in Table 2. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 3 OCM over La,O, catalyst calcined at 1023 K catalyst reaction temperature/K CH, conversion (%) C, selectivity (%) C2HJC2H6 CO/CO, Ib 116 873 923 973 1023 873 883-923 15.7 21.5 23.5 24.0 23.1 oscillations observed 29.7 47.0 53.6 55.5 49.3 0.37 0.71 0.92 1.03 0.83 0.56 0.39 0.33 0.25 0.29 IIIb 973 1023 873 923 973 24.9 26.5 11.2 14.7 18.9 46.3 47.6 36.5 38.3 46.3 0.89 0.98 0.56 0.56 0.66 0.31, 0.28 0.56 0.54 0.45 IVb 1023 873 923 973 1023 23.7 12.7 20.8 22.0 22.2 54.1 53.6 55.2 54.9 55.2 0.85 0.86 0.97 1.21 1.60 0.39 0.50 0.43 0.37 0.33 Vb 873 923 973 6.9 22.7 22.2 10.0 42.7 48.5 0.63 0.71 <0.1 0.74 0.45 0.40 1023 23.7 51.4 1.oo 0.39 873 17.0 38.9 0.70 0.47 923 19.8 39.8 0.72 0.37 973 22.3 46.2 0.73 0.36 1023 25.3 53.0 0.98 0.29 ~~ Reaction conditions: amount of catalyst = 0.1 g, CHJO, ratio = 4.0 and space velocity (at STP) = 103 200 an3g-'h-'.The very low surface area and catalytic activity and selec- tivity of IVa (Table 2) may be due to sintering of the catalyst, probably related to the presence of traces of sodium in the La carbonate (I) obtained by precipitation with sodium carbon- ate.The increase in the catalyst calcination temperature from 1023 to 1223 K resulted in a large decrease in the surface area and basicity (both total and strong basicity) but, except for IV, a significant increase in both the methane conversion activity and selectivity, and also caused a change in the cata- lyst order for their surface and catalytic properties. The above comparison and the results in Fig. 1-7 and Tables 1-3 reveal a strong influence of the catalyst prep- aration conditions on the surface properties and catalytic activity and selectivity, in steady OCM.From the comparison of the catalysts for their surface acidity and/or basicity with that for the catalytic activity, C, selectivity or C, yield, it can be noted that there is no direct relationship between the surface acidity/basicity and the catalytic activity or selec-tivity, as observed for rare-earth oxides.' Also, in our earlier studies on the OCM over alkali- or rare-earth-metal pro- moted MgO and Ca02593' and alkali- or alkaline-earth-metal promoted rare-earth metal oxides,,' no direct relation- ship between the surface and catalytic properties was observed. The overall OCM process is very complex. It involves a number of catalytic (i.e. surface-catalysed) and non-catalytic (homogeneous or surface-initiated homo-geneous) reactions occurring sim~ltaneously.~~-~~ The con- tribution of the homogeneous reactions to the observed conversion and selectivity in the OCM process is quite appre- ciable and hence it is difficult to obtain a direct relationship between the surface properties and the catalytic activity and selectivity.Unsteady Reaction Behaviour The results indicate that only 11, prepared from La acetate, shows unsteady reaction behaviour in OCM over a narrow temperature range. In the earlier st~dies'-'*~~-~' oscillations were not observed. In our earlier paper,23 the results showing periodic fluctua- tions in reaction temperature and 0, concentration (in the product) indicate symmetric oscillations in OCM over IIb above 823 K but below 973 K.Additional results showing the influence of process parameters and catalyst parameters on the oscillatory behaviour in the OCM process over I1 are presented in Tables 4-6. The minimum and maximum of the 0, concentration oscillation correspond to the maximum and minimum, respectively, of the temperature oscillation. Note that the results of OCM under the unsteady conditions are indicative of changes in the conversion and product selec- tivity measured at close to the minimum and maximum of the oscillating reaction temperatures. It is extremely difficult to obtain results exactly at the minimum and maximum of the oscillating temperature. Efect of Process Parameters The influence of CH4/02 ratio (in the feed) on the steady/ unsteady reaction behaviour for the OCM process over I1 is shown in Table 4.The temperature and 0, concentration oscillations are observed only over a narrow temperature range, depending upon the CH4/02 ratio. For ratios 3.0, 4.0 and 5.0, the oscillations are observed only at 843 < T/ K < 943, 863 < T/K < 933 K and 873 < T/K < 903, respec-tively; for CH,/02 2 6.0, no oscillations were observed. Both the observed temperature and 0,concentration oscillations were symmetrical and sustained for a long period without affecting their amplitude Lie. T,,, -Tmin and C02(max) -COz(minJand cycle period. The recorded oscillations were given el~ewhere.~~.~~ In general, the cycle period and ampli- tude of both the oscillations are found to decrease with increasing reactor temperature.A comparison of the results in Table 5 with that in Table 4 for a CHJO, ratio of 4.0 show a high dependence on the amount of catalyst on the reaction behaviour. When the amount of catalyst in the reactor is changed from 0.1 g (Table 4) to 0.2 g (Table 5), then, at the same space velocity (103 200 cm3g-' h-') oscillations are observed at 883,903 and 923 K only for the smaller amount of catalyst. This indicates that when the linear or superficial gas velocity is doubled, the J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 4 OCM over I1 under steady and unsteady conditions at different CHJO, ratios ~ ~~~ CHJO, reactor temperaturew oscillatingtemperature/K oscillating 0, concentration in product (molyo) methane conversion (%) ~ C, selectivity C,H&,H, ~ CO/CO, IIb 3.0 853 800 f 5 (min) 18.0 f0.5 (max) 16.3 27.2 0.95 0.32 1005 f6 (max) 5.5 f0.4 (min) 28.6 43.0 1.06 0.10 873 838 f3 (min) 15.3 f0.2 (max) 18.9 29.1 0.83 0.72 1003 f 2 (max) 5.6 f0.2 (min) 22.4 42.6 1.01 0.20 923 924 f2 (min) 8.7 f0.2 (max) 22.4 33.1 0.76 0.25 964 f5 (max) 7.2 f0.2 (min) 23.4 37.6 0.86 0.33 973" 973 f5 6.2 f0.1 24.4 40.9 0.86 0.23 1023" 1023 f2 4.7 f0.1 28.6 44.0 1.20 0.15 4.0 883 885 f2 (min) 10.2 f0.1 (max) 23.2 48.8 0.88 0.20 989 f6 (max) 5.2 f 0.3 (min) 24.0 48.3 0.85 0.25 903 913 f2 (min) 8.8 f0.2 (max) 22.5 47.0 0.82 0.23 975 k4 (max) 6.0 f0.2 (min) 22.0 43.8 0.83 0.30 923 933 f3 (min) 7.7 f0.1 (max) 20.8 42.2 0.80 0.35 971 f6 (max) 6.2 f0.1 (min) 22.2 44.9 0.82 0.26 973" 973 f2 6.0 f0.2 24.9 46.3 0.89 0.31 1023" 1023f3 4.9 f 0.1 26.5 47.6 0.98 0.28 5.0 883 868 f2 (min) 13.1 f0.2 (max) 5.9 16.0 0.2 1 0.70 931 f5 (max) 9.7 f 0.3 (min) 7.7 30.0 0.27 0.28 923" 923 f5 9.2 f0.2 10.0 30.4 0.31 0.50 973" 973 f 3 6.5 f 0.1 12.0 30.8 0.35 0.77 1023" 1023 f2 4.9 f0.3 19.1 50.4 0.69 0.26 6.0 883" 883 f2 11.5 f 0.1 3.7 5.7 -0.9 1 923" 923 f 3 8.5 f0.1 8.5 25.8 0.2 1 0.59 973" 973 f 5 6.0 f0.2 14.4 42.8 0.43 0.41 1023" 1023 & 2 4.9 f0.1 16.8 51.3 0.60 0.28 IIC 4.0 873" 873 f 2 2.6 f0.1 23.1 49.3 0.83 0.29 923" 923 f 3 2.2 f0.1 23.7 48.9 0.94 0.29 973" 973 f1 1.9 f 0.3 23.5 46.9 1.02 0.33 1023" 1023 f1 1.8 f0.1 22.9 41.7 1.15 0.39 IIU 4.0 873" 873 f 2 No reaction 923" 889 f 2 (min) 5.2 14.1 0.15 0.50 996 f 3 (max) 18.2 46.7 0.74 0.27 973" 973 f3 23.9 48.0 0.82 0.23 1073" 1073 & 4 24.5 53.3 0.96 0.17 Reaction conditions; amount of catalyst, 0.1 g; feed, CH,-O, mixture of pure CH, and 0, and space velocity 103200 cm3 g-' h-I." Oscillations not observed. Table 5 OCM over IIb under steady and unsteady conditions at different space velocities oscillating 0, run reactor oscillating concentration in methane no. temperaturew tempe!rature/K products (molY0) conversion (%) C, selectivity (%) C,H&,H, CO/CO, ~ space velocity 51 600 cm' g-' h-' 823 823 f 2 no reaction 883" 883 f2 3.5 f0.1 26.4 46.7 0.82 0.22 903" 903f5 3.0 f0.2 28.8 46.9 0.87 0.27 923" 923 f 3 2.8 f0.3 27.4 49.9 0.9 1 0.16 973" 973 f3 2.5 f0.1 28.6 49.0 0.99 0.26 1023" 1023 f 5 2.9 f 0.1 28.2 50.7 1.05 0.17 823 782 f 1 (min) 9.2 f0.2 (max) 18.5 42.5 0.82 0.40 925 f5 (max) 3.0 f0.1 (min) 27.6 41.3 0.89 0.44 8 773 773 f1 20.0 f0.1 no reaction space velocity 103 200 cm3 g-' h-' 823 823 f 2 no reaction 873" 873 f5 7.5 f0.2 19.6 46.1 0.96 0.40 923" 923 f3 4.0 f0.1 24.0 43.1 1.21 0.54 973" 973 f3 3.5 f 0.3 23.7 41.3 1.35 0.55 1023' 1023 f 2 3.0 f 0.2 23.2 40.2 1.52 0.55 833" 833 f3 12.5 f0.3 11.1 47.8 0.78 0.2 1 808 784 f 3 (min) 16.0 f0.2 (max) 5.8 32.8 0.56 0.49 878 f 5 (max) 11.5 f0.1 (min) 13.3 45.3 0.62 0.42 8 773 773 f 1 20.0 & 0.1 no reaction Reaction conditions: amount of catalyst, 0.2 g and CHJO, ratio in feed, 4.0." Oscillations not observed. 