摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 3605-3611 Synthetic Hydroxyapatites as Inorganic Cation Exchangers Part 2 B Y TAKASHI SUZUKI,* TOSHIAKI HATSUSHIKA AND MICHIHIRO MIYAKE Department of Applied Chemistry, Yamanashi University, Takeda-4, Kofu-shi 400, Japan Received 19th March, 1982 Ihe removdl of c d k i u h as Pb2+ Mn2+, Co2+ and Cu2+ in aqueous solution by four synthetic hydroxyapatites (S- 1, S-2, S-3, S-4) has been investigated using both batch and column methods. The removal is due not only to an adsorption effect but also to an ion-exchange reaction between the cations in solution and the Ca2+ ions of the apatites. The order of the ions according to the amount exchanged was as follows: Pb2+ > Cu2+ > Mn2+ 1: Co2+. Pb2+ ions were readily removed by the apatites and the maximum value for the exchange of Pb2+ ions was 230 mg per g of S-4 apatite.The apatites, particularly S-4, would seem to be possible agents for the removal of toxic Pb2+ ions. The selectivity of the apatites for the cations can be explained by considering the radii and the electronegativities of the ions. It is well known that hydroxyapatite [Ca,,(PO,),(OH),] is the major inorganic constituent of biological hard tissues1 such as bones and teeth and consequently its surface characteristics have been investigated from various standpoints. Macromolecules2 such as polypeptides and proteins may be adsorbed, in the usual sense of physical adsorption or chemisorption, on its surface, while some ions may ex- change with the ions of the hydroxyapatite lattice. Some of these exchange reactions, i.e.anion-exchange reactions3 between the hydroxy ions of the lattice and anions in solution, have been extensively studied. However, little attention has been given to cation-exchange r e a c t i o n ~ ~ - ~ between calcium ions of the lattice and cations in solution. In a previous paper,' we noted that some synthetic hydroxyapatites could be used as inorganic cation exchangers for metallic cations. The purpose of the present paper is to examine the cationic characteristics in more detail and to clarify the selectivity of the apatites for various cations. EXPERIMENTAL MATERIALS Four hydroxyapatite samples were prepared as follows. Sample 4 (S-4) was prepared by precipitation from hot water, i.e. by dropwise titration of phosphoric acid solution into a boiling calcium hydroxide solution, according to the method of Brown et a1.8~9 Sample 3 (S-3) was prepared by a similar precipitation method as described by Moreno et 111.l~ Sample 1 (S-1) and TABLE I.-SPECIFIC SURFACE AREAS AND Ca/P MOLAR RATIOS OF FOUR SYNTHETIC HY DROXY APATITES apatite s- 1 s-2 s-3 s-4 ~ ~~~~~ Ca/P molar ratio 1.60 1.56 1.66 1.66 surface area/m2 g-l 30.0 11.4 30.7 36.0 36053606 HYDROXYAPATITE CATION EXCHANGERS sample 2 (S-2) were commercially obtained materials (S-1 from Wako Pharmacy Co and S-2 from Nihon Chemical Co).S-1 and S-2 were synthesized for chromatographic use by a precipitation method. The specific surface areas of the samples were determined by the B.E.T. method using nitrogen adsorption at liquid-nitrogen temperatures, assuming the cross-section area of a nitrogen molecule to be 16.2 A2.The specific surface areas and Ca/P molar ratios of the four samples are summarized in table 1 . All chemicals were analytical reagent grade commercial materials and used without further purification. BATCH METHOD 1 g of each sample was mixed and stirred with 400 cm3 of 100-1000 ppm aqueous solutions containing the cations Pb2+, Mn2+, Co2+ and Cu2+ in the form of nitrates at 20 "C. After 2 h, the supernatant solutions were analysed by atomic absorption spectroscopy and EDTA titration method. Thus, the amount of Ca2+ released into solution and the amount of cation absorbed by the samples were determined. COLUMN METHOD In order to study the removal of the cations by the apatites in detail, the samples (1, 3 and 5 g) were weighed into a glass column (length, 300 mm; diameter, 10 mm) and 400 cm3 of the 100-1000 ppm solutions was passed through the column at a flow rate of 3 cm3 min-' at 20 "C.After washing with 100 cm3 of distilled water, the eluents were analysed. In addition, structural changes in the samples before and after uptake were investigated using powder X-ray diffraction. RESULTS AND DISCUSSION Table 2 summarizes the results for the removal of Pb2+ ions from 100-1000 ppm aqueous solutions (400 cm3) by S-4 (1 g) apatite at 20 OC using the batch method. Note that the molar ratio Pb2+/Ca2+ is almost equal to 1.0. This means that the uptake of Pb2+ is an ion-exchange reaction between Pb2+ ions in solution and Ca2+ ions in the apatite, as in the Cd2+-Ca2+ system discussed previ~usly.~ Moreover, the removal ratio of Pb2+ ions from 100-200 ppm solutions by S-4 is loo%, and the amount of Pb2+ ions removed from the 1000 ppm solution is ca.230 mg, even though the removal ratio is 57%. This is interesting because the maximum removal of Cd2+ ions, found to be one of the most easily removed ions,l' is only 80 mg under the same conditions. In other words, Pb2+ ions are readily removed by S-4, which has potential as an agent for the removal of Pb2+ ions. Table 3 shows the removal ratios of Pb2+ ions from 400 cm3 of 100 ppm solutions by 1 g of various synthetic hydroxyapatites (S-1 to S-4) at 20 "C using the batch method. As shown in the table, the removal ratios are different for the different TABLE 2.-REMOVAL OF Pb2+ IONS FROM 100-1000 ppm SOLUTIONS BY s-4 APATITE AT 20 "c USING THE BATCH METHOD concentration (PPm) 100 200 500 1000 Pb2+/mg per 400 cm3 39.9 77.9 171.9 231.1 ( x lop4 mol) (1.93) (3.76) (8.26) (1 1.2) ( x mol) (1.92) (3.75) (8.16) (11.1) Pb2+ : Ca2+ 1: 1 1 : l 1 : 0.99 1 : 0.99 molar ratio Ca2+/mg per 400 cm3 7.7 15.0 32.7 44.4 removal ratio (%) 100 100 86 57T.S U Z U K I , T. HATSUSHIKA A N D M. MIYAKE 3607 TABLE 3.-REMOVAL OF Pb2+ IONS FROM 100 ppm SOLUTION BY VARIOUS APATITES AT 20 " c USING THE BATCH METHOD sample s- 1 s-2 s-3 s-4 Pb2+/mg per 400 cm3 34.2 17.5 39.9 39.9 ( x mol) (1.65) (0.84) (1.93) (1.93) ( x mol) (1.65) (0.82) (1.92) (1.94) molar ratio of Pb2+ (%) Ca2+/mg per 400 cm3 6.6 3.3 7.7 7.8 Pb2+ : Ca2+ 1 : l 1 : 0.98 1 : l I : 1 removal ratio 89 45 100 100 apatites; e.g.the removal ratio of S-2 is only 45%. S-2 was found to be a mixture of Ca,,(PO,),(OH), and Ca,(PO,),. H20 by powder X-ray diffraction and to have a small surface area of 11.4 m2 g-l (cf. table 1). These seem to be the cause of the small removal ratio. However, granular S-2 seems to possess an advantage over the other apatites in that it can be used for chromatography without applying pressure. Fig. 1 shows the results for the removal of Pb2+ ions from 100-1000 ppm solutions (400 cm3) by 1 g of the four samples using the column method. It is clear that the removal of Pb2+ ions is similar to that observed using the batch method and that the order of the removal ratios for the samples at each original concentration is as follows : s-4 > s-3 > s-1 > s-2.Furthermore, no structural changes were detected by powder X-ray diffraction. Thus these samples, especially S-4, could be used as agents for the removal of Pb2+ ions. Indeed, it was found that only 3 g of S-4 is required to remove 400 mg of Pb2+ ions using the column method, as shown in fig. 2. Table 4 shows the uptake phenomena of Mn2+ ions by S-4 under the same conditions as in table 2. From table 4 it was calculated that the molar ratio '0°1 I 8 Kl 0 X 0 t t X 0 A 0 I X I I I I 0 LO 80 200 LOO original amount of Pb*+/mg per 400 cm3 FIG. 1.-Removal of Pb2+ ions at 20 O C using the column method. 0, S-4; A, S-3; x , S-2; 0, S-I.3608 HYDROXYAPATITE CATION EXCHANGERS i A t 0 200 LOO amount of Pb2+ in original solution/mg per 400 cm3 FIG.2.-Removal ratios of Pb2+ ions on 3 g of 0, S-4 and A, S-3 apatites using the column method. TABLE 4.-REMOVAL OF Mn2+ IONS FROM 100-1000 ppm SOLUTIONS BY s-4 APATITE AT 20 OC USING THE BATCH METHOD original Mn2+ concentration (ppm) 100 200 500 1000 Mn2+/mg per 400 cm3 10.2 13.8 20.2 24.2 ( x mol) (1.86) (2.51) (3.68) (4.40) Ca2+/mg per 400 cm3 7.7 10.1 15.2 17.7 ( x mol) (1.93) (2.53) (3.79) (4.42) Mn2+: Ca2+ 1:1.04 1:1.01 1:1.03 1:1.01 molar ratio removal ratio 25 17 10 6 of Mn2+ (%) Mn2+/Ca2+ is almost equal to 1.0, as in the case of Pb2+/Ca2+, but that the removal ratios of Mn2+ ions are much smaller than those of Pb2+ ions (cf. the fifth line of table 4). The trend of smaller amounts of Mn2+ being exchanged was also found for S-1, S-2 and S-3.Fig. 3 shows the removal ratios of Mn2+ ions by 5 g S-4 and S-3, which have high removal ratios for the cations, from 100-1000 ppm solutions (400 cm3) using the column method. Note that a large amount of S-4 ( 5 g) is needed to remove completely the Mn2+ ions in the solution with a low concentration of 100 ppm (40 mg per 400 cm3). These results suggest that the synthesized hydroxyapatites have the same selectivity for metallic cations as that found previously.'T. SUZUKI, T. HATSUSHIKA A N D M. MIYAKE 3609 E h ' O 0 I $ f 2 0 A 0 A 0 A 0 A t 0' I I I I 40 80 2 00 L 00 original amount of Mnz+/mg per 400 cm3 FIG. 3.-Removal ratios of Mn2+ ions on 5 g of 0, S-4 and A, S-3 apatites using the column method. 0 1 I I I I I 100 200 300 LOO 500 original concentration (ppm) FIG.4.-Removal of Co2+ ions from 100-500 ppm solutions by 0, S-1, A, S-3 and 0, S-4 apatites using the batch method.3610 HYDROXYAPATITE CATION EXCHANGERS Fig. 4 shows the results for the removal of Co2+ ions by 1 g of S-4, S-3 and S-1 from 400 cm3 of 100,200 and 500 ppm solutions using the batch method. It was found that the amount of Co2+ ions removed and the order of the apatites (S-4 > S-3 > S-1) are almost the same as in the case of Mn2+ ions. Fig. 5 presents the removal of Cu2+ ions from 400 cm3 of 100-500 ppm solutions by 1 g of S-4, S-3 and S-1 using the batch method. The amounts of Cu2+ ions removed I 1 I are greater than those of Mn2+ and Co2+ ions, but smaller than those of Pb2+ ions. In other words, the order of the removal ratios of the cations is as follows: Pb2+ > Cu2+ > Mn2+ N Co2+.Fig. 6 shows the relationship between the ionic radius12 and ele~tronegativityl~ of various cations, where MI Ca and M2 Ca represent the ranges of radii of two different Ca2+ ions1* in the hydroxyapatite. The following results are obtained from fig. 6. First, the radii of easily removed ions such as Pb2+ and Cd2+ fall within the range of the radius of M2 Ca (0.9-1.3 A) and they have large electronegativity values. Secondly, Mg2+ and Ba2+ ions, which are hardly removed by the apatites,' have radii which fall outside the range and low electronegativity values. Thirdly, Cu2+, Co2+ and Mn2+ ions show intermediate behaviour, i.e. their radii fall outside the range but they have high electronegativity values.From these results, we conclude that if cations have large electronegativity values and radii close to the range 0.9-1.3 A then they are more easily removed by the hydroxyapatites. This work was supported by a Grant in Aid for Research from the Japanese Government Ministry of Education. We thank Mrs Hideko Hatsushika for her secretarial help during the preparation of the manuscript.T. S U Z U K I , T. HATSUSHIKA A N D M. MIYAKE 361 1 2.0 1 0 CdvD 0 MI Ca(VI) - I I 0.6 0.7 0.8 0.9 1.0 1 . 1 1 . 2 1.3 1.4 ionic radius/A FIG. 6.-Relationship between electronegativity and ionic radius. See, for example, C. L. Kibby and W. K. Hall, The Chemistry of Biosurfaces (Marcel Dekker, New York, 1972), vol. 2, p. 663. V. Hlady and H. F. Milhofer, J.Colloid Interface Sci., 1979, 69, 460. See, for example, I. Zipkin, A. S. Posner and E. D. Eanes, Biochim. Biophys. Acta, 1962, 59, 255; R. A. Young, Proc. 2nd Int. Congr. Phosphorus Compounh (IMPHOS, Paris, 1980), p. 73. C. Y. C. Pak and F. C. Bartter, Biochim. Biophys. Acta, 1967, 141, 410. T. Suzuki, T. Hatsushika and Y. Hayakawa, Proc. Zndlnt. Congr. Phosphorus Compounds (IMPHOS, Paris, 1980), p. 165. T. Suzuki, T. Hatsushika and M. Miyake, Proc. 3rd Congr. Inorganic Phosphorus Chemistry (The Chemical Society of Japan, Tokyo, 1981), p. 16. T. Suzuki, T. Hatsushika and Y. Hayakawa, J. Chem. Soc., Faraday Trans. 1, 1981, 77, 1059. Y. Aunimelech, E. C. Moreno and W. E. Brown, J. Res. Nut1 Bur. Stand., Sect. A , 1973, 77, 149. H. Mcdowell, T. M. Gregory and W. E. Brown, J. Res. Natl Bur. Stand., Sect. A, 1977, 81, 273. lo E. C. Moreno, T. M. Gregory and W. E. Brown, J . Res. Natl Bur. Stand., Sect. A , 1968, 72, 773. l 1 T. Suzuki and Y. Hayakawa, Proc. 1st Int. Congr. Phosphorus Compounds (IMPHOS, Paris, 1977), l 2 R. D. Shannon and C. T. Prewitt, Acta Crystallogr., Sect. B, 1969, 25, 925. l 3 L. Pauling, The Nature of the Chemical Bond (Cornell University Press, Ithaca, New York, 3rd edn, l4 M. I. Kay, R. A. Young and A. S. Posner, Nature (London), 1964, 204, 1050. p. 381. 1960). (PAPER 2/475)