3364 J. CHEM. SOC. FARADAY TRANS.,1994, VOL. 90 Table6 OCM over IIb with particle size of 30-72 mesh, under steady and unsteady conditions ~ run no. reactor temperaturew oscillating temperaturew oscillating 0, concentration in products (mol%) methane conversion (%) C, selectivity (%) C,HJC,H, CO/CO, 823 823 & 2 no reaction 873" 873 f2 3.8 f0.2 26.0 44.9 923" 923 f 5 3.5 f0.1 26.7 45.8 973" 973 f3 3.2 f0.2 28.1 47.0 1023' 1023 f 5 3.1 f0.1 28.3 47.7 833" 833 f2 6.8 f0.2 23.8 45.3 773" 773 f 5 9.0 f0.1 20.8 42.5 723 673 715 f8 (min) 765 f6 (ma) 673 f 1 16.7 f0.1 (max) 12.5 f0.2 (min) 20.0 f 1 13.3 15.1 no reaction 40.1 40.7 Reaction conditions: amount of catalyst,0.1 g, CHJO, ratio in fed, 4.0 and space velocity 103 200 cm3g-1.01 0.28 1.08 0.27 1.05 0.27 1.18 0.26 0.93 0.30 0.85 0.36 0.77 0.36 0.78 0.39 h-l." Oscillations not observed. unsteady reaction behaviour changes drastically. It is also interesting to note from the results in Table 5 that, for both space velocities the catalyst do not show any activity initially at or below 823 K but after the reaction at higher tem- peratures, it shows activity with oscillations at 823 K for a space velocity 51 600 cm3 8-l h-' and at 808 K for velocity 103 200 cm3 g-h-'. These observations indicates a possi- bility of formation of active sites responsible for the unsteady behaviour during the reaction at the higher temperatures.The increase in the space velocity resulted in a decrease in the temperature at which the oscillations occurred and also a decrease in the amplitude of both the oscillations. Efect of Catalyst Parameters A comparison of the results on IIa, IIb and IIc for CH,/02 = 4.0 (Table 4) shows that the catalyst calcination conditions (temperature and gas atmosphere, i.e. N, or 0,) have a strong influence on the unsteady reaction behaviour and also on the catalytic properties in steady OCM. In the OCM over IIc, no oscillations were observed and the C, selectivity was found to decrease with increasing reac- tion temperature. With IIa and IIb, oscillations are observed and the C2 selectivity is found to increase with increasing reaction temperature.These observations reveal a strong influence of the gas atmosphere used in the catalyst calcina- tion on the reaction behaviour and selectivity in the OCM process. The results on IIa and IIb (for CHJO, = 4.0) indicate that, when the catalyst calcination temperature is increased from 1023 to 1223 K, no reaction occurs at 873 K and the oscil-lations are observed only at 923 K, thus narrowing drasti- cally the temperature range for the unsteady reaction. A comparison of the results in Tables 4 and 6 shows that for IIb (at CHJO, = 4.0) when the particle size of the cata- lyst is changed from 22-30 mesh to 30-72 mesh, no unsteady reaction behaviour in the OCM was observed at 873-923 K but it was observed at 723 K after the reaction at higher temperatures.The results in Table 6 also reveal that, in this case, no catalytic activity was observed at 823 K but the cata- lyst showed activity at lower temperature after the reaction at higher temperatures, indicating the creation of new catalytic active sites on the catalyst during the OCM reaction at the higher temperatures. The above observations reveal that the unsteady reaction behaviour in the OCM process over I1 is very complex and strongly influenmd by both the process and catalyst param- eters. Aim, new sites, active at lower temperatures, are fornied on the catalyst during the OCM at higher tem- peratures. However, the nature of these sites has not been identified; further work is necessary for this.The observed unsteady reaction behaviour is likely to be due to a large difference in the C2 selectivity of the La203 catalyst in the OCM process at lower (<973 K) and higher (>973 K) temperatures (Table 4); the selectivity is decreased with decreasing temperature. In the pulse reaction of methane over nit-in presence of free 0,(CHJO, = 2.9), the C, selectivity at lower temperatures (823-923 K) was found to be much smaller (almost negligibly small) than at higher temperatures ( 2973 K).,' With IIc, the selectivity decreases with increasing temperature (Table 4) and, therefore, no oscil- lations are observed. In addition to the above the unsteady- state reaction behaviour could also be due to a change in the nature of the active sites on the catalyst.There is a possibility of the formation of surface La oxycarbonate (which is stable at lower temperatures) during the reaction at higher tem- peratures. Since, the unsteady reaction behaviour is observed only in the case of a particular catalyst sample, the tem- perature controller is not expected to have a significant role in the observed oscillations. Conclusions The following important conclusions have been drawn from the present investigation on the surface properties, the cata- lytic activity and selectivity, and the steady or unsteady reac- tion behaviour in OCM of La203 catalysts prepared using different catalyst precursors and calcination conditions : (1) The surface properties (uiz.surface area, crystal size and morphology, acidity and acid-strength distribution, basicity and base-strength distribution) and the catalytic activity and selectivity in OCM are strongly influenced by the catalyst precursor [uiz. hydrated La203, La nitrate, La acetate, La carbonate and La hydroxide]. The calcination temperature also has a strong influence on the surface and catalytic properties of La203 obtained from the different precursors. There is no direct relationship between the surface acidity/ basicity of La203 and its catalytic activity and selectivity in OCM. (2) Among the La203 catalysts, only the one obtained from La acetate and calcined in the presence of N, (but not in 0,) shows unsteady reaction behaviour with symmetrical tem- perature and concentration oscillations in the OCM process at below 933 K in a narrow temperature range.(3) The unsteady reaction behaviour in OCM on the La203 catalyst obtained from La acetate is very complex and strongly influenced by both the process parameters (uiz. tem-perature, CHJO, ratio in feed, space velocity and linear or superficial gas velocity) and catalyst parameters (uiz. particle size and calcination temperature and gas atmosphere used in the catalyst calcination). There is a strong possibility of for- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3365 mation on the catalyst of new sites catalytically active at lower temperatures (below 873 K) during the OCM at higher temperatures.The authors are grateful to Dr. Kuber and Dr. Belhekar (National Chemical Laboratory, Pune) for their help in the catalyst characterization by XRD and SEM. 17 18 19 20 21 Z. Kalenik and E. E. Wolf, Catal. Today, 1992, 13,255. H. Yamashita, Y. Machida and A. Tomita, Appl. Catal., 1991, 79, 203. V. R. Choudhary, S. T. Chaudhari, A. M. Rajput and V. H. Rane, J. Chem. SOC.,Chem. Commun., 1989,555, 1526. V. R. Choudhary, S. T. Chaudhari, A. M. Rajput and V. H. Rane, J. Chem. SOC., Chem. Commun., 1989,605. V. R. Choudhary, S. T.Chaudhari, A. M. Rajput and V. H. Rane, Catal. Lett., 1989,3, 85. References 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 K. Otsuka, K. Jinno and A. Morikawa, Chem. Lett., 1985,499. C. H. Lin, K. D. Campbell, J. X. Wang and J. H. Lunsford, J.Phys. Chem., 1986,90,534. K. Otsuka, K. Jinno and A. Morikawa, J. Catal., 1986,100, 353. K. D. Campbell, H. Zhang and J. H. Lunsford, J. Phys. Chem., 1988,92, 750. S. J. Korf, J. A. Roos, H. M. Diphoorn, R. H. J. Veehof, J. G. van Ommen and J. R.H. Ross, Catal. Today, 1989,4,279. R. P. Taylor and G. L. Schrader, Znd. Eng. Chem. Res., 1991,30, 1016. V. R. Choudhary and V. H. Rane, J. Catal., 1991,130,411. A. Kooh, H. Mimoun and C. J. Cameron, Catal. Today, 1989,4, 333. R. Burch, G. D. Squire and S. C. Tsang, Appl. Catal., 1988, 43, 105. Y. Tong, M. P. Rosynek and J. H. Lunsford, J. Catal., 1990, 126, 291. J. M. Deboy and R. B. Hicks, J. Chem. SOC., Chem. Commun., 1988,982. J. M. Deboy and R. B. Hicks, J. Catal., 1988,113, 517. Y. Feng, J. Niiranen and D.Gutman, J. Phys. Chem., 1991, 95, 6558; 6564. J. M. Deboy and R. B. Hicks, Ind. Eng. Chem. Res., 1988, 27, 1577. E. E.Gulcicek, S. D. Colson and L. D. Pfefferle, J. Phys. Chem., 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 S. Becker and M. Baerns, J. Catal., 1991,128,512. V. R. Choudhary and V. H. Rane, Catal. Lett., 1992,16,211. V. R. Choudhary and V. H. Rane, Catal. Lett., 1990,4, 101. V. H. Rane, Ph.D. Thesis, University of Pune, 1992. S. Bernal, R. Garcia, J. M. Lopez and J. M. Rodriguez-Jzquierdo, Collect. Czech. Chem. Commun., 1983,48,2205. I. Carrizosa, J. A. Odriozola and J. M. Trillo, Inorg. Chim. Acta, 1984,94, 114. J. B. Kimble and J. H. Kolts, CHEMTECH., 1987, 501. V. R. Choudhary and V. H. Rane, J. Catal., 1994,145,300. V. R. Choudhary, A. M. Rajput, D. B. Akolekar and V. A. Selez- nev, Appl. Catal., 1990,62, 171. S. T. Chaudhari, Ph.D. Thesis, University of Bombay, 1993. E. Morales and J. H. Lunsford, J. Catal., 1989, 118,255. R. A. Geisbrecht and T. E. Daubart, Znd. Eng. Chem. Process Des. Deu., 1975, 14, 159. T. Ito, J. X.Wang, C. H. Lin and J. H. Lunsford, J. Am. Chem. Soc., 1985, 107, 5062. V. R. Choudhary, S. T. Chaudhari and A. M. Rajput, AIChE J., 1991,37, 915. V. R. Choudhary and V. H. Rane, J. Catal., 1992,135,310. S. Lacombe, J. G. Sanchez, M. P. Delichere, H. Mozzanego,J. M. Tatiouet and C. Mirodatos, Catal. Today, 1992,13,273. C. Louis, T. L. Chang, M. Kermarec, T. L. Van, J. M. Tatiouet and M. Che, Catal. Today., 1992,13,283. T. L. Van, C. Louis, M. Kermarec and J. M. Tatiouet, Catal. Today, 1992,13,321. 16 1990,94,7069. M. Xu and J. H. Lunsford, Catal. Lett., 1991, 11,295. Paper 4/01718H ;Received 22nd March, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003357