ISSN:0300-9599    

DOI:10.1039/F19827803605

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3613-3617 Comparison of Nernst-Planck Equations and Miller’s LN Approximations for Countercurrent. Electrolysis in a Thin Porous Membrane BY KYOSTI KONTTURI* AND PIRKKO FORSSELL Helsinki University of Technology, Laboratory of Physical Chemistry, SF-02 150 Espoo 15, Finland AND AARNE H. S I P I L A ~ Helsinki University of Technology, Computing Centre, SF-02 1 50 Espoo 15, Finland Received 29th March, 1982 Two models which describe the transport of electrolytes in a porous membrane during countercurrent electrolysis are compared with experimental results. One model is based on Nernst-Planck equations and the other is based on Miller’s LN approximations. The comparison was made in the ternary system NaCl+HCl+H,O at a total concentration of 0.01 mol dm-3.Both models give errors of approximately the same order of magnitude, and they both predict the behaviour of the system fairly accurately. The errors are surprisingly high compared with the errors in the transport quantities. This is thought to be due to the accumulation of small errors in the integration over the membrane. When convection is in the opposite direction to electrical conduction, separation of ions can be achieved., This kind of separation process, called countercurrent electrolysis in a thin porous membrane, is modelled as follows. A system consists of two electrolyte solutions (a and /3) separated by a membrane. This membrane is of a wide-pore type so that there is no interaction between the membrane and electrolytes.We have a flow of solution through the membrane, i.e. convection, which is determined by the pressure difference. The electric current is determined by the electromotive force and the resistance of electric circuit. The values of these quantities are constmt during the transport process. The electrolyte solutions (a and p) are kept homogeneous by stirring. The membrane is so thin that the flows of ions in the transport process can be assumed to be constant. To understand this process a theoretical model must be known, i.e. we have to be able to determine the flows of ions across the membrane when convection, electric current and boundary concentrations on both sides of the membrane are known. Since the membrane does not affect the transport process, it can be regarded as an unstirred layer between two electrolyte solutions.This is in fact means that the transport can be described by the equations used to describe in electrolyte solution. Within the framework of the thermodynamics of irreversible proce~ses,~-* the flows of ions with the same charge in a ternary system rl,CrlaA + rZcDrzaA + H,O are Present address : Digital Equipment Corporation, Finland. I I7 FAR 1 361 33614 COUNTERCURRENT ELECTROLYSIS I N A POROUS MEMBRANE where subscripts 1 and 2 denote the ions with the same charge, e.g. cations, j , and j , are the ionic fluxes, u is the velocity of solvent with reference to the membrane, t, and t , are the transport numbers, Dij. are the diffusion coefficients in Hittorfs reference system, c, and c, are concentrations, z, and z, are charge numbers, and rlc and rZc are stoichiometric numbers.By integrating eqn (1) and (2) over the membrane the ionic fluxes are obtained. These fluxes depend on the electric current, the convection and the boundary concentrations on both sides of the membrane. However, this is not feasible, except in a few cases, because the dependences of diffusion coefficients and transport numbers on concen- trations in multicomponent systems are not k n ~ w n . ~ ~ The lack of measured data for transport quantities means that we have to estimate the concentration dependences of these quantities. A simple and efficient way to estimate diffusion coefficients and transport numbers in multicomponent systems is offered by Nernst-Planck equations. There are also more sophisticated and accurate estimation methods, e.g.those of Miller.' These methods are usually based on the thermodynamics of irreversible process; for a review see ref. (8). The Nernst-Planck equations are based on a kinetic m0de1.~J~ Invoking the condition of electroneutrality and the relationship between the ionic fluxes and electric current the Nernst-Planck equations for a ternary system can be presented, after elimination of the electric-potential gradient, as jl = j 2 = where A,, A2 and A3 are the molar conductivities of the ions. By comparing eqn (3) and (4) with eqn (1) and (2) we can obtain the values of the diffusion coefficient and the transport numbers as given by the Nernst-Planck approximation. The transport quantities so obtained are the same as those given by Wendt," Gosting12 and Millers in dilute solutions.The diffusion coefficients in eqn (3) and (4) are the ternary analogues of the Nernst-Hartley equations. More accurate methods of estimating the diffusion coefficients are the so-called LN approximations of Miller. These approximations are based either on purely binary data or on partly ternary data, and they take into account the cross-effects of ions. These effects are neglected in the simple methods referred to above. Estimating the transport numbers according to McInnes13 and the diffusion coefficients by LN approximations we can solve the transport problem more accurately than by using the Nernst-Planck approximation. However, the modelling of transport phenomena in countercurrent electrolysis with Miller's method is considerably more difficult than modelling with the Nernst-Planck equations.We will thus compare these different approaches with experimental results to discover whether these approximations are worth using in place of the Nernst-Planck equations.TABLE 1.-EXPERIMENTAL AND CALCULATED RESULTS FOR THE SYSTEM NaC1+ HCl +H,O; cga+ = c&+ = 0.5 x mol dm-3 Miller Nernst-Planck measured approximation approximation V” - 4 %a+ Y&+ S l o g s KNa+ K H + KNa+ KH + KNaf KH+ 1.62 3.68 0.075 0.80 11 1.03 -0.114 - 1.21 -0.111 -1.13 -0.083 - 1.18 3.35 5.29 0.0175 0.85 49 1.70 -0.027 -1.29 -0.032 -1.19 0.0004 - 1.27 4.47 7.35 0.0074 0.85 115 2.06 -0.011 -1.29 -0.012 -1.97 0.0085 -2.03 6.02 9.14 0.0017 1.00 590 2.77 -0.0026 -1.52 0.0037 - 1.39 0.0056 - 1.36 7.49 10.58 0.00052 0.93 1800 3.25 -0.0007 -1.40 0.0016 - 1.69 0.0027 - 1.67 Definitions of dimensiqnless parameters, when c, = 0.01 mol drnp3, Do = 2 x 10-5 cm2 s-l, A/1= 10 cm, F is Faraday’s constant and I is the thickness = ci/c, (concentration); S = Y& Ya,/Y& Yga (selectivity of the membrane: vk = V’I/AD, (convection); I,. = Il/c,D,F (electric current density); ratio); KH = j, l / D , c,,, KNa = jNa / / D o C, (ion fluxes).0 Z 43616 COUNTERCURRENT ELECTROLYSIS I N A POROUS MEMBRANE 3 2 % - T 1 0 - Vk FIG. 1.-Plot of the logarithm of the selectivity ratio ( S ) as a function of convection (uk); see table 1. EXPERIMENTAL The experiments were performed using the ternary system NaCl + HCl + H,O. The apparatus used has been described in detail previ~usly.'~ Measurements were carried out by feeding pure water into compartment a with volume flqw Voand by taking solution out of compartment a with volume flow Va.Convection was V k = V o - Va. The concentrations in compartment D were kept constant during the measurements by circulating fresh solution through this compartment. The Ag, AgCl electrodes were used as both cathode and anode. The membrane used in our measurements was Millipore SC with pore size 8 pm, reinforced with a nylon net (thickness ca. 0.3 mm). The electric current through the cell was controlled by potentiostat type 120 from Sycopel Scientific Ltd. The solution flows were regulated by the peristaltic pump Ismatec ip-12 and the flow rates were determined by weighing solution volumes.Hydrogen ion concentration was analysed by titration and sodium ion concentration by flame photometry. RESULTS AND DISCUSSION The results of our measurements are reported in table 1, which also includes the fluxes calculated according to Miller7 and Nernst-Planck. In fig. 1 we have plotted the exponential dependence on the selectivity ratio ( S ) on experimentally verified convection, as this relationship has special interest when considering the countercurrent electrolysis in a thin porous membrane as a separation pr0cess.l The calculations of hydrogen ion and sodium ion fluxes in these different models were made by the shooting method.15 We tried different Miller approximations7 but no significant differences were noticed. The tabulated values were calculated using the LN approximation recommended by Kim et aZ.,16 modified by taking the activitiesK.KONTTURI, P. FORSSELL AND A. H. SIPILA 3617 of the mixtures according to Guggenheim.17 The approximation of Franck and Thomsonls was used for the activity of sodium chloride, required as a reference in Guggenheim’ s method . On studying the results in table 1 we can deduce the following. As would be expected, Miller’s approximation is more accurate than the Nernst-Planck approxi- mation. The actual differences are relatively large, even though the total concentration is only mol dmP3. The steeper the concentration profiles, the greater the errors become. This is due to the fact that small errors in the diffusion coefficients and the transport numbers are accumulated in the integration over the membrane causing large errors in the calculated fluxes.However, both Miller’s approach and the Nernst-Planck flux equations gave results whose magnitudes are of the same order. Taking into account the fact that using the Nernst-Planck equations in the form of eqn (3) and (4) is easier than using Miller’s method we conclude that in most cases the behaviour of the system can be described sufficiently well using the Nernst-Planck equations. We thank the Neste OY Foundation for financial support. K. Kontturi, P. Forssell and A. Ekman, Sep. Sci., submitted for publication. R. Haase, Thermodynamics of Irreversible Processes (Addison-Wesley, London, 1969). T. F~rland and S. Ratkje, Electrochim. Acta, 1980, 25, 157 and references therein. A. Ekman, S. Liukkonen and K. Kontturi, Electrochim. Acta, 1478, 23, 243. E. Cussler, Multicomponent Diflusion (Elsevier, Amsterdam, 1976). H-J. Schonert, 2. Phys. Chem. N.F., 1967, 51, 196. D. Miller, J. Phys. Chem., 1967, 71, 616. D. Miller, Faraday Discuss. Chem. SOC., 1977, 64, 295. W. Nernst, 2. Phys. Chem., 1888, 2, 37. lo M. Planck, Ann. Phys. Chem. N.F., 1890, 39, 161 and 561. R. Wendt, J. Phys. Chem., 1965, 69, 1227. l2 L. Gosting, Adv. Protein Chem., 1956, 11, 536. l 3 D. McInnes, J. Am. Chem. Soc., 1933, 55, 996. l4 A. Ekman, P. Forssell, K. Kontturi and G. Sundholm, J. Membr. Sci., in press. l5 A. Sipila, A. Ekman, and K. Kontturi, Finn. Chem. Lett., 1979, 97. l 6 H. Kim, G. Reinfelds and L. Gosting, J. Phys. Chem., 1973, 77, 934. l 7 E. Guggenheim, Applications of Statistical Mechanics (Clarendon Press, Oxford, 1966), p. 165. lA H. Franck and P. Thompson, in The Structure of Electrolytic Solutions, ed. W. Hamer (John Wiley, New York, 1959), p. 113. (PAPER 2/536)

ISSN:0300-9599    

DOI:10.1039/F19827803613