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3367-3372 Comparative IR-Spectroscopic Study of Low-temperature H, and CO Adsorption on Na Zeolites Silvia Bordiga, Edoardo Garrone, Carlo Lamberti? and Adriano Zecchina" Dipartimento di Chimica Fisica, Chimica lnorganica e Chimica dei Materiali, via Pietro Giuria 7, 10125 Turin, Italy Carlos Otero Arean Departamento de Qdmica, Universidad de /as lslas Baleares, 07071 Palma de Mallorca, Spain Vladimir B. Kazansky and Leonid M. Kustov N. D. Zelinsky Institute of Organic Chemistry, Russian Academy of Sciences, Moskow, Russia Extraframework cation sites in the sodium forms of the zeolites ZSM-5, mordenite, Linde-4A and faujasite-type X and Y have been investigated by using low-temperature adsorption of dihydrogen and carbon monoxide as IR spectroscopic probes.The extent of H-H and C-0 bond polarization was found to be dependent not only on the cation electrostatic field, but also on the neighbouring oxygen atoms of the zeolite framework. The influence of these oxygen atoms is most keenly felt by adsorbed molecular hydrogen, but they also affect the IR frequency shift of the stretching vibration of adsorbed carbon monoxide. The Si : Al ratio of the zeolite framework modu- lates the basic strength of the oxygen atoms, and this was found to be reflected in the IR stretching frequency of both adsorbed molecules, H, and CO. The possible role of extraframework cations as catalytically active sites has been postulated and discussed in the first papers'-4 on catalytic applications of zeolites.According to these ideas, such ions can activate adsorbed molecules due to their high polarization power. The extraframework cations protrude into the void internal space of the zeolite and expose adsorbed guest molecules to electric fieldssp7 in the order of 107-108 V cm-'. These intense fields are thought to be responsible for the activation and reactivity of adsorbed guests. Furthermore, the zeolite framework has a negative charge arising from A10,-units which replace neutral SiO, units. While the total negative charge (per unit cell) depends on the Si :A1 ratio, the local charge is a function of the ion- icity of the lattice. Negatively charged oxygen atoms sur-rounding an extraframework cation provide dual acid-base sites (of the Lewis type) which play an important role498.9 in the proposed mechanisms for many catalytic processes medi- ated by zeolites.The usual approach to an experimental estimation of polarization effects in zeolites is the use of adsorbed mol- ecules which act as spectroscopic probes. IR spectroscopy can then monitor the perturbation of the probe molecules in the adsorbed state. However, a suitable choice of different adsorbed molecules is needed for a better understanding of the combined effect of extraframework cations and their neighbouring (framework) oxide anions. This should ultim- ately lead to a better understanding of the catalytically active sites in zeolites. We present here a combined IR study of low- temperature adsorption of molecular hydrogen and carbon monoxide on several sodium zeolites, which have different Si : A1 ratios.The probe molecules were chosen on the basis of their specific (or distinct) interactions with the adsorption sites. Dihydrogen, being a symmetric molecule, has an H-H bond vibration which is IR-inactive. Therefore, molecular hydrogen does not show any absorption bands in the IR unless polarized, inside the zeolite cavities, by the electric field created by cations and surrounding anions. Thus, molec- ular hydrogen has the advantage of the absence of back- ground absorption from the gas phase. In addition, the H-H t Also at: INFN Sezione di Torino, Italy. stretching frequency exhibits, in the adsorbed (polarized) state, very large bathochromic shifts,"." up to 200 cm-'.Finally, this molecular probe has a diameter', of only 2.3 8, and can penetrate even inside the sodalite cages, which have six-membered-ring windows of cu. 2.6 8, in diameter. As a second molecular probe we used carbon monoxide. This molecule has a small dipole moment (0.1 DS) and a rather high polarizability. Therefore, the stretching frequency of CO adsorbed on coordinatively unsaturated cations (via the carbon atom) exhibits hypsochromic shifts' 3-1 resulting from polarization, which can be used to probe the corre- sponding electrostatic field. Reported shifts' 6-1 for CO adsorbed on metal oxides and halides are in the range of 10-100 cm- '. With transition-metal ions, possible charge- transfer effects can also result in shifts of the C-0 stretching frequency.However, for Na' ions this latter effect is not present. Experimental Commercially available sodium forms of ZSM-5 (Si : A1 = 39, mordenite (Si : A1 = 5.0), Y (Si :A1 = 2.35) and X (Si :A1 = 1.25) faujasites, and Linde-4A (Si : A1 = 1.0) were used. Diffuse reflectance IR spectra of adsorbed hydrogen were recorded for zeolite powders at liquid-nitrogen tem-perature, using a Beckman Acta N-7 spectrophotometer equipped with a home-made diffuse reflectance attachment. Further details were given elsewhere.20*2 ' FTIR transmittance spectra of adsorbed carbon monoxide were also recorded at liquid-nitrogen temperature, using a Bruker 88 spectrometer equipped with a MCT cryodetector and a silica cell with NaCl windows.The background created by the zeolite framework was subtracted from all the spectra recorded. Both techniques allowed in situ activation of the zeolites and IR measurements to be made at different equilibrium pressures of adsorbed gases. For activation, the zeolite powders (diffuse reflectance study) or self-supporting wafers (FTIR transmission spectra) were heated under vacuum at 673 K for 2 h. $ 1 D x 3.33564 x C m. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Results and Discussion Low-temperature Hydrogen Adsorption Diffuse reflectance IR spectra of molecular hydrogen adsorbed at 77 K on sodium forms of the different zeolite samples are depicted in Fig.1. Several bands are evident in the H-H bond vibration region. They originate from H, molecules perturbed by Na ions (and surrounding anions) + in different positions inside the zeolite micropores. The spectrum of Na-ZSM-5 [Fig. l(a)] is the simplest since only one IR absorption band, at 4110 cm-', was observed. This shows that there is only one major site of Na' location in Na-ZSM-5. No crystallographic data are avail- able on cation distribution in this zeolite, but indirect strongly suggest that channel intersections are the preferential location of extraframework cations. The spectrum of Na-mordenite at low hydrogen pressure [Fig. l(c)] is quite similar to that of Na-ZSM-5, with a slightly different position of the absorption maximum; 4108 cm-'.At higher pressure [Fig. l(b)] a shoulder develops at 4125 cm-', indicating the existence of at least two different sites of sodium location. This is consistent with current struc- tural Sodium ions in Na-mordenite are prefer-entially located at the centre of highly distorted eight-membered rings, and H, adsorption on these ions must be responsible for the 4108 cm-'band. The shoulder at 4125 cm-' may then be assigned to more loosely bonded Na' ions distributed among other available sites., 5-27 The IR spectra of H, on Na-A [Fig. l(d) and (e)] show only a major band at 4075 cm- ', which develops a shoulder (4110 cm-') at high H, equilibrium pressure. Compared to Na-mordenite, both of these IR absorptions are shifted to lower frequency and are more separated from each other.This is also consistent with structural data25*28 according to which the preferential location of the Na+ ions is near the 4110 I centre of the six-membered rings forming the sodalite cages. The IR absorption band at 4075 cm-' must be assigned to these cations. The shoulder at 4110 cm-' would thus corre- spond to Na' ions located in the remaining (less populated) positions. The IR reflectance spectra of hydrogen adsorbed on sodium forms of faujasites are more complex. For Na-Y [Fig. l(g) and (h)] they show three bands, the relative intensity of which changes with hydrogen pressure. The distribution of sodium ions in dehydrated Na-Y has been studied by X-ray diffraction It was concluded that Na+ prefer- entially occupies s,, sites inside the large cavities and, to a lesser extent, S,.sites inside sodalite cages and S, sites inside hexagonal prisms. This correlates with the IR spectrum of Fig. l(g) if the 4102 cm-' band (which shows the largest shift from free H,) is ascribed to dihydrogen perturbed by Na+ ions in S,, positions, while the bands with maxima at 4125 and 4150 cm-' are assigned to H, molecules under the influ- ence of Na' ions located at SItand S, sites, respectively. Such an interpretation is further confirmed by the spectra of hydrogen adsorbed on Na-X [Fig. l(f)] where similar bands of H, molecules perturbed by Na' at S,, and SIPsites are observed, while the band corresponding to S, sites is absent.This suggests, in agreement with X-ray diffraction that in dehydrated faujasite-type zeolites site S, is less populated in Na-X than in Na-Y. Further details on the assignment of IR bands for molecular hydrogen adsorbed on these zeolites are given elsewhere.,' Low-temperature Carbon Monoxide Adsorption Transmittance IR spectra of carbon monoxide adsorbed at 77 K on Na-ZMS-5, Na-mordenite, Na-Y, Na-X and Na-A zeo- lites at increasing micropore filling are presented in Fig. 2-6. The most prominent feature of these spectra is the presence of two major absorptions, one of them shifted to higher fre- quency (as compared to free CO)and the other one appear- ing at about 2138 cm-'. This latter band is a~signed~,-~~ to non-specifically adsorbed CO which, at relatively high equi- librium pressure, behaves as a liquid-like phase inside the zeolite cavities or channels.This attribution has been con- firmed by an accurate band fitting analysis performed on H-ZSM-5 at different CO equilibrium pressures. The relative intensity of this band is highest for Na-ZSM-5 and lowest for Na-X and Na-A. This could be due to the fact that Na-X and Na-A have the highest extraframework cation population, leaving less space available for the intake of CO in a hindered rotational state. More quantitative consider- ations cannot be made because the amount of weakly Fig. 1 Diffuse reflectance IR spectra of H, adsorbed at 77 K on 2300 2200 2100 . 2000 Na-ZSM-5 (a),Na-M (b) and (c), Na-A (d) and (e), Na-X (f) and wavenumbers/cm -' Na-Y (9)and (h).P,, = 0.3 for (a), (b), (e), (f) and (9)or 13 kPa for (c) Fig. 2 FTIR spectra in the CO-stretching region of Na-ZSM-5 at and (6).Spectrum (h) was taken after outgassing (77 K, 1 min) the increasing CO equilibrium pressure, from ca. 3 Pa up to 3 kPa. The Na-Y zeolite samples with preadsorbed hydrogen. zeolite blank has been subtracted from each spectrum. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1 .o Q, t 2s: 0.5 9 0 1 2300 2200 2100 ’ 2000 wavenumber/cm-Fig. 3 FTIR spectra in the CO-stretching region of Na mordenite at increasing CO equilibrium pressure. Conditions as in Fig. 2. 2-2300 2200 2100 2000 ’ wavenum ber/cm - Fig. 4 FTIR spectra in the CO-stretching region of Na-Y at increasing CO equilibrium pressure.Conditions as in Fig. 2. adsorbed CO is critically dependent upon small fluctuations of the actual temperature of the zeolite wafer. The high- frequency band must correspond to CO polarized by the Na’ ions of the zeolites. In a recent of low-temperature CO adsorption on alkali-metal-exchanged ZSM-5 zeolites (Si :A1 = 14), the position of this band was found to shift gradually from 2188 cm-’ for Li-ZSM-5 down to 2157 cm-’ for Cs-ZSM-5 showing that this high-frequency band is cation specific. For Na-ZSM-5 the cation-specific band (at 2178 cm-’) is narrow and single (see Fig. 2), which is consistent with the fact that this zeolite has only one type of extraframework cation site, as already found in the study of dihydrogen adsorption.A detailed discussion on the minor absorptions present in Fig. 2 has already been performed else~here~~.~~?~~and is not relevant in this context. For Na-mordenite, Fig. 3, the CO spectra are quite complex as two main bands are observed at 2177 and ca. 2160 cm-’. This is consistent with the structure of mordenite3* where two types of channels are present: the first similar to the linear channel present in ZSM-5; the second perpendicular to the first, obstructed by narrow windows located at ca. 8 A from the intersection with the main channel and forming elongated nanopockets. The 2177 cm- similar to the single 2178 cm-’ band observed on Na-ZSM-S/CO spectra, is so assigned to CO interacting with Na+ cations located in the main channel, while the band at ca.2160 cm-’ is attributed to CO interacting with Na+ ions in the lateral obstructed nanopockets. It is most noticeable that two bands are also observed in the H, adsorption on Na-mordenite I Q, C n 2300 2200 2100 2000 wavenumber/cm-’ Fig. 5 FTIR spectra in the CO-stretching region of Na-X at increasing CO equilibrium pressure. Conditions as in Fig. 2. 2300 2200 2100’ 2000 wavenumber/cm - Fig. 6 FTIR spectra in the CO-stretching region of Na-A at increasing CO equilibrium pressure. Conditions as in Fig. 2. [vide supra Fig. l(a)]. According to structural data,”-,’ CO penetrating even partially into these narrow obstructed cavi- ties is expected to undertake interactions not only with one cation but also with the walls.In presence of such multiple interactions, the relationship between the CO frequency shift and the electric field generated by an isolated cation” becomes too rough an approximation and the role played by the neighbour framework atoms cannot be neglected. On the contrary, the validity of this approximation has been proven in the case of a more diluted zeolite37 (ZSM-5 with Si :A1 = 14). In fact, the study of the interaction at 77 K of CO with a complete series of alkali-metal-exchanged ZSM-5 zeolites indicated that a quasi-linear relation exists between the CO frequency shift and the variable l/(Rx + R,,)’, where R,, represents the CO radiu~”*~~and R, the cation radius4’ (X = Li, Na, K, Rb, Cs).The similarity of the frequency of the first CO band with that of Na, ZSM-5, indicates that this relation holds also for alkali-cation-exchanged mordenites. A more detailed discussion of this problem has been performed el~ewhere.~’Note that the assignment of the 2177 cm- band to Na+...CO interactions involving cations in the main channel, and the 2160 cm-’ band to similar interaction with cations placed at the bottom of lateral pockets, is in agree- ment with, (i) the presence of a similar high-frequency band (at 2178 cm-’) for Na-ZSM-5 where only the main channel sites are available, and (ii) the fact that cations in side pockets are more shielded by negatively charged oxygens and there- fore exert a smaller polarization on adsorbed CO molecules.Only one cation-specific CO band (at 2172 cm-’) was found for Na-Y (Fig. 4) while H, gave two or three bands [Fig. l(g) and (h)]. This is due to the fact that the larger CO molecule, which has a diameter12v3' of ca. 3.5-4.2 A, cannot penetrate inside the sodalite cages and hexagonal prisms, thus only the S,, sites (in the supercage) are probed with CO. For Na-X the situation is more complex. Fig. 5 shows that the cation-specific band has a maximum at 2164 cm-and a shoulder at about 2170 cm-'. We assign the maximum to CO molecules interacting with Na+ ions at S,, sites, while the shoulder is probably due to a similar interaction with Na+ at (less populated) S,,, sites also within the supercage.Site SnIis located slightly over a four-membered ring, therefore the cor- responding cation is more coordinatively unsaturated than that at site S,, (which protrudes from a six-membered ring). This would explain why the Na+ ion at S,, polarizes the CO molecule more strongly than that at S,, thus giving a high- frequency shoulder in the IR absorption band. The foregoing assignment is also consistent with the single cation-specific band observed for CO or Na-Y, since in Y-type zeolites S, sites are usually vacant. The fact that dihydrogen shows no resolution of the band assigned to Na+ ions in the supercage [Fig. l(f)] seems to suggest that H, is a less sensitive spec- troscopic probe. A detailed discussion on and relative assign- ments of the minor absorptions present in the Na-Y and Na-X spectra shown in Fig.4 and 5 has been presented else- where.34 For Na-A only a main band at 2163 cm-' is observed, which corresponds to the most intense peak at 4075 cm-'of dihydrogen, shown in Fig. l(d) and (e). The observed peak corresponds to CO on Na+ ions at the centre of six-member rings of the sodalite cages. Owing to the high cationic popu- lation (compared to ZSM-5 and mordenite) and to the small dimensions of the cavities (compared to faujasites), it is quite probable that CO adsorbed on the Naf sites is suffering from simultaneous perturbing effects from other neighbouring cations. Also, in this case a simple relation between the CO frequency shift and the electric field generated at a single cationic position could be an exceedingly rough approx-imation.In this context, we note that due to the stronger interaction of CO with Na+ ions (as compared to H,) the bands of spe- cifically adsorbed molecules are more evident even at lower dosages. Moreover, the band maxima shift slightly toward lower frequencies when going from low to higher CO cover-age. This is a consequence of a gradual decrease of the elec- trostatic charge surrounding the cations when more dipolar CO molecules are admitted into the internal space of the zeloite~.