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 3619-3627 Effect of Solvent on the Kinetics of the Reaction between Dialkylbenzimidamides and 4-Nitrophenylnitromethane BY COLIN D. HUBBARD*, DELBERT L. HARRIS 111, DAVID W. HOOPER AND ARTHUR F. TUCCI Department of Chemistry, University of New Hampshire, Durham, New Hampshire 03824-3 598, U. S .A. Received 2nd April, 1982 Rate constants have been determined for the proton-transfer reaction of 4-nitrophenylnitromethane, N02C,H,CH,N0,, with two NN-dialkylbenzimidamides, HN=C(C,H,)NR,, where R = 1 -butyl, or I-propyl, in the temperature range 273-303 K in some aprotic solvents of low polarity. In the concentration ranges used the reaction to form an ion-pair product is first-order in acid and first-order in benzimidamide. The reactions were monitored spectrophotometrically employing the stopped-flow method. An examination of the relationships of the forward proton-transfer rate constants at 298 K, in six solvents, for the dibutyl compound to a solvent permittivity function and to the solvent polarity parameter (ET) shows that a polar activated complex develops; this is compatible with reaction in which ions are produced.The corresponding kinetic parameters for the dipropyl analogue in three solvents are consistent with this pattern. There is no significant effect upon the kinetic parameters of variation of the alkyl group in the benzimidamide. The energies of activation for forward proton transfer are small, in the range 10-23 kJ mol-', while in the reverse direction the energies of activation range is 65-1 14 kJ mol-l.A Brransted-type plot using the forward rate constants and the derived equilibrium constants at 298 K yields a coefficient which is indicative of a significant degree of proton transfer in the transition state. However, the highly exothermic nature of the reaction and the low energies of activation in the forward direction suggest a reactant-like activated complex. Provisional results (forward rate constants and energies of activation) for deuteron transfer from the deuterated acid to both benzimidamides in di-l-butyl ether and in chlorobenzene are reported. The hydrogen-ion transfer results are compared with those reported for reaction of the title acid with I, 1,3,3-tetramethylguanidine. The kinetics of the reaction between 4-nitrophenylnitromethane (4-NPNM) and 1,1,3,3-tetramethyIguanidine (TMG) have been studied in several aprotic solvents by Caldin and Mate0.l They observed large kinetic isotope effects and anomalous Arrhenius parameters and concluded that the carbon acid releases a proton to the base by a pathway .which involves quantum-mechanical tunnelling.2-6 The variations of the kinetic and activation parameters with solvent were interpreted as indicating that the transfer of the proton is coupled to motions of solvent molecules for the reaction in polar solvents such as acetonitrile, but not in non-polar solvents such as toluene.Studies by Caldin and coworkers using other bases, notably quinuclidine, triethylamine and tributylamine, in both toluene and other solvents (anisole and chlorobenzene) yielded results which were consistent with the solvent dependence noted earlier. The kinetic isotope effect, which was reported to be 45 at 298.0 K for TMG and 4-NPNM in toluene, took a much smaller value for the reactions with tertiary although still larger than could be attributed to zero-point energy differences only.This suggested that the barrier width for a given height may be less for the imine bases. The virtually constant value of the volume of activation, AV*, for the reaction of TMG and 4-NPNM in five solvents of low polarity was also interpretedll as support for the earlier analysis of solvent influence upon kinetic isotope effects. The kinetic isotope effects and activation parameters for the reactions 36193620 PROTON TRANSFER BETWEEN 4-NPNM AND TMG of 4-NPNM with two alkylimidarnides,l0 a cyclic imidamide12 and penta- methylg~anidinel~ were of magnitudes which were also consistent with the need for a tunnelling correction.Subsequently the reaction of TMG was further investigated, and the consequences of exchangeable hydrogen in TMG were exarnined.l4-l6 If the assumed mechanism of deuteron transfer was modified to incorporate the assumption of rapid exchange between the two forms of hydrogen on the imine nitrogen of the ion pair, then the kinetic isotope effect derived for the reaction in toluene was much lower (ca. 11 at 298.0 K).157 l6 Concern has been raised12 over the consequences of the presence of extraneous water, which is very difficult to remove from TMG, upon the kinetics.Residual water does not significantly affect the proton transfer to TMG, but the apparent rate of the deuteron transfer may be affected.17 Extraneous water does not appear to affect formation of the ion pair formed by pentamethylguanidine or its deuterated analogue in toluene and acetonitrilels or the loss of tritium from tritiated 4-NPNM reacting with TMG in t01uene.l~ Before the appearance of these latest publications we undertook a study of the reaction of 4-NPNM with two benzimidamide bases, NN-di- 1 -propylbenzimidamide and NN-di- 1 -butylbenzimidamide,t in various solvents with a view to determining the relationships of kinetic parameters with solvent .20 On the basis of results subsequently reported for the reaction of pentamethylguanidine and 4-NPNM in two solvents it has been suggestedlg that the interpretation of the dependence of the kinetic parameters upon solvent proposed earlier is not correct.This report is concerned primarily with the effect of solvent upon the kinetics of the proton-transfer reaction, but we report also some preliminary results on the deuteron-transfer reaction. EXPERIMENTAL MATERIALS 4-NPNM was synthesized by a literature methodz1 and stored in a darkened desiccator. [2H]4-NPNM was supplied by Dr C. J. Wilson. DPBA and DBBA were synthesized according to a literature method.,, (The latter imidamide has not been reported previously.) Deuterated DPBA, [2H]DPBA, was prepared by shaking a ten-fold excess of D,O with a sample of DPBA and then removing the solvent using a rotary evaporator.Following a repetition of this process the product was purified by vacuum distillation. Characterization of bases was by elemental analysis and proton n.m.r. spectroscopy. Bases were stored in darkened desiccators over dry nitrogen. Solvents used were purified from reagent-grade commercial products by standard methods using extra precautions to exclude atmospheric moisture. Purified solvents were kept in dark containers with, when appropriate, drying agents within the solvent. An atmosphere of dry nitrogen was maintained over the solvent, the containers of which were stored in sealed vessels over dehydrating agents. In experiments with carbon tetrachloride as solvent it was found that kinetic data for the reaction of DBBA with 4-NPNM were unchanged within experimental variation when Aldrich Gold Label spectroscopic grade solvent was used in place of solvent samples purified in our laboratory. METHODS Solution spectra, equilibrium measurements and some deuteron-transfer kinetics were recorded using a Cary 14 spectrophotometer.A wavelength of 450 nm was used for monitoring purposes. The temperature of solutions was controlled by using a Lauda K-2/R thermostatting sys tern. Most kinetic studies werecarried out using a Durrum-Gibson stopped-flow spectrophotometer 7 NN-di- 1 -propylbenzimidamide and NN-di- 1 -butylbenzimidamide are abbreviated to DPBA and DBBA, respectively.c. D. HUBBARD, D. L. HARRIS 111, D. w. HOOPER AND A. F. TUCCI 3621 with a 2 cm cuvette and a Kel-F flow system. Oscilloscope photographic records of the reactions monitored at 450 nm were analysed to determine the rate constants by a method previously All data recorded yielded first-order kinetic plots linear for at least three half-lives, except in some cases at low base concentrations when the formation of product is not favoured and the signal amplitudes are very small, or when a mixture of 4-NPNM and [2H]4-NPNM was deliberately used.The values reported for rate constants are averages of usually four to seven replicate determinations. The temperature was maintained constant throughout a given kinetic run by a circulating fluid whose temperature was controlled by a Neslab RTE-8 bath. The accuracy with which the temperature of solutions is known is reduced the further that is from the ambient temperature in the type of stopped-flow instrument used.RESULTS PROTON TRANSFER The kinetic data for reaction (I), kf 4-NPNM + DBBA (4-NPNM)- (DBBA)+. . . (1) kh the ion-pair product of which has an intense yellow colour, are consistent with Here kobs is the first-order rate constant obtained from reaction traces under the condition of base concentration in large excess over acid concentration. (Parallel equations may be written for reaction of DPBA.) Plots of kobs against total base concentration are linear over the range of base concentration studied, yielding the second-order forward and first-order reverse rate constants (k, and kb) from the slope and intercept, respectively. The kinetic data and the concentration conditions for reaction (1) are reported in table 1. A summary of kinetic results and derived activation parameters together with the standard deviations for these is presented in table 2.The excess of base was typically five to ten fold, usually in the region 1-10 mmol dmP3. The particular concentration range that can be used for a given reaction depends on the reaction rate, the need for sufficient base excess and the position of equilibrium for the temperature and solvent employed. Over the ranges studied the plots reveal no systematic deviation from linearity. Given that it is difficult to remove water from TMG then it seems unlikely that. benzimidamide bases are readily rendered anhydrous. Although normal precautions to exclude water were taken during the preparation and purification of reagents, the purification of solvents and the preparation of solutions and during kinetic experiments, undoubtedly trace amounts are present.The interpretation of results, however, is based on the assumption that extraneous water has no significant effect upon the proton-transfer reactions studied since no effect is detected. In one reaction, that between DPBA and 4-NPNM in carbon tetrachloride, solvent for the base was saturated with water. At 293 K k,H, the forward rate constant (48 mol-l dm3 s-l), is within experimental error of the value for the dry solvent, and the reverse rate constant is changed from ca. 2.5 to 2.0 s-l, but this change is almost within the range of experimental error. Other studies of the effect of added water upon the kinetics of carbon acid-base reactions have shown a similar result to that reported here; uiz.the rate of proton transfer from 4-NPNM to 1,8-bis(dirnethylamino)-2,7-dimethoxy- naphthalene in chlorobenzene is unaffected when the solvent is saturated with water but the reverse rate is lowered by about ~ n e - h a l f . ~ ~ A comparison with some proton-transfer rates to a benzimidamide in anisole is possible. In an earlier study, k,H (for NN-diethylbenzimidamide) was reported as eqn (2): kobs = k,[DBBA] + k,. (2)3622 PROTON TRANSFER BETWEEN 4-NPNM A N D TMG TABLE 1 .-RATE CONSTANTS FOR REACTIONS OF 4-NITROPHENYLNITROMETHANE WITH DBBA AND DPBA IN VARIOUS SOLVENTS solvent carbon tetrachloride di- 1 -butyl ether anisole chlorobenzene tetrahydro fur an dichloromethane [DBBA]/IO-' -kk,H/mol-l temp/K mol dm-3a7 a v .kYbs/S-' dm3 s-" kF1s-l 283 288 293 298 28 1 284.6 288.4 293.9 280.7 284.9 289.0 290.9 293 298 303 280.5 285.4 29 1.4 298 304 275.8 278.6 291.3 298 303 278.3 290.6 295.5 298.5 5.0-25.0 7.0-27.0 7.0-27.0 10.0-33.0 3.0-30.0 7.0-30.0 3.0-30.0 3.0-30.0 3.0-24.0 2.0-26.0 3 .O- 1 2.0 2.93-24.4 3.0-24.7 1.82-24.4 2.93-18.6 3.86-8.69 2.90- 19.3 2.90- 17.4 1.93- 17.4 4.83-21.2 0.43-2.68 0.43-3.42 3.35-21.4 8.43-20.1 3.35-21.4 2.30-9.20 4.90-24.5 2.4- 12.0 2.3-9.2 0.6 1 - 1.42 1.33- 1.97 1.95-2.83 2.35-4.27 1.57-4.58 3.13-5.73 3.32-7.01 5.76-9.9 1 2.02-7.36 3.27-7.15 4.56-6.65 5.01-10.4 6.92-10.9 9.60- 14.9 1 2.6- 1 6.8 2.46-3.64 3.80-7.56 5.2 1-9.20 9.45-13.7 17.5-22.4 0.508-0.963 0.646- 1.39 3.69-8.67 6.5 1 - 10.4 7.67- 1 5.0 1.50-3.69 5.02-14.3 6.27- 10.2 5.34-9.14 43.2k 4.3 33.0 f 2.1 40.8 f 3.9 83.9 & 25.0 11 1.9f2.3 1 19.5 f 7.4 141 f9.9 138f 18 208 f 8.30 141 f 25 226 f 87 238f 13 183k 11 228 f 30 245 & 27 242f 16 229 f 1.1 271 f 8.0 265 f 43 320 k 59 211f33 259 f 33 275 f 5.2 329 f 43 421 f 4 7 339 f 25 484f 116 451 f.87 530 f 27 0.37 f 0.072 1.07 k 0.039 1.70 & 0.072 1.14f 0.060 1.27 f 0.04 2.30 f 0.14 3.06 f 0.18 5.92 k 0.33 1.63 f 1.24 2.87 f 0.35 3.68 f 0.74 4.36 f 0.20 6.54 & 0.17 8.82 f 0.40 12.1 f0.3 1.55 k 0.10 3.14f0.01 4.47 f 0.09 9.24 f 0.45 16.1 f0.8 0.446 f 0.058 0.533 f0.065 2.76 f 0.07 3.90 f 0.64 6.32 f 0.63 0.610f0.164 4.10 f 1.80 5.33 f0.66 4.28 k 0.17 [DPBA]/ solvent temp/K mol dm-3a. carbon tetrachloride 293 di- 1 -butyl ether 28 1 284.6 293.9 chlorobenzene 279.9 283.8 288.1 294.6 299.1 300.0 10-25 5.0-30 3.0-30 10-30 4.0-20 4.0-20 2.5-20 6.0-20 4.0-20 3.0-25 3.1-4.0 1.49-4.55 1.80-5.67 7.62-10.4 2.20-6.72 2.69-8.74 3.16-11.2 9.26-15.6 13.0-20.6 16.9-27.9 62+ 10 112f8 130k 19 148 f 18 292 f.22 393 f 34 434 f 34 463 f 59 442 f. 109 533 f 67 2.54 k 0.18 1.10f0.14 1.55 f 0.39 6.14 f 0.39 0.96 f 0.30 1.25 f 0.44 2.46 k 0.40 5.97 f 0.82 12.8f 1.4 13.5f 1.1 a The ranges of initial concentrations of the two benzimidamides used are given. The initial mol dm-3 except in a few cases where an initial concentration of mol dmP3 was used. The value of kobs was independent of initial acid concentration within Values of k p and kb were obtained from plots of kobs against base concentrations concentration of 4-NPNM was 5.0 x 1 .O x experimental error.according to eqn (2). Uncertainties are standard deviations. 