~~ Comparison of Adsorbed Dihydrogen and Carbon Monoxide Table 1 summarizes the observed absorption frequencies of specifically adsorbed dihydrogen and carbon monoxide on the different Na-zeolites.To facilitate discussion, the corre- sponding frequency shifts from the free molecules are also reported. In the case of molecular hydrogen, when more than one band was observed in the spectra the shifts refer to the J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 lowest frequency line (highest bathochromic shift). For CO, when more than one cation-sensitive band was observed, the quoted shifts correspond to the highest frequency (highest hypsochromic shift). The data reported in Table 1 demonstrate a clear depen- dence of the observed frequency shifts upon the aluminium content of the zeolite framework. This applies to both molec- ular probes. However, the variation of the shifts for adsorbed H, follows a trend which is opposite to that observed for adsorbed CO.The bathochromic shift of the H-H stretching vibration is largest for the most basic Na-A zeolite and smal- lest for Na-ZSM-5, while for the hypsochromic shifts of the C-0 stretching vibration Na-ZSM-5 shows the largest effect and Na-A the smallest one. This is a consequence of different mechanisms of perturbation of the adsorbed molecules, which can be explained from the results of quantum chemical analysis of the interaction of CO and H, with A10 clusters mimicking A13 + ions at the surface of aluminium oxide. Semi-empirical calculations 14*42 concerning CO adsorp- tion on A13+ ions in trigonal or square pyramidal coordi- nation provide support for a linear interaction through the carbon end as shown in Scheme 1.No tendency was found for interactions involving the oxygen end of the CO molecule. Maps of the molecular elec- trostatic potential (MEP) of CO, calculated including corre- lation effects,I4 show that the oxygen atom has a second negative zone, albeit less pronounced than that at the carbon end of the molecule. Therefore, CO could, at the most, coor- dinate to two Lewis acid sites through both ends of the mol- ecule. No tendency to interact with Lewis basic sites is expected, because the positive lobes in the MEP are not sub- stantial, particularly in the neighbourhood of the oxygen atom.l4 In conclusion, carbon monoxide probes exclusively the Lewis acid sites which polarize the molecule (through the carbon end) increasing the C-0 For dihydrogen adsorption the situation is quite different. Ab initio SCF MO calculations are available43 for an adsorp- tion site modelled by the Al(OH), cluster with standard bond angles and bond lengths, as shown in Fig.7. These calcu- lations were performed using the GAUSSIAN 80 program and 3-21-G basis set, and optimizing the adsorption complex geometry with respect to the H-H distance and the position of H, relative to the adsorption site. The resulting structure of the molecular complex (Fig. 7) shows that dihydrogen acts as a probe for the surfaee acid-base pairs, where the contri- bution of both basic oxygen and low coordinated cations are equally important. The adsorbed molecule becomes slightly Scheme 1 Table 1 Stretching frequencies of adsorbed H, and CO, and corresponding frequency shifts from the free molecules ~~ ~~~ H-H stretching C-0 stretching frequency of Avu -: downward frequency of adsorbed Av,: upward shift zeolite Si : A1 ratio adsorbed H,/cm-' shift from H, gas4 co/cm - from CO gas* Na-A Na-X Na-Y Na-MOR Na-ZSM-5 1.o 1.25 2.35 5.0 35 4075, 41 10 (sh) 4096, 4125 (sh) 4102,4125,4150 4108,4125 (sh) 4110 -88 -67 -61 -55 -53 2163 2164 2172 2177 2178 +20+21+29+34+35 v~~ =-4163 m-'.vc0 = 2143 m-'. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 H I H Fig. 7 Model for H, adsorption polarized, which renders it IR active and the H-H stretching frequency is lowered. The calculated adsorption heat43 was 29 kJ mol-' which agrees with the fact that molecular hydro- gen is reversibly adsorbed, even at liquid-nitrogen tem-perature.The results shown in Table 1 can now be easily explained. Bathochromic shifts of the H-H stretching frequency should depend both on the polarizing power of the extraframework cation and on the basic strength of neighbouring oxygen atoms. The latter is the highest for Na-A, which results in the strongest perturbation of adsorbed dihydrogen. The weakest perturbation corresponds to Na-ZSM-5 where the basic strength of the oxygen atoms is lowest because the increased Si : A1 ratio increases the covalent character of the zeolite framework. We note that the different frequency shifts observed for adsorbed H, (Table 1) cannot be explained on the basis of different geometries of the corresponding adsorp- tion sites, since S,, sites in both Na-X and Na-Y possess the same geometry.For adsorbed carbon monoxide at an isolated cation site the perturbation depends only on the net positive charge at the extraframework site. However, this is influenced by the surrounding oxygens, which provide a compensating negative charge. Neighbouring basic oxygens in the more ionic zeolite frameworks, which correspond to lower Si : A1 ratio^,^^.^^ decrease the positive electric field in the proximity of the cation sites. This is why the smallest shift of the C-0 stretching frequency is observed for Na-A, while the largest effect corresponds to Na-ZSM-5.The C-0 stretching frequencies of adsorbed carbon mon- oxide (Table 1) can be compared with those of CO adsorbed at low temperature on alkali-metal halides. For CO on NaCl films, several author^'^.^^ have reported an IR absorption band at 2159 cm-' and assigned it to CO interacting with coordinately unsaturated Na+ ions. Hauge et ~l.,~'who studied the interaction of CO with some alkali-metal fluo- rides in an argon matrix, found a C-0 stretching frequency of 2172 an-' for the CO/NaF system. It thus appears that Na' ions in Na-X have about the same polarizing power as in halide films, while in Na-ZSM-5 this polarizing ability is greater, and more similar to that found in single alkali-metal fluoride molecules. This increased polarizing power for adsorbed CO must be correlated with the lower basicity of framework oxygen atoms in Na-ZSM-5, as compared to Na-A, a fact also revealed by the experiments on the adsorp- tion of molecular hydrogen. In summary, it was shown that by using H, and CO as IR spectroscopic probes, not only the different extraframework cation sites, but also their environments can be examined with confidence.The ionicity of the zeolite framework increases with decreasing Si : A1 ratio, and this affects the IR stretching frequencies of the adsorbed molecules. Note, however, that this effect tends to level out when the Si : A1 ratio is greater than five. This is reflected (Table 1) in the very small increase (2 cm-')of the adsorbed H, stretching fre- quency when going from Na-MOR (Si : A1 = 5) to Na-ZSM- 5 (Si :A1 = 35).Furthermore, in a recent IR study3, of CO 337 1 adsorption on two Na-ZSM-5 samples with Si : A1 = 14 and Si : A1 = 35, respectively, a value of 2178 cm-was found in both cases for the C-0 stretching frequency. Support from the CNR (Progetto Finalizzato Chimica Fine 11) and from MURST (40%)is gratefully acknowledged. References 1 J. A. Rabo, C. L. Angell, P. H. Kasai and V. Shoemaker, Discuss. Faraday SOC., 1966,41, 328. 2 R. M. Barrer and R. M Gibbons, Trans. Faraday SOC., 1965,61, 948. 3 H. W. Haynes, Catal. Rev.-Sci. Eng., 1978, 17, 273. 4 J. A. Rabo, Catal. Reu.-Sci. Eng., 1981, 23, 293, and references therein. 5 E. Dempsey, J. Phys. Chem., 1969,73,3660. 6 E.Preuss, G. Linden and M. Peuckert, J. Phys. Chem., 1985, 89, 2955. 7 T. Yamazaki, I. Watanuki, S. Ozawa and Y. Ogino, Langmuir, 1988,4,433. 8 W. W. Kaeding and S. A. Butter, J. Cutal., 1980,61, 155. 9 T. Mole, J. Catal., 1983,94,423. 10 V. B. Kazansky, V. Yu. Borovkov and L. M. Kustov, in Pro-ceedings of the 8th International Congress on Catalysis, Berlin, Dechema Verlag Chemie, Weinheim, 1984, vol. 3, p. 3. 