283f 7 mol-1 dm3 s-l at 298 K,1° while in the present study a value of k,H of 228+30 mol-1 dm3 s-l has been determined for the dibutyl analogue at the same temperature. This supports the present kinetic data, particularly at temperatures close to ambient where temperature control is satisfactory.C. D . HUBBARD, D . L. HARRIS 111, D . W. HOOPER A N D A. F. TUCCI 3623 TABLE 2.-sUMMARY OF KINETIC AND ACTIVATION PARAMETERS FOR REACTIONS OF DBBA AND DPBA WITH CNPNM AT 298 K (EXCEPT AS NOTED) k,H / 1 0-2 mol-l solvent dm3 s-l KF/s-l E$/kJ mol-l E,Hb/kJ mol-' KH(k) KH(sp) cc1, DBE C,H,OCH, CB THF CH,Cl, cc1, DBE CB 0.84f0.25 1.48 f 0.19 2.28 f 0.30 2.65 k 0.43 3.29 f 0.43 5.24 f 0.26 0.62 f 0.10 1.61 k0.20 4.18 f 0.79 (293) DBBA 1.14 f 0.59 7.51 f0.42 8.82 _+ 0.40 9.24 f 0.45 3.90 0.64 4.15 f0.17 DPBA 2.54f0.18 8.12 f 0.52 13.4+ 1 .1 (293) - - 12.3 f4.6 79.5 + 5.3 9+7 63f4 11.2f3.5 65.2 f 5.8 14.3 f 3.2 69.8 f4.5 23.2 f 7.3 114f26 13.8 f 3.9 94.0f 8.8 15.1 f4.1 96.9 f 4.9 - 19.7 25.9 28.7 84.4a 71.9b (297.6) 126 19.8 31.2 a KH(k), (range 276-303 K) A H e = - 55.3 f 6.6 kJ mol-l; b KH(sp), (range 279-303 K) AHe = - 58.1 f 3.4 kJ mol-'. DEUTERON TRANSFER It has been noted earlier that extraneous water affects the deuteron-transfer reaction more than the proton ionisation for the TMG system in t01uene.l~ Therefore the limited data reported below for the deuteron-transfer reaction, for the benzimidamides with 4-NPNM in chlorobenzene and in di-1-butyl ether and for DBBA in anisole, are probably more subject to the effect of the presence of water, even though such an effect is not readily apparent, and are to be treated as provisional.For the reaction in anisole there is a further reason for caution. After being allowed to stand in anisole for a few hours a sample of [2H]4-NPNM loses a significant fraction of initial deuterium. This is shown by the appearance of two steps widely separated in time in the stopped-flow spectrophotometer. The first step corresponds to the proton-transfer reaction (in the tenths of a second range), and the second stage occurs over several seconds, which corresponds to the loss of a deuteron from the acid. The results reported are from experiments conducted within about an hour of making the anisole solution of [2H]4-NPNM, before the hydrogen exchange on the acid has reached a very significant proportion.This rapid exchange process did not occur with other solvents, in which the deuterated acid appears to be stable virtually indefinitely. It is not known what component of the solvent or impurity facilitates this rapid exchange in anisole. The effect of this spurious exchange upon the exclusive proton-transfer reaction (k?) would be negligible. At 298 K apparent values of k,D are 5.3240.04 and 4.82f0.40 mol-l dm3 s-l in di-1-butyl ether and 14.1 1.0 and 12.5k0.3 mol-l dm3 s-l in chlorobenzene for reactions of DBBA and DPBA, respectively. These are extrapolated from results at other temperatures. In anisole the value of k,D, measured at 298 K for the reaction of DBBA, is 7.98 4 0.17 mol-1 dm3 s-l.The range of excess base and concentrations of acid and base are similar to those used in the proton-transfer studies. For deuteron transfer to deuterated DPBA in chlorobenzene apparent values of k,D are 15.8 0.6 mol-1 dm3 s-l at 294.9 K and 12.9 1 .O mo1-1 dm-3 s-l at 297.5 K for two independent determinations. When eqn (2) is used for deuteron-transfer reactions it is not possible to determine k? reliably because the intercept seems to be anomalously low, a situation noted by others for similar systems.? t For example ref. (1).3624 PROTON TRANSFER BETWEEN 4-NPNM A N D TMG ACTIVATION PARAMETERS A N D EQUILIBRIUM CONSTANTS The Arrhenius activation parameters, where determined, are subject to a higher probable error than those determined for the TMG system.This arises because standard deviations in kf and k, are generally larger than those reported earlier; in turn this can be attributed in part to the less favourable equilibrium position (the benzimidamides are weaker bases than TMG) yielding smaller signal amplitude changes and in part to the less satisfactory temperature control of the reactant and product solutions with the available instrument. The relatively small possible temperature range that can be employed also contributes to the magnitude of error in the activation parameters. Despite these difficulties the ranges of activation energy are clear (table 2). For forward transfer of the proton the range is E z = 10-23 kJ mol-1 and for the reverse reaction the activation energy EFb is 65-1 14 kJ mo1-1 for DBBA in five solvents.For reactions of DPBA in di-1-butyl ether and chlorobenzene, E$ is 14-15 kJ mol-1 while E,H, is 94-97 kJ mol-l. Provisional values (all in kJ mol-l) of the activation energies for the deuteron-transfer reactions are for reaction of DBBA : E,D = 37 +4 (six temperatures, 272-300.2 K) in di- 1-butyl ether, E s = 22 f 3 (three temperatures, 277.6-298 K) in anisole, and EnDf = 33.4+ 7.5 (four temperatures, 277.6-293 K) in chlorobenzene; and for reaction of DPBA, E$ = 39 f 5 (five temperatures, 280.5-299.8 K) in di-1-butyl ether, EaDf = 42.5 1I3 (four temperatures, 273.4-300.3 K) in chlorobenzene. A series of determinations at four temperatures using deuterated DPBA in chlorobenzene, over the range 28 1.6-297.5 K, yields a provisional value of EaDf = 42f8.8 kJ mol-l.No determinations of Epb are possible since k; is not calculated, as indicated above. Values of the equilibrium constant for the formation of the protonated base, derived from the ratio of the forward and reverse rate constants, range from 20-30 mol-l dm3 in very low-polarity solvents such as di-1 -butyl ether (D = 3.06) and carbon tetrachloride (D = 2.28) to ca. 130 mol-l dm3 in the relatively high-polarity solvent dichloromethane (D = 9.08). The values of KH (table 1) are subject to a significant error (not reported) as a derived value, but a general trend of increasing strength of the acid-base reaction with increase in dielectric constant is observed. In one solvent, tetrahydrofuran, the equilibrium constant has been measured spectro- photometrically, and the value obtained, 72 mol-1 dm3 at 298 K, is reasonably consistent with the kinetically derived value (84 mol-1 dm3).In this solvent the reaction enthalpy change, A@, is determined as - 58.1 f 3.4 and - 55.3 f 6.6 kJ mol-l by the equilibrium and kinetic methods, respectively, in satisfactory agreement. Considering the error associated with the determination of k g , calculation of KD is not warranted. If KD has a similar magnitude to KH, as would be expected, then the intercept of the plot of kobs against benzimidaniide concentration, although small, would be detectable. However, it is very small and not distinguishable from zero. There are several possible reasons for the observation being at variance with the prediction, but because of the provisional status of the deuteron-transfer results these will not be described.Use of deuterated NN-di- 1 -propylbenzimidamide in an effort to eliminate isotopic exchange yields results which also give rise to low intercepts on the kobs axis. This, however, is an ambiguous result since traces of water present could facilitate the dedeuteration of the base.C. D. HUBBARD, D. L. HARRIS 111, D. W. HOOPER AND A. F. TUCCI 3625 DISCUSSION The forward proton-transfer reaction is characterized by rate constants in six solvents which range from 84 dm3 mol-1 s-l in carbon tetrachloride to 524 dm3 mol-1 s-' in dichloromethane, for DBBA at 298 K. The dipropylbenzimida- mide reaction with 4-NPNM has rate constants similar to those for DBBA in the three solvents in which it has been studied.The results are consistent with a pattern of partial transfer of charge in the activated complex, and parallel those of Caldin and Mateol for the reaction of TMG with 4-NPNM. In the latter reaction with the stronger base, a greater range of solvent dielectric was used but the magnitude of solvent dielectric influence on rate was smaller. In di-I-butyl ether and in chlorobenzene, at 298 K, k p and k? values vary little with varying alkyl group on the base, indicating that the bulkiness of these offer similar hindrance or no hindrance to the acceptance of the proton. This pattern is repeated in comparing the values of k p in anisole for DBBA and the diethyl analogue,l0 and is sustained in consideration of the deuteron-transfer reactions to both DBBA and DPBA in di-1 -butyl ether and in chlorobenzene.The equilibrium constants, reflecting the completed acid-base reaction, increase with increase in solvent polarity, are in agreement with prediction and mirror the results for the substituted guanidine reaction. Energies of activation for the forward proton-transfer reaction are very small (10-23 kJ mol-l) and imply that the entropies of activation are large and negative. For the reaction with DBBA a plot of the logarithm of k p against the solvent permittivity function (D - 1)/(2D + 1) (employing the relationship derived by K i r k ~ o o d ~ ~ ) is reasonably linear (fig. 1) with a positive slope, indicating the development of a polar activated complex and supporting the model of solvent dielectric effects upon the kinetics of reaction of neutral molecules where non-electrostatic influences of solvent upon the activated complex are similar to those upon the reactants.25 The range of values of ET, the empirical solvent polarity parameter, for the solvents used26 is limited, but the plot is consistent with a linear dependence of log k,H on ET (fig 2), as expected for a reaction producing ions.Both plots are reasonably similar to those found for the corresponding reaction with TMG. Results for proton and deuteron transfer in toluene from 4-NPNM to various bases, where isotope exchange on the base has been incorporated in the kinetic scheme and where extraneous water is considered to have a negligible effect, have been plotted by Rogne et al.14 in the form of a Brarnsted-type plot (log k,H and log kF against log KH).The results reported here for k p and KH in di-1 -butyl ether and in anisole (which have dielectric constants fairly close to that of toluene) fit well to the lower of the two lines described. The slope gives a Brarnsted coefficient of ca. 0.9, indicative of a high degree of transfer of the proton in the transition state. This is difficult to reconcile with the very low energy barrier and the highly exothermic nature of the reaction in the forward direction, which would give rise to the prediction of a reactant-like configuration in the transition state. It appears that the proton transfer is anomalous in that the energies of activation (10-23 kJ mol-l) are barely greater than that characteristic of diffusion control, yet the rate constants are seven or eight orders of magnitude lower than typical values for diffusion control at 298 K.In contrast the provisional values of E z are 30-40 kJ mol--l. The isotope effect (kfH/kf") based on provisional values of kF is in the range 20-30 at 298 K. Both of these results, uiz. AEa = E,D,-E,H, > 4-5 kJ mol-l and kfH/kp > 7, are thought? to be criteria for tunnelling, and lead to the requirement of a tunnelling correction for the reaction rate. Hence these results would imply that t For example chap. 4 of ref. (5).3626 PROTON TRANSFER BETWEEN 4-NPNM AND TMG 2.71 = i (D -1)(2D + I ) FIG. I.-Plot of logk? for proton transfer from 4-NPNM to DBBA against the solvent permittivity function, ( D - 1)/(2D+ l), in various solvents at 298 K.See text for the basis of this plot. The points in order of increasing value along the abscissa are for the reaction in carbon tetrachloride, di- 1 -butyl ether, anisole, chlorobenzene, tetrahydrofuran and dichloromethane. FIG. 2.-Plot of log kB for proton transfer from 4-NPNM to DBBA against the empirical solvent polarity parameter, ET, in various solvents at 298 K. The points in order of increasing value along the abscissa are for the reaction in carbon tetrachloride, di- 1 -butyl ether, anisole, tetrahydrofuran, chlorobenzene and dichloromethane. tunnelling may occur in the transfer of a proton from 4-NPNM to benzimidamides. Since, however, the deuteron results are as yet only provisional, a conclusion regarding tunnelling is not justified at present.We have considered the application to the present work of the two principal criticisms levelled at earlier work in this area where tunnelling has been invoked. The influence of extraneous water upon the reactions reported here can be predicted to be insignificant for proton transfer but may not necessarily be disregarded for deuteron transfer. The effect of rapid hydrogen exchange upon the imine nitrogen of the base acceptor has been fully discussed.15~ l9 For the benzimidamide systems, rapid isotopic exchange would have even greater consequences for producing an apparent rate constant which does not truly represent kp, since the ratio of k t to k,D is much higher than it would be in the guanidine system. Prevention of isotopic mixing can be achieved by deuterating the amine nitrogen before reaction occurs.Our limited dataC. D. HUBBARD, D. L. HARRIS 111, D. W. HOOPER AND A. F. TUCCI 3627 for the reactions of deuterated base with deuterated 4-NPNM are compatible with a scheme in which isotopic exchange is not rapid. However, it is conceivable that reactions of both deuterated and non-deuterated base could be affected by extraneous water. We hope to report later the results of further measurements made with special precautions to exclude water. The authors are indebted to Dr C. J. Wilson for discussions prior to this work and to Prof. E. F. Caldin for sharing results prior to their publication. C.D.H. also acknowledges with gratitude extensive personal discussion and correspondence with Prof.Caldin. Neil Linsky and Elizabeth Corless are thanked for their assistance with preliminary experimental work, and Philip Duce is thanked for help with some calculations. This research was supported in part by the Central University Research Fund of the University of New Hampshire, Graduate School and Research Office. E. F. Caldin and S . Mateo, J. Chem. SOC., Faraday Trans. I , 1975, 71, 1876. R. P. Bell, Proc. R. SOC. London, Ser. A, 1933, 139, 466. R. P. Bell, The Proton in Chemistry (Chapman and Hall, London, 2nd edn, 1973), chap. 12. R. P. Bell, The Tunnel Effect in Chemistry (Chapman and Hall, London, 1980). L. Melander and W. H. Saunders, Reaction Rates of isotopic Molecules (Wiley, New York, 1980), pp. 36, 75, 140 et seq. E. F. Caldin and S . Mateo, J . Chem. SOC., Faraday Trans. I , 1976, 72, 112. E. F. Caldin and C. J. Wilson, Faraday Symp. Chem. SOC., 1976, 10, 121. E. F. Caldin, D. M. Parbhoo, F. A. Walker and C. J. Wilson, J . Chem. SOC., Faraday Trans. I , 1976, 72, 1856. C. D. Hubbard, C. J. Wilson and E. J. Caldin, J . Am. Chem. SOC., 1976, 98, 1870. * E. Wigner, 2. Phys. Chem., Teil B, 1932, 19, 203. lo E. F. Caldin, D. M. Parbhoo and C. J. Wilson, J . Chem. Soc., Faruday Trans. 1, 1976, 72, 2645. l2 E. F. Caldin, 0. Rogne and C. J. Wilson, J. Chem. SOC., Faraday Trans. I , 1978, 74, 1796. l3 E. F. Caldin and 0. Rogne, J. Chem. SOC. Faraday Trans. I , 1978, 74, 2065. l4 I. Heggen, J. Lindstrom and 0. Rogne, J . Chem. SOC. Farada-y Trans. I , 1978, 74, 1263. l5 0. Rogne, Acta Chem. Scand., Part A , 1978, 32, 559. l 7 E. F. Caldin, S. Mateo and P. Warrick, J. Am. Chem. SOC., 1981, 103, 202. I s J. H. Blanch, 0. Rogne and L. I. Rossemyr, J. Chem. Soc. Faraday Trans. I , 1980, 76, 1905. l y A. J. Kresge and M. F. Powell, J. Am. Chem. SOC., 1981, 103, 201. 2o C. D. Hubbard and D. W. Hooper, Phys. 106, Abstracts of 2nd CIC/ACS Meeting, Montreal, P.Q., Canada, June 1977. *I R. G. Cooke and A. K. McBeth, J. Chem. Soc., 1938, 1024. 22 R. P. Hullin, J. Miller and W. F. Short, J . Chem. SOC., 1947, 394. 23 C. D. Hubbard and J. F. Kirsch, Biochemistry, 1972, 11, 2483. 24 S. Mateo and C. Mateo, Abstracts of the Second international Conference on Mechanisms of Reactions 25 J. G. Kirkwood, J . Chem. Phys., 1934, 2, 351. 26 C. Reichardt, Angew. Chem., Int. Ed. Engl., 1965, 4, 29. J. H. Blanch and 0. Rogne, J . Chem. SOC., Faraday Trans. I , 1978, 74, 1254. in Solution, Canterbury, Kent, 1979. (PAPER 2/567)

ISSN:0300-9599    

DOI:10.1039/F19827803619

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 3629-3643 Solvation of Silver Ions in a Range of Binary Solvent Systems Radiation and Electron Spin Resonance Study? BY MARTYN C. R. SYMONS* AND COLIN K. ALESBURY Department of Chemistry, The University, Leicester LE1 7RH Received 2nd April, 1982 Exposure of dilute solutions of silver perchlorate in a range of pure and mixed solvents to 6oCo y-rays at 77 K gave Ago and Ag" centres with reduced silver hyperfine coupling indicating delocalisation onto solvent molecules. In certain cases this was confirmed by the detection of hyperfine coupling to ligand nuclei. Solvents used included water, methanol, cyanomethane, acetone, methyltetrahydrofuran, dimethyl- sulphoxide, dimethylformamide, ethylamine, triethylamine, pyridine, ammonia, benzene and toluene.Results for the Ago centres are interpreted in terms of primary species, which retain the solvation of the original Ag+ ions, and secondary centres, which have lost one or more of the primary solvent molecules of the cation, the limiting species being generally the unsolvated atoms. In some cases Ag:, Ag$+ and Ag:+ clusters were also detected either at 77 K or after annealing. A linear correlation between the isotropic hyperfine coupling to lo9Ag and the solvent donor number was followed by the majority of pure solvents, the most marked exceptions being pyridine and cyanomethane. Reasons for this trend and the exceptions are discussed. Other species with relatively low isotropic hyperfine coupling to lo9Ag were detected, and possible structures are suggested.We have recently shown that Ago and AgI1 centres formed from Ag+ ions in dilute solutions in cyanomethane both retain the four primary solvent molecules coordinated to Ag+ prior to irradiation, as evidenced by hyperfine coupling to 14N.l Using mixed (MeCN+ H,O) and (MeCN + MeOH) solvent systems,2 we went on to show that Ag+ ions are remarkably strongly bound to MeCN molecules, the tetrasolvate dominating to beyond the 50% mole-fraction range. That Ag+ ions are preferentially solvated by MeCN in (MeCN+H,O) systems has also been adduced by Cox et al.,3 and there is conductometric evidence that Ag+ coordinates to four MeCN molec~ies.~ Our conclusions, that in these radiolyses electron-capture by Ag(MeCN)a gives Ago(MeCN), and electron-loss gives Ag(MeCN)i+, have recently been q~estioned.~ In fact, the argument centres on the silver-atom species only.The most interesting result of exposure of solutions of AgClO, in MeCN to X-rays at 4 K was that the yield of Ago(MeCN), centres was reduced by a factor of ca. 1 /6, but that on annealing to 77 K this increased six-fold. The most probable reason for this is that at 4 K MeCN- or (MeCN); centres6 are formed, and that these transfer their electrons to Ag+ on annealing.5 However, because a silver-atom centre containing only one MeCN molecule was formed on photolysis, it was concluded that the Ag+ ions are only solvated by one MeCN molecule, the Ago(MeCN), centre being the primary centre, with Ago(MeCN), formed therefrom because of an atom solvation process. Our reasons for rejecting the conclusions of Li and Kevan will be outlined el~ewhere.~ Our major case rests on the fact that the compound Ag(MeCN), C10, gives only the centres Ago(MeCN), and AgII(MeCN), on irradiation.Herein we use our original assignments and treat the Ago(MeCN), centre as a partially desolvated unit, intermediate between Ago(MeCN), and unsolvated Ago. 7 Taken as Solvation Spectra, Part 70. 36293630 SOLVATION OF SILVER IONS EXPERIMENTAL Deuterated solvents, including [2H6]acetone (Me2CO), [2H6]dimethylsulphoxide (DMSO), [2H3]cyanomethane (MeCN), [zH,]NN-dimethylformamide (DMF) and [2H,]toluene, which were all supplied by Merck, Sharp & Dohme (Canada) Ltd, were of the highest grade available. They were used as supplied after drying and removing oxygen by freeze-thaw procedures.Solution handling, irradiation and e.s.r. spectroscopic methods were as described previ~usly.~ RESULTS AND DISCUSSION We consider first the range of silver centres formed in the pure solvents, before turning our attention to the mixed-solvent systems. No attention is given to the range of solvent radicals also detected in the g = 2 region of the spectra. PURE SOLVENTS In all cases studied, replacing the normal solvent with its perdeuterated form resulted only in line-narrowing. Most of the following discussion refers to perdeu- terated media. ACETONE As can be seen in fig. 1, the major silver centres detected at 77 K were Ago, Ag*I and Agl (data are given in tables 1-3). The Agi centre is characterised by a major triplet, the central features being partially concealed beneath intense features from solvent radicals.* The outer (MI = f 1) features appear as triplets.This is because TABLE 1.-Ago CENTRES DETECTED IN PURE- AND MIXED-SOLVENT SYSTEMS AT 77 K solvent species g,, 14N lo9Ag CD,CN (CD3)2C0 DMSO MeTHF nitromethane triethylamine 33% aq. NH, form amid e DMF toluene CD3CN + D,O CD,CN + (CD,),CO A B A B A B A B 1.997 2.001 2.003 2.003 2.003 2.002 2.001 1.998 2.001 2.001 2.001 1.995 1.996 1.999 1.998 2.000 2.000 2.001 2.003 2.002 2.000 +6 -532 -612 - 643 - 705 - 679 - 626 -601 - 578 - 636 - 597 - 661 -511 - 455 - 647 - 603 - 679 +6 -587 - 665 +7 -679 -721 - 555 a G = 10-4 T.M. C. R. SYMONS A N D C. K. ALESBURY 363 1 1 32506 A (a ) FIG. I b 1 3250G A - 2 SOG , I * ' A I Gain x 10 FIG.1.-First-derivative X-band e.s.r. spectra for AgC10, in (CD,),CO after exposure to 6oCo prays at 77 K : (a) before annealing, showing features assigned to Ago and Ag" centres, and (b) after annealing to 120 K, showing additional features assigned to A&+ centres. Central features for solvent radicals are not shown.3632 SOLVATION OF SILVER IONS TABLE 2.-Ag" CENTRES DETECTED IN PURE- AND MIXED-SOLVENT SYSTEMS AT 77 K AisoIG 14N 10SAg 14N 109Ag solvent gl g , ga, CN,CN (CD3)2C0b C DMSO nit rome thane pyridine 33% aq. NH, formamide DMF CD,CN + (CD,),CO mole fraction (CD,),CO: 0.76 C CD,CN +DMSO mole fraction DMSO: O.25-0.tl6 C 2.312 2.312 2.361 2.293 2.360 2.189 2.159 2.305 2.321 2.31 1 2.360 2.283 2.302 U 2.067 2.067 2.059 2.076 2.037 2.062 2.065 - 2.067 2.067 2.058 2.058 - 2.149 2.165 2.137 - 2.148 2.165 2.133 2.137 21 46 26 - 26 - 26 34 - 30 32 - 25 29 22 - 18 36 27 - 27 32 - 27 - - - - - - - - - - - - 32 - 24 35 - 30 - - 24 - 25 32 - 25 - - - 26.0 31.3 27.3 - 26.7 31.7 24.7 27.3 a Features not resolved; before annealing; after annealing.centres containing lo9Ag+lo9Ag, lo9Ag+lo7Ag and lo7Ag+lo7Ag give different splittings. (Both lo9Ag and lo7Ag have I = 4. They are present in nearly 50% abundance. Both have negative magnetic moments, that for lo9Ag being slightly greater in magnitude than that for lo7Ag.) These features exhibit little anisotropy. On annealing the features for Agz first narrowed and then decayed irreversibly, with features assigned to Agi+ growing-ing [fig. 1 (b)]. These features together with those for the remaining centres decayed at ca. 155 K.A range of Ago centres were detected at 77 K. Unfortunately we were unable to irradiate at 4 K so cannot judge conclusively which of these centres is the primary centre. However, as with MeCN systems, we assume that, in general, as the lo9Ag hyperfine splitting increases, so the solvent molecules originally solvating the Ag+ ions are being discarded. As stressed elsewhere, however, this is not necessarily true, since 5 s - 5 ~ admixture will result in a fall in the magnitude of the isotropic coupling and this is most likely to occur for unsymmetrical complexes.1o However, this is expected to result in g-anisotropy, with a shift to g-values < 2.00. Such shifts were not detected for the present centres.Thus the centre labelled A is probably the primary species. Centre F may not be a genuine species, since its features were always so weak that they could not be clearly resolved. The most atom-like centre, C (spin density ca. 0.99), is probably a monosolvate. Spectra for unsolvated silver atoms were not detected in this solvent. These features had all decayed at ca. 135 K. These results are interesting since they clearly establish the presence of a range of solvated silver units, but the work of Kevan and coworkers establishes that it will be necessary to work at lower temperatures before the full picture can be e~tablished.~? l1 The AgI1 spectrum changed markedly on annealing to 130 K, the parallel lines moving from 2.312 to 2.361, All(logAg) increasing from 26 to 34 G concomitantly.This probably indicates a relaxation towards the square-pyramidal arrangement of solvent molecules expected for this 4d complex. It is significant that no such change occurredM. C. R. SYMONS A N D C. K. ALESBURY 3633 TABLE 3.