11 V. B. Kazansky, V. Yu. Borovkov and A. V. Zaitsev, in Pro-ceedings of the 9th International Congress on Catalysis, Calgary, ed. M. J. Phillips and M. Ternan, The Chemical Institute of Canada, Ottawa, 1988, vol. 3, p. 1426. 12 Handbook of Chemistry and Physics, ed. R. C. Weast, H. J. Astle and W. H. Beyer, CRC Press, Boca Raton, FL, 66 edn., 1985.13 N. S. Hush and M. L. Williams, J. Mol. Spectrosc., 1974,50, 349. 14 P. Ugliengo, V. P. Saunders and E. Garrone, J. Phys. Chem., 1989,93,5210. 15 G. Pacchioni, G. Cogliandro and P. S. Bagus, Znt. J. Quantum Chem., 1992,42, 11 15. 16 E. Escalona Platero, D. Scarano, G. Spoto and A. Zecchina, Faraday Discuss. Chem. SOC.,1985,80, 183. 17 D. Scarano and A. Zecchina, J. Chem. SOC., Faraday Trans. I, 1986,82,3611. 18 A. Zecchina, E. Escalona Platero and C. Otero Arean, J. Catal., 1987,107,244. 19 D. Scarano, G. Spoto, S. Bordiga, S. Coluccia and A. Zecchina, J. Chem. SOC., Faraday Trans., 1992,88,291. 20 L. M. Kustov, A. A. Alexeev, V. Yu. Borovkov and V. B. Kazansky, Bull. Acad. Sci. USSR, 1981,261,1374. 21 L.M. Kustov and V. B. Kazansky, J. Chem. SOC., Faraday Trans., 1991, 87, 2675. 22 J. R. Anderson, T. Mole and V. Christov, J. Catal., 1980,61,477. 23 N. Y. Topsoe, K. Pedersen and E. G. Derouane, J. Catal., 1981, 70,41. 24 J. Caro, M. Bulow, J. Karger and H. Pfeifer, J. Catal., 1988, 114, 186. 25 W. J. Mortier, Compilation of Extraframework Sites in Zeolites, Butterworth, Guildford, 1982. 26 W. M. Meier, 2. Kristollogr., 1961, 115,439. 27 J. L. Schlenker, J. J. Plut and J. V. Smith, Muter. Res. Bull., 1979, 14,961. 28 R. Y. Yanagida, A. A. Amaro and K. Seff, J. Phys. Chem., 1973, 77, 805. 29 G. L. Eulenberger, G. P. Shoemaker and G. G. Keil, J. Phys. Chem., 1967,71, 1812. 30 K. Hseu, Ph.D. Thesis, University of Washington, 1972. 31 J.V. Smith, in Zeolite Chemistry and Catalysis, ed. J. A. Rabo, ACS Monograph 171, Washington, 1976. 32 S. Bordiga, E. Escalona Platero, C. Otero Arean, C. Lamberti and A. Zecchina, J. Catal., 1992, 137, 179. 33 A. Zecchina, S. Bordiga, G. Spoto, D. Scarano, G. Petrini, G. Leofanti, M. Padovan and C. Otero Arean, J. Chem. SOC., Faraday Trans., 1992,88,2959. 34 S. Bordiga, D. Scarano, G. Spoto, A. Zecchina, C. Lamberti and C. Otero Arean, Vib. Spectrosc., 1993, 5, 69. 35 C. Lamberti, S. Bordiga, G. Cerrato, C. Morterra, D. Scarano, G. Spoto and A. Zecchina, Comput. Phys. Commun., 1993, 74, 119. 36 C. Lamberti, C. Morterra, S. Bordiga, G. Cerrato and D. Scarano, Vib. Spectrosc., 1993,4, 273. 3372 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 37 A. Zecchina, S. Bordiga, C. Lamberti, G. Spoto, L. Carnelli and 43 I. N. Senchenya and V. B. Kazansky, Kinet. Katal., 1988, 29, C. Otero Arean, J. Phys. Chem., in the press. 1131. 38 W. M. Meier and D. H. Olson, Atlas of Zeolite Structure Types, 44 D. Barthomeuf, Mater. Chem. Phys., 1987,17,49. Butterworth-Heinemann, London, 3rd edn., 1992. 45 A. G. Palmenshchikov, E. A. Paukshtis, V. G. Stepanov, V. I. 39 G. E. Leoni, G. E. Ewing and G. C. Pimentel, J. Chem. Phys., Paylov, E. N. Yurchenko, K. G. Ione, G. M. Zhidomirov and S. 1964,40,2298. Beran, J. Phys. Chem., 1989,93,6725. 40 R. D. Shannon, Acta Crystallogr., Sect. A, 1976,32, 751. 46 A. Zecchina, D. Scarano and E. Garrone, Surf. Sci., 1985, 160, 41 S. Bordiga, C. Lamberti, F. Geobaldo, A. Zecchina, G. Turnes 492. Palomino and C. Otero Arean, Langmuir, submitted for pub- 47 R. H. Hauge, S. E. Garnsden and J. L. Margave, J. Chem. SOC., lication. Dalton Trans., 1979, 745. 42 I. N. Senchenya, N. D. Chuvilkin and V. B. Kazansky, Kinet. Katal., 1986, 27,608. Paper 4/02905D; Received 16th May 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003367

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3373-3375 Techniques and Mechanisms in Electrochemistry. By P. A. Christensen and A. Hamnett. Blackie, Glasgow, 1994.Pp. x + 379.Price f24.94.ISBN 0-7514-0129-3 There has been a flurry of textbooks on electrochemistry in the last two or three years, all emphasizing electrode pro- cesses, but none so clearly focussed on modern developments as this one. The approach is novel and well planned. There are three chapters : 1. Introduction to modern electrochem- istry ; 2. Techniques giving mechanistic information; and 3. Examples of the application of electrochemical methods. Chapter 1 provides a quick scamper through interfacial structure and electrode kinetics in 35 pages. It seems to me doubtful whether an advanced undergraduate would be able to follow this without preparation from a more explanatory text, especially as there are some unhelpful misprints, for example, in the section on reactions with adsorbed inter- mediates (pp.29-33). Chapter 2 is centrally placed and is clearly central to the philosophy of the book. It is a very clearly written intro- duction to many of the techniques now being used by electro- chemists. There are several specialist review series which cover this material, but this will provide a valuable intro- duction to them. It is a little surprising to find the chapter starting with the capillary electrometer which is probably the oldest electochemical apparatus still in use. Also, the descrip- tion of the study of solid surfaces by the piezoelectric method may prove to yield important results, but so far the direct measurement of capacity, which is scarcely mentioned, has provided much more detailed and quantitative information.The first half of the chapter continues with important in situ probes which are surface specific :voltammetry, scanning tun- nelling miroscopy and spectroscopy, atomic force micros- copy, IR and Raman spectroscopies, ellipsometry, X-ray methods and impedance. The second half of Chapter 2 deals with in situ probes of the near-electrode region, and begins with the classical elec- trochemical techniques involving consideration of diffusion, followed by rotating disc and ring, EPR, UV-VIS spectro-scopies, the quartz microbalance, FTIR and electrochemical mass spectrometry.There is a final, short section on ex situ techniques dealing with emersion into UHV and mass spec- trometry. This chapter as a whole is an outstanding summary of the ways in which a modern electrochemist can select techniques which will help to solve a particular problem. The principle of each method is simply and lucidly explained. There is clear information about the practice of each method and its strengths and weaknesses as well as real examples of its use. Chapter 3 provides a number of case histories. The first two, and to some extent the third, are concerned with Pt elec- trodes. The literature on such electrodes is now so enormous that a comprehensive discussion of H and 0 adsorption and methanol oxidation could easily fill a volume this size, espe- cially if one starts from Bowden’s work in 1928, as here.The selection made is reasonable with its emphasis on voltam- metry, ellipsometry and IR studies of polycrystalline Pt. Although some elegant work on single crystal surfaces is mentioned, full advantage of the work of the previous decade is not taken, for example, in explaining the nature of strongly and weakly bound H on the polycrystalline electrodes. This is surprising when the book starts with a fairly detailed account of surface structure. The third section on CO, reduction is similarly detailed, but the complexities of the passive film on Fe are described rather briefly in the 4th section; will a student know about mixed potentials and the significance of Fig.3.66(b)? The penultimate section on polymer (largely poly-pyrrole) films shows a skillful selection of the vast material now accumulated and the final section on enzymes in films points to the furture, some of which is already with us. No book is perfect; there are some things I would have liked to have changed: the perpetuation of the unhelpful n, introduced by Nicholson and Shain (p. 174), the statement that adsorbed H is always in equilibrium with H, (p. 233), the use of diagrams taken from published work without rational- ising the plotting of the potential scale. But despite these, and some other trivial criticisms, I have no hesitation in recom- mending this well produced book with enthusiasm.One of its strongest features is the continual use of real experimental data which keeps the reader in contact with the way electro- chemists carry out fundamental research at the present time. R. Parsons Received 22nd March, 1994 Physical Methods of Chemistry, 2nd Edition, Volume 1XA. Investigations of Surfaces and Interfaces. Ed. B. W. Rossiter and R. C. Baetzold. John Wiley and Sons Ltd, Chichester, 1993.Pp. xii + 516. €98.00. Before a surface can be fully characterised, several questions need to be answered. These include: How much surface is there, measured by which probe? What is the geometry of that available surface? What is its chemical composition? How do these things affect the energetics of the surface? This volume, which arrives towards the end of a well known series, is one of a pair planned to cover the field.There are difficulties, therefore, in assessing this single volume. It does not pretend to answer all of the above ques- tions, since there is hardly a mention of surface area measure- ment, porosity or everyday materials. That the book covers techniques of study and measurement, with no experimental results on materials, as such, takes a little getting used to. The first chapter covers the determination of surface tension of liquids, and the second the study of adsorption from solution: these two chapters are slightly out of tune with the rest of the book since they are ‘older’ techniques. They are brought up to date well by the discussion of more recent methods of measurement. Most of the volume is about assessment of well defined solid surfaces, with many metallic and semiconductor exam- ples.A good chapter on scanning tunneling microscopy covers the widening range of techniques for the atomic resolution of surfaces across a wide range of materials. The bulk of the book is taken up with techniques to study chem- istry at surfaces. These include surface spectroscopic tech- niques such as EELS and SIMS, the use of molecular beams, and laser-induced thermal desorption. These last two discuss reactions of molecules sorbed on metal surfaces rather than the surfaces themselves. The individual chapters are all well worth reading. The starting level of each of them is sensible, and it is useful to be reminded of the limits of each technique, but to be encour- aged to push some techniques further.Many of the chapters have references up to the late eighties and occasionally the early nineties. This is probably reasonable in view of the orientation of the series towards techniques. The editors’ aim ‘to produce an introduction to the field.. . to give a clear understanding ... of the value, potential and limitations of the respective techniques’ has been pretty well met even though this single volume does not give a rounded picture of even most of the techniques available: final judge- ment must wait the arrival of the second volume. M. C. Ball Received 27th April, 1994 Cluster Ions. Wiley Series in Ion Chemistry and Physics.Ed. Cheuk-Yiu Ng, Tomas Baer and Ivan Powis. John Wiley 8 Sons Ltd, Chichester, 1993.Pp. xiv + 479.Price f80.00.ISBN 0-471-93830-0. This book contains an interesting series of articles on both experimental and theoretical approaches to the study of gas-phase clusters. The topics chosen cover both weakly and strongly bound clusters, where ‘clusters’ range from weakly bound triatomic systems to very large assemblies. The introductory review by Kamke gives a detailed treat- ment of the photoelectron-photoion coincidence method for cluster studies. This included an extensive assessment of the experimental problems associated with such methods, with a detailed look at one of them, threshold photoelectron pho- toion coincidence, TPEPICO.In the second half of the review examples of weakly bound clusters are given. In the third chapter, by contrast, Jarrold deals with the important strongly bound class of molecules, the silicon cluster ions. He gives a comprehesive treatment of the cluster dissociation paths, of cluster structure and of reactivity. Both the articles by Lisy (Chapter 4) and the longer one by Farrar (Chapter 5) deal with the spectroscopy of solvated ions. These are studied either by the IR techniques described by Lisy or via their electronic transitions as described by Farrar. Taken together these articles give a detailed and up to date overview of this rapidly developing field. In Chapter 7 Brunetti and Vecchiocattivi, give a detailed description of both the observation of the autoionisation of collisional complexes and also of the theory used to interpret them.The complexes treated here are small ones consisting of atom-atom or atom-molecule systems, and are much smaller than those described in the rest of the reviews. The theoretical sections commence with Chapter 2 in which Lifshitz gives a detailed review of the unimolecular and collision-induced decomposition of both proton-bound and carbon-ion clusters. This is followed in Chapter 6 by a detailed review by Last and George of rare gas clusters con- taining charge atoms. In summary this series of reviews contains a wealth of information concerning the present status of cluster ions from both an experimental and a theoretical point of view.It will be very useful to anyone who wants an overview of the current status of this important field of ionised clusters. G. Duxbury Received 27th April, 1994 Applied Laser Spectroscopy. Techniques, Instrumen- tation, and Applications. Ed. David L. Andrews. VCH, Weinheim, 1992.Pp. ix + 471. Price DM 198, €81.00. ISBN 3-527-28072-3. My initial feeling before I opened this book was one of fore- boding. Here was a book with a title similar to many others that had gone before and are now long forgotten. This depression was reinforced by the publisher’s ‘blurb’ on the back cover stating that this book ‘provides a solid introduction’ to laser spectroscopy which requires no back- ground in lasers, only a general understanding of spectroscopy’.How many times does one read similar pro- mises for a book only to be disappointed? However, my initial reaction was misplaced and my fears groundless. This is a truly excellent book that lives up to its claims and even exceeds them. This is the book to give to graduate students who are about to begin a research career involving any aspect of laser spectroscopy. (In part, it has taken me so long to write this review because I have continually had to reclaim my review copy from my research students!) It would form J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 an excellent textbook for postgraduate and post-experience courses in laser spectroscopy. Established scientists who wish to inform themselves about new developments in laser tech- nology and new techniques in laser spectroscopy will find this book invaluable.The value of the book comes from its range of coverage and the high standing of its contributors who are expert practitioners in the fields which they describe. The strength of the book undoubtedly comes from the brief given to the con- tributors by the editor that they should cover the potential and problems associated with their areas of expertise; a ‘warts-and-all’ approach that is largely ignored by most review articles, but which is most important background information for newcomers. The principles of each applica- tion are described together with the specialised instrumen- tation and illustrated with typical applications drawn from chemistry, physics and biology in both academic and com- mercial situations.The first two chapters describe the fundamental applica- tions of laser spectroscopy (Andrews) and review general aspects of laser instrumentation (McCoustra). Techniques in electronic absorption spectroscopy, multiphoton absorption spectroscopy and laser-induced fluorescence are described by Demtroder, Goodman and Philis and by Pfab, respectively. Laser methods in high-resolution IR spectroscopy are detailed by Howard and Brown. Raman spectroscopy forms the basis for two chapters; one detailing modern techniques (Morris) and the other introducing developments in non-linear Raman spectroscopy (Berger, Lavorel and Millot). Laser mass spectroscopy and its applications are reviewed by Ledingham and Singhal and the relatively new and rapidly expanding area of ultrafast spectroscopy is reviewed by Anfinrud, Johnson, Sension and Hochstrasser.The book con- cludes with an invaluable appendix detailing a wide range of laser acronyms. The editor is to be congratulated for the range of topics that have been selected giving the book a wide appeal and for persuading leading experts to write with authority. This book will remain a useful introduction to the several fields described for many years to come. It deserves a place in every science library and in the offices of research groups involved in any application of laser spectroscopy. It should be required reading for those entering into a research career involving laser spectroscopy.A case could easily be made for a low cost paperback version to be produced to make it more accessible to the people for whom it is intended and who will benefit most from it. J. C. Whitehead Received 19th May, 1994 Molecular Magnetism. By Olivier Kahn. Verlag Chemie, Weinheim, 1993.Pp. xvi + 380.Price €63.00. ISBN 3-527-89566-3. ~ ~ ~ ~~~~ The scientific potential of magnetochemistry is two-fold: on the one hand, the investigation of macroscopic magnetic properties such as magnetisation and magnetic susceptibility can serve to elucidate the microscopic, electronic and chemi- cal structure of materials, on the other hand, theoretical mag- netochemical concepts can guide synthetic attempts to control the macroscopic magnetic properties.For example, understanding the interaction of paramagnetic centres in molecular solids is of particular importance in designing new ferromagnetic materials. The interest in this problem has instigated dedicated efforts in the synthetic chemistry of coor- dination compounds of paramagnetic ions of transition metals, making it an intriguing and vigorously growing area of research. The present monograph by Olivier Kahn, himself a very active and renowned researcher in the field of molecu- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 lar magnetism, provides a systematic treatise of the theoreti- cal principles together with representative and illuminating experimental examples taken from current research on the magnetism of mono- and poly-nuclear transition metal com- plexes including infinite linear chains. In the words of the author, the book is primarily written to serve chemists working in the field of molecular magnetism or intending to do so.Regarding the prerequisite knowledge, the author assumes that the reader possesses a good background in sym- metry, ligand field theory and molecular orbital theory, as well as a familiarity with the operator and matrix notations used in quantum mechanics. The book is organized in 13 chapters and includes an appendix with some useful compilations of frequently required matrix representations of spin and angular momen- tum. After establishing the general relations between mag- netic susceptibility and the magnetic field dependence of molecular energy levels, the treatise starts with the consider- ation of molecules (usually coordination compounds) pos- sessing a unique magnetic centre elaborating, in succession, on cases without and with first order orbital momentum and treating the effects of zero field splitting and spin-orbit coup- ling.Two chapters are then devoted to cases of energetically close configiirations of different spin (low-spin-high-spin transition, intermediate spin, spin-admixed state). The major part of the book deals with situations where several magnetic centres interact. Since the general case of many interacting centres is systematically built on the interaction between close pairs, the situation of the two centres in a dinuclear complex is considered in detail, proceeding from isotropic interactions through the anisotropic and antisymmetric inter- actions in such systems, to the central quantum chemical models representing the present day understanding of the relation of orbital structure and spin exchange energies.One chapter then presents the methods for dealing with the new aspects and concepts arising from trinuclear complexes and compounds of higher nuclearity. This is followed by a detailed consideration of magnetic chain compounds usually built from a regular or alternating chain of magnetic metal ions connected by bridging ligands. A critical presentation of theoretical concepts and examples concerning three-dimensional magnetic ordering and the design of molecular- based magnets forms another chapter, while the final one takes up the important problem of ‘double magnetic exchange’, i.e.the combination of Heisenberg spin exchange and spin-dependent electron delocalization in mixed valence compounds. The book has been prepared with great care, although a few inevitable, usually typographical, errors have escaped the final proof reading. It is written in a very clear and rational style. The theoretical concepts are systematically developed and presented in didactically excellent fashion with efficient support by many instructive figures and diagrams. As a rule, for all the compounds discussed as examples, pertinent struc- tural information is given in graphical form.Theory and experimental examples are well balanced. This book fully achieves the author’s intent to provide a useful guide for those seeking a systematic access to the field of molecular magnetism. I find this an excellent monograph which may be expected to become a standard reference in the field. U. E.Steiner Received 13th June, 1994 Introduction to Modern Vibrational Spectroscopy. By Max Diem. John Wiley and Sons Ltd, New York, 1993. Pp. xiii + 285.Price f49.50.ISBN 0-471-59584-5. There are a number of excellent recent texts which provide an introduction to molecular spectroscopy, as well as the classic more advanced texts by Herzberg and by Wilson, Decius and Cross. However, there are few which deal comprehensively with modern vibrational spectroscopy at an introductory level.Ideally, such books should contain the theoretical back- ground necessary to understand the principles of vibrational spectroscopy, together with a description of the more impor- tant modern experimental spectroscopic techniques and com- prehensively illustrated by selected applications. To accomplish this in a text of reasonable length is a challenging task. The author has largely achieved this aim although the balance of material could perhaps be improved particularly for advanced undergraduate use. Following a short introduction, the second chapter out- lines the elementary quantum mechanical background of rotational and vibrational spectroscopy for simple molecular systems. Although limited in mathematical detail, this chapter covers most of the relevant theory required by the remainder of the text and also provides an introduction to more com- prehensive texts dealing with the quantum mechanical basis of spectroscopy.The vibrations of polyatomic molecules are treated in Chapter 3. Elementary classical and quantum treatments of vibrations in polyatomic molecules are present- ed, followed by an introduction to the theory and techniques used for the computation of vibrational frequencies and normal modes of vibration. The water molecule is used as an example, and frequencies are computed using three types of force fields. IR absorption intensities are also discussed briefly at the end of this chapter.The symmetry of molecular vibrations is introduced in the next chapter. The basic con- cepts of group theory relevant to vibrational spectroscopy are outlined in this section. This basic treatment should be sufi- cient to enable the interpretation of the vibrational spectra of many molecules and the derivation of the appropriate selec- tion rules to be carried out. In the following chapter, the theory of Raman spectroscopy is introduced. This section is well written and also contains brief accounts of resonance, non-linear and surface enhanced Raman spectroscopy. The instrumentation for both IR and Raman spectroscopy is covered in Chapter 6. A discussion of interferiometric tech- niques, including Fourier and Hadamard transforms, is included.Modern Raman instrumentation including multi- channel and Fourier transform spectrometers are described in some detail. The remaining one third of the book is concerned with applications of vibrational spectroscopy beginning with selec- ted small molecules and graduating to larger, more compli- cated, mainly biological, molecules. Considerable space is devoted to an in-depth treatment of the biological applica- tions of vibrational spectroscopy. Clearly, this is an area of major interest to the author and provides a good intro-duction to this still rapidly developing area of spectroscopy. Together with vibrational optical activity, this section occupies the major part of the section dealing with applica- tions. In an introductory text, this represents a considerable bias towards the biological area at the expense of other topics such as gas-phase or high-resolution spectroscopy which might have been included. Although vibrational optical activ- ity is also an important topic which should be included, the space devoted to this technique in a general introductory text is probably excessive. This book is accurately and well presented, and in spite of the above reservations, can be recommended as an up to date introduction to modern vibrational spectroscopy both for advanced undergraduate and postgraduate readers. The price of E49.50 for the hardcover edition is, however, too high for undergraduate purchase as a recommended text. A cheaper paperback edition would be welcome and more realistic for this sector of the market. It would also encourage graduate students to purchase the book for personal use. R. T. Bailey Received 16th June, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003373

出版商:RSC    

年代:1994

数据来源: RSC