-sILVER CLUSTER AND OTHER CENTRES DETECTED IN PURE- AND MIXED-SOLVENT SYSTEMS AT 77 K solvent species giso Aiso(’”Ag)lG CD,CN (CD3)2C0 DMSO MeTHF ni trome thane [2H2]ethylamine pyridine 33% aq. NH, formamide DMF benzene toluene CD,CN + CD,OD CD,CN + (CD,),CO mole fraction (CD,),CO: 0-0.4 CD,CN + DMSO mole fraction DMSO: 0.25-0.8 CD,CN + MeTHF mole fraction MeTHF: ca. 0.75 DMF CD,CN + (CD,),CO mole fraction CD,CN: 0.5-0.9 1.971 1.971 1.929 1.979 1.965 1.951 1.964 1.996 1.968 1.996 - - 1.967 1.968 1.957 1.973 1.968 1.973 1.954 1.972 1.971 1.971 1.963 1.953 1.978 1.965 1.976 1.965 2.002 2.002 - 267 - 280 -112 - 273 - 129 - 138 - 148 - 225 - 157 - 200 - 130 - 100 - 290 - 267 - 137 - 194 -281 - 194 - 135 - 267 - 134 - 267 - 182 - 124 -271 - 132 - 276 - 135 - 100 - 82 a See text.on annealing above 77 K for the AgII(MeCN), centre, our conclusion being that the expected square-planar structure had already been acquired at this temperature.2 Interestingly, features for AglI at 4 K were very weak and poorly defined.5 This may reflect the inability of the tetrasolvate to change from its original configuration to the square-planar form at 4 K. The features for the 77 K centre grew in markedly on annealing, as e~pected.~ DIMETHYLSULPHOXIDE Unexpectedly, the main centre formed by electron addition at 77 K was Agi (fig.2). This probably implies some aggregation prior to irradiation since extensive mobility at 77 K is unlikely. Since DMSO is a strong ligand, as indicated by the low3634 SOLVATION OF SILVER 1 1 IONS FIG. 2.-First-derivative X-band e.s.r. spectra of AgCIO, in (CD,),SO after exposure to 6oCo y-rays at 77 K, showing Ago, Ag: and AgU centres. Central features for solvent radicals are not shown. value of AisO(lo9Ag) for the Ago centre, it is unlikely that Ag+-ClO; aggregates are present in the solutions. Only a single Ago centre was detected, in relatively low yield. The Agl* centre was well defined, and did not change on annealing. As usual, on annealing the Agi features gave way to Ag;+ features before the medium became fluid.METHYLTETRAHYDROFURAN This solvent was selected as typical of monoethers because good glasses are obtained on cooling. Two types of Ago centres were formed, as indicated in fig. 3, these species being dominant at 77 K. The spin densities of 0.90 (A) and 0.84 (B) indicate considerable delocalisation onto solvent ligands. The very low relative yield of AglI centres is expected for this solvent, which is unable to react with electrons, but traps the primary electron-loss centres efficiently.12 This is well illustrated by the fact that in the absence ofAgC10, the ejected electrons are shallowly trapped, giving a blue-black coloration to the glasses. In the presence of AgCl, the glasses become pale yellow, showing that electrons are very efficiently scavenged at 77 K.On annealing, Agi centres and later Ag;+ centres were formed, but no unsolvated silver atom features were detected. NITROMETHANE Pure MeNO, gives Me’ and ‘NO, radicals on irradiation at 77 K, neither electron-gain nor electron-loss centres bring trapped.13 Hence, as expected, both Ago and AglI centres were formed in high yield from solutions of AgClO, in this solvent. The Ago centre features were extremely broad, the large hyperfine coupling to lo9Ag ( - 661 G) indicating relatively minor delocalisation, as expected for this weakly basic solvent. The g-value of 1.995 may reflect slight admixture of 5p orbitals into the wavefunction, so that the actual spin density could be even higher than the value of ca. 92% deduced from Aiso. Agi+ centres grew-in on annealing, without any clear intermediate formation of Agz or Agi+ centres.M.C. R. SYMONS A N D C. K. ALESBURY 3635 ? 1 Gain x 100 I Sdwnt radicals FIG. 3.-First-derivative X-band e.s.r. spectra of AgClO, in MTHF after exposure to 6oCo prays at 77 K, showing features assigned to AgO (A and B), Ag; and central features for MTHF radicals. ETHYLAMINE A N D TRIETHYLAMINE The only well-defined product using EtNH, or EtND, was an Agi centre. For triethylamine solutions very broad Ago features were obtained, with a low value for Ai,,(lOgAg) (ca. -511 G). In this case no cluster centres were detected, even on annealing, nor were any AglI features detected. The contrast between these two solvents is most interesting. Triethylamine is an aprotic, strongly basic solvent, able to solvate Ag+ ions strongly, but unable to form hydrogen bonds to perchlorate ions.It might therefore be expected that ion clusters would be favoured in this solvent rather than in ethylamine, which can form hydrogen bonds to anions. Possibly ion clusters such as that in I (see later) tend to form in methylamine. Solvent-shared ion pairs do seem to be favoured in basic-protic media, especially at low temperatures. l4 PYRIDINE The major centre in this solvent was an AglI species for which the perpendicular feature showed well-defined hyperfine coupling to 14N nuclei (probably four). The very broad parallel feature was, unfortunately, unresolved in part because of overlap with broad features assigned to Agi (fig. 4). The central well-defined triplet is due to 1 -pyridyl radi~a1s.l~ These results suggest that pyridine may be a good solvent for forming electron-loss centres in preference to electron-gain centres.AQUEOUS AMMONIA Although liquid ammonia was not studied, aqueous ammonia glasses up to ca. 30% ammonia were used. Since silver halides are soluble in this medium they were used in addition to AgClO,; the initial results were independent of the anion, showing that ion-pairing or complex formation was not important. Very broad Ago features dominated the spectra at 77 K. Although hyperfine coupling to I4N is almost certainly3636 SOLVATION OF SILVER IONS n i , 32506 Gain x l o 4 Solvent radicals FIG. 4.-First-derivative X-band e.s.r. spectra of AgC10, in pyridine after exposure to 6oCo y-rays at 77 K and slight annealing, showing well-defined features for Ag" centres and the triplet features for pyridyl radicals.log 2+ I I I I I I I I I Ag, + 4 +3+2 +1 0 - 1 - 2 - 3 - 4 107 2+ I I I I I I I I I Ag, +4+3+2+1 0 - 1 - 2 - 3 - 4 FIG. 5.-First-derivative X-band e.s.r. spectra of AgF in 33% (as.) NH, after exposure to 6oCo y-rays at 77 K and annealing, showing well-defined features for Ag" centres.3637 responsible for this width, unfortunately it was not resolved. This probably means that a range of solvates containing both H,O and NH, ligands is present, each one capable of adding electrons. That ammonia ligands must be present is established by the remarkably low value for Aiso (lo9Ag) (ca. -451 G), corresponding to only ca. 64% spin density on silver. On annealing, features for Ag: became better defined, and again showed unusual width and an exceptionally low value for Aiso (- 200 G).These differences presumably arise because NH, ligands are retained in the dimer, delocalisation still being quite extensive. When the central signals, due mainly to 'NH, radicals,lG were lost on annealing, there was a concomitant growth in features assigned to an Agl* complex (fig. 5). As with pyridine, the perpendicular features were well resolved, showing hyperfine coupling to three or four 14N nuclei. Our preferred analysis is based on the presence of four ligands (fig. 5), but three equivalent 14N nuclei are equally probable. The possibility that only two NH, ligands were present can be ruled out. This complex, which is clearly a square-planar complex having a (- - - 4d$-Y2) electron configuration, is presumably formed by attack of 'NH, radicals on Agl complexes.Since these radicals are not likely to be efficient electron-acceptors, it is probable that these radicals add to the AgI, possibly remaining as ligands. Thus if, as expected,17 the major Agl complex is a linear complex, Ag(NH,)i, presumably with water molecules less strongly coordinated in the equatorial plane, ' NH, radicals would convert these into complexes having three equatorial ligands containing 14N nuclei rather than four. M. C. R. SYMONS A N D C. K. ALESBURY FORMAMIDE A N D DIMETHYLFORMAMIDE For formamide at 77 K, broad features assigned to Ago, Ag: and AglI centres were obtained. On annealing, the Agz features became better resolved and two different Ago centres were apparent.No 14N splitting was observed for these centres, but the features were broad. For [2H7]DMF, resolution was better, and several types of Ago centres were detected (fig. 6). Also a new species (a) was detected, which is discussed below. In addition A i B I 3250G FIG. 6.-First-derivative X-band e.s.r. spectra of AgC10, in [*H,]DMF after exposure to 6oCo y-rays at 77 K, showing features assigned to Ago (A and B) and Ag" centres, together with features (a) assigned to species having structure II.3638 SOLVATION OF SILVER IONS to weak unsolvated atom features, very narrow features (A) for a weakly solvated species, and intense, broad features for a strongly solvated species (B) were obtained. Species (A) grew at the expense of (B) on annealing, probably as a result of loss of primary solvent molecules.A most intriguing aspect of the spectra for B is the appearance of hyperfine splitting in the outer regions of all four features. Only 3-5 features were resolved, but high-resolution spectra reveal the presence of several further shoulders. Since the splitting of cu. 6 G is equal to that found for 14N coupling in the Ago(MeCN), complexes, it is tempting to assign this coupling to two or more 14N nuclei. This result is most unexpected since it is generally found that amides coordinate uiu oxygen. Indeed, amides are locally nearly planar at nitrogen; the nitrogen 2p-electrons, being part of a delocalised ;n-system, are not normally available for coordination to cations. We hope to study this possibility using infrared spectroscopy.BENZENE AND TOLUENE The unusually high solubility of silver salts in aromatic solvents arises because Ag+ ions form weak complexes with unsaturated systems. Indeed, well-defined com- plexes such as C,H,*AgC10,,18 are well known. This complex consists of -C,H,-Ag+-C,H,-Ag+- chains, each benzene ring being coordinated to two Ag+ ions lying above and below opposite bonds of the ring.19y20 We find that the isolated complexes are very radiation resistant, the only para- magnetic silver centre detectable after long exposure at 77 K being an Agz+ unit. However, as reported previously,21 frozen solutions do give fair yields of electron- excess centres on irradiation. The results of the present work, which was directed towards mixed-solvent systems, differ from those previously reported, for no clear reasons.We now find that solutions 1 32506 Solvent radicals Gain x 10 ( 107 Ago + ';go M I = +'h FIG. 7.-First-derivative X-band e.s.r. spectra of AgClO, in toluene after exposure to 6oCo prays at 77 K, showing features assigned to Ago centres and solvent radicals. Weak features for Agl are also seen.M. C. R. SYMONS A N D C. K. ALESBURY 3639 in dry benzene give, after irradiation, well-defined Agi+ centres with clear resolution of the four sub-components from species containing differing proportions of lo9Ag and lo7Ag. Dry toluene solutions gave an Ago centre, not previously detected in aromatic solvents (fig. 7). We suspected that the reason why different types of silver centres were formed in benzene under different conditions might be connected with phase separation, so we also studied the effect of ionizing radiation on the isolated complex (C,H, * AgC10,).In fact, the radical yield was very low, the only detectable silver species being Agi+, the MI = +E line showing the expected four features from different lo9Ag+ lo7Ag combinations. It seems probable that dilute solutions at low temperatures comprise aggregates containing three Ag+ ions and probably three benzene molecules, and these share one electron between them. Toluene gives good glasses on rapid freezing, which probably explains why monatomic centres can be formed. The high %-character for these previously unknown centres (ca. 94%) shows that delocalisation into the aromatic ring is minor.In other words the 5s(Ag) orbital is not greatly involved in the bonding for these complexes. When the silver(1)-toluene complex was isolated and irradiated, yields were again low, A&+ and A&+ being the dominating electron-excess centres. The relatively low A(lo9Ag) values for these complexes (A&+, - 163 G, corresponding to ca. 697: spin density on silver; A&+, - 122 G, also corresponding to ca. 69% on silver) suggests that the aromatic rings are still involved in these structures, and that delocalisation is facilitated relative to the monomer units. The data for the Agg+ complex are very close to those for the benzene complex, suggesting comparable structures. CORRELATION OF Ai,,(109Ag) WITH SOLVENT DONOR NUMBERS The donor number (DN) of a solvent is a widely used empirical measure of its ability to coordinate via a lone-pair of electrons.It is based upon the heat of formation of a single coordinate bond to SbC1,.22 We use it because such bonding is thought to be involved when Agl is solvated by the solvents used herein. It must be borne in mind that the correlation with Ai,,(109Ag) is only expected to hold for complexes having the same solvation number. Clearly, as the solvation number increases so the extent of electron delocalisation must increase. Hence only those species which appear to be primary species have been utilised in the correlation which is shown in fig. 8. That the correlation is reasonable for some 10 systems supports the idea that bonding into the outer vacant (for AgI) 5s and 5p orbitals dominates solvation for Ag' systems.It is interesting that data for silver nitrate in cyanomethane and cyanoethane lie close to the line, whereas those for silver perchlorate are well removed therefrom. This fits in with our suggestion that silver is coordinated to oxygen ligands from two nitrate ions.z3 In that case, the donor numbers for MeCN or EtCN are no longer appropriate. If we move these points horizontally onto the correlating line, we obtain a donor number for NO; of ca. 18, close to those for water or methanol. This seems to us to be entirely reasonable. The data for solutions in MeCN and EtCN suggest that these molecules coordinate to silver ions far more strongly than one would expect for these weakly basic solvents.In our view, this can best be explained in terms of back-donation of silver 4d (n) electrons into the otherwise vacant n* (x, y ) orbitals of the N-C group (the ' back-Chatt ' effectz4). Such bonding will reduce the silver-nitrogen bond length, thereby increasing a-overlap and delocalisation. It will also increase the positive charge on silver and the negative charge on the N-C group, thus also increasing the strength of the a-bonds.3640 450. SOLVATION OF SILVER IONS 0 EtCN MeOH MeTHFDMF DMSO HMPA I I 1 I I I I I c3 '= 0 ' 2 v 0 650 - 0 550 t ( AgClO,) 0 1 Me *, \ Triethylamine is not able to participate in such bonding, but pyridine can, and it is noteworthy that once again the A(logAg) value is far lower than predicted by the correlation. MIXED SOLVENTS WITH CYANOMETHANE MeCN + Me,CO The most remarkable result, as with the mixed MeCN + H,O and MeCN + MeOH system^,^ was that the AgO(MeCN), and Ag"(MeCN), complexes persisted up to the 0.5 mole fraction region before resolution into the nine (14N) hyperfine components was lost.In the 0.5-0.87 mole fraction region these features lost all resolution, and features for Agz centres grew in. No clear evidence for specific mixed solvates was obtained, results for the Agl* centre suggesting that damage to the Ag*(MeCN), centres may have been preferred over other mixed complexes since features for the tetracyanomethane derivative were still detectable in the 0.8 mole fraction region. Ultimately features characteristic of solutions in pure acetone were obtained in the > 0.9 mole fraction region.In most systems, on annealing, features for Agi, Agi+ and Ag:+ were detected. Other species with relatively low hyperfine coupling to lo9Ag, formed in some of these mixed-solvent systems, are discussed below. MeCN + DMSO We expected, from the correlation in fig. 8, that DMSO would displace MeCN far more readily than Me,CO because of its high donor number. This was indeed the case. Beyond ca. 0.1 mole fraction (DMSO) the 14N hyperfine features began to broaden with loss of resolution. This probably reflects absorption from AgO(DMSO),(MeCN),, AgO(DMSO),(MeCN,) and other mixed solvates, the net effect of a range of coupling constants and lines in different positions being simply to broaden out the features, as observed.M. C. R. SYMONS A N D C.K. ALESBURY 364 1 However, as with pure DMSO, these mixtures in the region > 0.8 mole fraction gave mainly Agi centres, which, on annealing, readily gave Ag;+ centres. Also, well-defined AglI centres characteristic of DMSO binding were obtained. Thus, in this mixed-solvent system there was no evidence for preferential solvation by MeCN. MeCN + MTHF In this case, preferential solvation by MeCN resembled that for acetone, the Ago(MeCN), and AgIl(MeCN), centres dominating the e.s.r. spectra up to ca. 0.6 mole fraction MTHF. From then on there was a gradual loss of resolution and increase in A(lo9Ag) for the Ago centres. Species B, detected in pure MTHF, grew in the 0.8 mole fraction region, prior to the appearance of species A. Features for AglI centres gradually diminished in the MTHF-rich region, being finally lost for the pure-solvent system.OTHER MIXED-SOLVENT SYSTEMS MTHF+ Et,N Again, as predicted from the correlation of fig. 8 and the high donor number of Et,N, the species found in pure Et,N dominated to beyond the 0.95 mole fraction MTHF range, indicating extreme preferential solvation by Et,N. C,H,+MTHF AND MeNO, On addition of 0.05 mole fraction MTHF, the Agg+ signals characteristic of benzene clusters were completely lost, and features characteristic of Ag-MTHF complexes were detected. Addition of MeNO, with a lower donor number required far greater concentrations to change the spectrum from that of pure benzene. These systems were not studied extensively, but it seems clear that the donor number is a fair measure of coordinating ability with the marked exceptions of cyanoalkanes and pyridines, and that the aromatic complexes are in no way exceptional in this regard, THE AglI CENTRES All the data obtained for AglI centres (table 2) are characteristic of d 9 complexes with a (- - - 4diz-,2) onf figuration.^^ However, in several cases, marked changes occurred on annealing involving, mainly, increases in gI1 and the magnitude of A II (lo9Ag).We suggest that these changes are the consequences of the complexes relaxing from a near-tetrahedral configuration fairly close to that of the parent Agl complexes, towards the square-planar structures that are favoured by a di2-y2 configuration. After allowing for the difference in spin-orbit coupling constants and nuclear magnetic moments, the changes are quite similar to those predicted by our correlation for copper(r1) (3d$-,2) complexes.26 OTHER PARAMAGNETIC SILVER CENTRES The cluster centres, Agi, Agz+ and Ag;+ centres have been discussed for each solvent system.A wide range of such complexes have now been studied, and the isotropic coupling constants to lo9Ag change quite markedly from one system to another. In many cases these trends resemble the trends shown in fig. 8 for Ago centres. This implies that coordination of solvent molecules is retained, as has been clearly established for the Agi+ centre detected in irradiated silver imidazole per~hlorate.~~ Other aspects of the data for these remarkable complexes will be considered later in connection with some studies on conduction-electron centres.A species previously described, probably having the limiting structure I,, was formed in many of the MeCN systems. This species has an isotropic coupling to lo9Ag3642 SOLVATION OF SILVER IONS of ca. 180 G. Yet another centre with a relatively small but isotropic coupling to lo9Ag and g-values close to 2.00 has now been detected. This species, with A(lo9Ag) = - 100 G, was formed in DMF solutions, as shown in fig. 6. There is some indication of 14N hyperfine splitting on these features, with A(14N) z 17 G, but this was never well resolved. This centre must differ from the Ago centres formed in this medium, so that simple coordination to nitrogen or oxygen is ruled out. Possibly some sort of n-bonding such as that depicted in I1 is involved.A similar species was formed in the 0.5-0.9 mole fraction (acetone) range for mixed MeCN + Me,CO systems, after annealing. This species, having IAi,,(logAg)l - 82 G UY I V (fig. 9, is thought to be a complex between Ago and acetone, since it was favoured in the high mole fraction (acetone) region, although it was not detected in pure acetone solutions. We think that the extra splitting clearly resolved in the high-field feature is due to g-value variation rather than 14N coupling but, unfortunately, we could not test this suggestion by studying the spectrum at Q-band frequencies because of low instrument sensitivity. A possible structure for this complex is shown by 111. 1 32506 FIG. 9.-First-derivative X-band e.s.r. spectra of AgCIO, in CD,CN + (CD,),CO (mole fraction 0.625) exposure to 6oCo y-rays at 77 K and annealing, showing features assigned to a species thought to complex between Ago and acetone (In).after be aM. C. R . SYMONS A N D C. K. ALESBURY 3643 Alternatively, these complexes may be formed from D,CCO(CD,) radicals [or D,cN(CD,)COD radicals for DMF] having structure IV. A species with such a structure (Ag-CH,OH) was detected in methanolic solutions, having Ai,O(lOgAg) e 130 G.2 CONCLUSIONS We conclude that most solvents bind to Ag+ via a-donation into 5s+ 5p manifold. On electron addition at low temperatures the primary solvent molecules are retained, the unpaired electron being extensively delocalised onto the ligands. The extent of delocalisation, measured by the reduction in Ai,O(lOgAg), correlates quite well with the donor number of the solvent, with the notable exceptions of cyanoalkane solvents and pyridine.In these cases back-bonding by 4d electrons is thought to be important. On annealing, solvent molecules are lost from the Ago centres, and extensive clustering usually occurs. The Ag" centres show evidence for relaxation from a near-tetrahedral to a square-planar configuration on annealing. We thank the S.E.R.C. and Kodak (Harrow) Ltd fgr a CASE Studentship (to C.K.A.) and Dr G. Farnell for helpful discussions. D. R. Brown, G. W. Eastland and M. C. R. Symons, Chem. Phys. Lett., 1979, 61, 92. C. K. Alesbury and M. C. R. Symons, J. Chem. SOC., Faraday Trans. 1, 1980, 76, 244. B. G. Cox, R. Natarajan and W. E. Waghorne, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 86. A. P. Zuur and W. L. Groeneveld, Recueil, 1967, 86, 1089. A. S. W. Li and L. Kevan, J. Phys. Chem., 1981, 85, 2557. R. J. England and M. C. R. Symons, J. Chem. SOC. A, 1970, 1326. M. C . R. Symons and G. W. Eastland, unpublished. R. S. Eachus and M. C. R. Symons, J. Chem. SOC. A, 1970, 1329; 1970, 3080. * L. Shields and M. C . R. Symons, Mol. Phys., 1966, 11, 57. lo M. C. R. Symons, f. Chem. Phys., 1978, 69, 3443. l 1 L. Kevan, H. Hase and K. Kawabata, J. Chem. Phys., 1977, 66, 3834. l 2 M. C. R. Symons, Pure Appl. Chem., 1981, 53, 223. l 3 Unpublished results. l4 I. M. Strauss and M. C. R. Symons, J. Chem. SOC., Faraday Trans. I , 1978, 74, 2146. l5 H. J. Bower, J. A. McRae and M. C. R. Symons, J. Chem. SOC. A, 1968, 2696. l6 K. V. S. Rao and M. C. R. Symons, f. Chem. SOC. A, 1971, 2163. l 7 F. A. Cotton and G. Wilkinson, Adzianced Inorganic Chemistry (Wiley-Interscience, New York, 3rd edn, 1972). A. E. Hill, J. Am. Chem. SOC., 1922, 44, 1163. l9 R. E. Rundle and J. H. Goring, J. Am. Chem. SOC., 1950, 72, 5337. 2o H. G. Smith and R. E. Rundle, J. Am. Chem. SOC., 1958, 80, 5075. 21 C. E. Forbes and M. C. R. Symons, Mol. Phys., 1974, 27,467. 22 V. Guttmann, Donor-Acceptor Approach to Molecular Interactions (Plenum Press, New York, 1978). 23 D. R. Brown, T. J. V. Findlay and M. C. R. Symons, f. Chem. SOC., Faraday Trans. I , 1976,72, 1792. 24 J. Chatt, Ric. Sci. Suppl., 1958, 23, 130. 25 M. C. R. Symons, Chemical and Biochemical Aspects of Electron Spin Resonance Spectroscopy (Van Nostrand Reinhold, London, 1978). 26 M. C. R. Symons, D. X. West and J. G. Wilkinson, J. Chem. SOC., Dalton Trans., 1975, 1696. 27 G. W. Eastland, M. A. Mazid, D. R. Russell and M. C. R. Symons, J . Chem. SOC., Dalton Trans., 1980. 1682. (PAPER 2/568)

ISSN:0300-9599    

DOI:10.1039/F19827803629

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1982, 78, 3645-3657 Pulse Radiolysis of Solutions of trans-Stilbene Radical-anions and Ion-pairs in Tetrahydrofuran BY JOHN R. LANGAN AND G. ARTHUR SALMON* University of Leeds, Cookridge Radiation Research Centre, Cookridge Hospital, Leeds LS16 6QB Receiced 5th April, 1982 On pulse radiolysis of solutions of trans-stilbene (t-St) in THF the radical-anion of t-St is formed by reaction of e; with t-St [reaction (S)] with k , = (1.16 k0.03) x 10" dm3 mol-' s-l. The transient absorption spectrum observed with I,,, at 500 and 720 nm is attributed to the unassociated St--. The subsequent decay of the radical-anion is accounted for by reaction with the counter-cation of THF formed on radiolysis and with radiolytically generated radicals; rate constants for these processes are estimated.Addition of sodium tetrahydridoaluminate (NAH) results in the radical-anion being associated with Na+ as a contact ion-pair and a shift of Amax to 490 nm. In the presence of the lithium salt the absorption spectrum of the radical-anion reverts to 500 nm and this is interpreted in terms of formation of solvent-separated ion-pairs. On pulse radiolysis of solutions containing NAH the main reaction forming St*- is that of (Na+, e;) ion-pairs with t-St [reaction (12)] with k , , = (1 -29 k0.04) x 10'" dm3 mol-I s-'. In addition there is a delayed formation of St*- over a period of microseconds and a possible mechanism for this is considered. The presence of tetrahydridoaluminate salts also greatly enhances the stability of St*- and at high doses per pulse little decay was observed over 700 p s .Reasons for the enhanced stability of St*- are considered. The variation of G(St'-) with "AH] was studied and was found to attain a plateau value of 2.0 at the higher concentrations. Pulse radiolysis of NAH/THF solutions containing crown ether yielded an absorption spectrum closely resembling that of e;, but kinetic studies on the solutions containing trans-stilbene indicated that the species is an ion-pair between e; and Na+ in which the sodium is complexed with the crown ether, i.e. (Na+C, e;). This species reacts with trans-stilbene with a rate constant of (3.86k0.07) x 10'" dm3 mol-l s-l. Pulse radiolysis of solutions of aromatic compounds (Ar) in tetrahydrofuran (THF) is known to result in the formation of the radical-anion of the aromatic compound by reactions (1) and (2)f-3 THF- THF(H)++e; (1) (2) e;+Ar -+ Ar*-.However, the lifetime of the radical-anions generated in this manner is generally short due to the occurrence of the fast neutralization reaction (3) between the radical-anion and the counter-ion, which is also formed in reaction (1)3 (3) Ar' - + THF(H)+ + Ar(H) + THF. The objective of the present study was to find conditions under which the lifetime of Ar*- is sufficiently extended that it would be possible to study the relatively slow reactions of the radical-anions of polymerizable aromatic olefins which are likely to be of importance in the early stages of the anionic polymerization of these compounds. trans-Stilbene (t-St) was chosen as the solute for this study since its radical-anion, which can be prepared in THF by reduction with metallic sodium, is stable and its absorption spectrum has been well ~haracterised.~* Previously studies1> have indicated that the tetrahydndoaluminate salts of sodium 118 3645 F A R 13646 PULSE RADIOLYSIS OF l'TanS-STILBENE or lithium are able to scavenge the positive ion formed on radiolysis of THF [reaction (4)], thereby extending the lifetime of any transient reduced species: (4) Pulse-radiolysis6-s and flash-photolysisg studies with solutions of alkali-metal salts in THF have demonstrated the formation of alkali-metal-cation-electron ion-pairs, (M+, e;), and alkali-metal anions, M-, which are therefore expected to be inter- mediates in the formation of radical-anions in the presence of tetrahydridolaluminate salts.THF(H)+ + A1H; -+ THF + AIH, or THF + AlH, + H,. EXPERIMENTAL The techniques of pulse radiolysis employed in this laboratory have been described else- where.1°-13 Chance-Pilkington filters were placed between the analysing light source and the sample to prevent photolysis of the sample in the cell, which was of 1 cm path length throughout the study. The methods of purification of THF6 (Hopkin and Williams, AnalaR) and trans-stilbenelO (Kodak) have been described. Dicyclohexyl- 18-crown-6 ether (Lancaster Synthesis, PCR Ltd) was used as supplied. The tetrahydridoaluminate salts (Alfa, 95 %) were used as supplied except when used in conjunction with the crown ether, when they were recrystallised from THF.Solutions of the salts were filtered under vacuum before preparing the samples for radiolysis. All solutions were degassed by three freeze-pumpthaw cycles. RESULTS SOLUTIONS OF tYaTZS-STILBENE IN T H F The end-of-pulse absorption spectrum observed on pulse radiolysis of solutions of trans-stilbene in THF, which is shown in fig. 1, agrees with that of the radical anion h/nm FIG. 1 .-(a) End-of-pulse absorption spectrum induced in a solution of lo-* mol dm-3 trans-stilbene in THF by a 50 ns, 15 Gy pulse. (b) Absorption spectrum 1 ps after irradiation of a mol dm-3 solution of trans-stilbene in NAH/THF with a 0.6 ps, 40 Gy pulse. of trans-stilbene.*q5 Using 25 ns, 10 Gy pulses the yield of radical-anion, assuming ~ 5 0 0 = 6.1 x lo4 dm3 mol-1 crnl * increased from G = 0.22 to 0.70 as the concentration of trans-stilbene was varied from 1.5 x mol dm-3.This dependence of G(St'-) on [t-St] was not studied in detail, but results from two effects, namely (i) at low concentrations of trans-stilbene the solute is unable to scavenge all the escaped electrons at the dose used and (ii) at high concentrations of solute additional spur electrons are scavenged. Two small absorption peaks at 360 and 390 nm were also toJ. R. LANGAN A N D G. A. SALMON 3647 observed which are probably due to radicals of trans-stilbene formed in the geminate neutralization process. All the absorptions decayed completely within several microseconds. The rate of reaction of e; with trans-stilbene [reaction ( 5 ) ] was measured by monitoring the formation of the stilbene anion at 500 nm e;+St + St*-.( 5 ) For trans-stilbene concentrations from 1.5 x mol dm-3 the formation of St*- following irradiation with 10 Gy pulses of 25 ns duration was found to obey first-order kinetics. The observed first-order rate constants so obtained were proportional to the concentration of trans-stilbene over the concentration range (2-30) x mol dm-3 and yielded k, = (1.16kO.3) x loll dm3 mol-1 s-l. The decay of the anions would be expected to follow second-order kinetics if they were distributed homogeneously through the medium and reacted only with cations. In fact the decay, monitored at 500 nm, gave second-order plots, but the slopes showed some variation with dose and trans-stilbene concentration and the extent of the decay which gave a linear fit varied from as short as 1.5 up to at least 2.5 half-lives.Comparable results were obtained when monitoring St*- at 710 nm. The deviation from a strict second-order rate law is ascribed, as elsewhere,1.3 to reaction of the radical-anion with radiolytically produced radicals. The observed values of the slopes? of the second-order plots varied from 3.9 x lo6 to 9.2 x lo6 s-l, with a mean value of 6.0 x lo6 s-l. The reaction scheme proposed to account for the formation and decay of St*- in the pulse radiolysis of solutions of trans-stilbene in THF is presented in table 1. to 3 x TABLE REACTIONS AND RATE CONSTANTS IN THE PULSE RADIOLYSIS OF ~~uw-STILBENE IN THF no. reaction rate constant /dm3 mol-l s-' ref. 5 e;+St+St'- 1.16 x 1011 this worka 7 e; + R -+ product 3.0 x 1 O 1 O 8 9 St'-+THF(H)+ 3 product 4.2 x 1011 this workb 10 St'-+R -+product 3.0 x 109 this workb 6 e; + THF(H)+ -+ product 2.0 x 10'2 2 8 R+R+R, 5 .o ~ 109 8 a Direct measurement; computer fit. Using these data and assuming G(e;) = 0.39, G(R) = 6 and &(St'-) = 6.1 x lo4 dm3 mol-l cm-l, a computer program has been written to calculate the concentration of each species formed as a result of the radiolysis. The program incorporates a least-squares refinement routine in two parameters for the evaluation of unknown rate constants. Reproductions of experimentally observed absorbance changes with the computed points superimposed upon them are given in fig. 2. The good agreement between experiment and calculation suggests that the reaction t k / d is the slope of the plot of A-' against t, where k is the rate constant in units of dm3 rno1-I s-1, E is the molar decadic extinction coefficient of the absorbing species in units of dm3 mo1-I cm-l and 1 is the optical path length of the cell in cm.118-23648 PULSE RADIOLYSIS OF tranS-STILBENE I" (v c 9 8 0 0 0 0 c OD 0 (D 0 U 0 (v 0 0 0 0 ameqiosqeJ. R. LANGAN A N D G. A. SALMON 3649 mechanism in table 1 is essentially correct, although it should be noted that the entity denoted by R may in fact represent more than one radical. The decay of the small absorption peak at 390 nm, which is believed to be due to stilbene radicals, was found to fit a second-order plot, although again there was a small variation of the slopes with dose, which suggests that reactions with stilbene radical-anion and possibly other radicals were occurring.The mean value of the slopes of the second-order plots was found to be 1.7 x lo6 s-l. SOLUTIONS WITH ADDED TETRAHYDRIDOALUMINATE SALTS The pulse radiolysis of a solution of trans-stilbene in THF saturated with sodium tetrahydridoaluminate (NAH/THF) resulted in the spectrum shown in fig. 1. Comparison with the spectrum observed in neat THF reveals a hypsochromic shift in the absorption peak of the radical-anion to 490 nm. The broad band around 710 nm was also observed. When doses per pulse in excess of 40 Gy are used the radical-anion is long lived and very little decay was observed over a period of 7OOp.s. Typical oscilloscope traces recorded during this experiment are given in fig.3 (a) and (b). The L t,,,,,,,,, FIG. 3.-Oscilloscope traces recorded in the pulse radiolysis of trans-stilbene in NAH/THF ordinate : % absorption pulse size dose/Gy I/nm per division (a) 0 . 6 , ~ 40 490 17.9 (b) 0 . 6 ~ s 40 490 17.0 ( d ) 50x1s 10 890 1.4 (e) 0 . 2 p 16 490 7.1 (f) 0.2 PS 16 490 7.1 (c) 50 ns 10 490 4.1 abscissa : ps per division 100 1 0.1 0.1 100 13650 PULSE RADIOLYSIS OF frUnS-STILBENE stability of the anion is indicative of the efficiency with which the AlH, ions remove the counter-ions formed during the pulse [reaction (4)]. The absence of absorptions below 400 nm due to radicals evident in the absence of NAH is probably also a consequence of this reaction. A similar spectrum was recorded when the sodium salt was replaced by lithium tetrahydridoaluminate (LAH), except that A, reverted to 500 nm.From these results it is deduced that in these systems St.- exists as a contact ion-pair (Na+, St.-) when associated with Na+, but as solvent-separated ion-pairs (Li+/ /St' -) when associated with Li+. For comparison purposes similar solutions were subjected to y-radiolysis and the spectra recorded are shown in fig. 4. It was noted that for the solutions containing 3 X/nm FIG. 4.-Radiation-induced absorption spectrum obtained by (a) y-irradiation of 1 0-2 mol dm-3 trans- stilbene in NAH/THF, (b) pulse radiolysis of mol dm-3 rrans-stilbene in NAH/THF, and (c) y-irradiation of 6.3 x lop3 mol dm-3 trans-stilbene in LAH/THF. Each spectrum is normalised to A,,,. NAH, Amax underwent a bathochromic shift with irradiation time to a limiting value of 508 nm.This shift is probably caused by the formation of an absorbing species which gives rise to the shoulder around 570 nm. The size of this shoulder is much smaller in solutions containing LAH or in solutions containing NAH which were irradiated with a pulsed electron beam. High dose rates thus effect a marked reduction in the formation of the species absorbing at around 570 nm. The nature of this species has not been investigated. However, in the y-irradiation of pyrene in LAH/THF1 the absorption of the anion peak (493 nm) was not seen, but instead one at 530 nm was observed which was attributed to the formation of a complex, probably Py*AlH,*THF-. In the pulse radiolysis of that system a post-pulse growth of absorption at 530 nm was noted as partial decay of the anion occurred.It may be that in the radiolysis of trans-stilbene solutions the absorption around 570 nm is due to the complex St*AlH,*THF-. However, no growth related to the absorption at this wavelength was observed when recording the spectrum (fig. 1) by pulse radiolysis. The dependence of the complex formation on dose rate may be explained if AlH, is removed by reaction (1 l), which would be more effective at the high dose rates involved in pulse radiolysis: 2AlH, + Al,H,. (1 1)J. R. LANGAN AND G. A. SALMON 365 1 Using a trans-stilbene concentration of 5 x lop3 mol dm-3 the yield of the radical- anion was measured as a function of NAH concentration using pulse radiolysis and the results are given in fig.5. At the higher salt concentrations a maximum value of G(St'-) = 2.0 _+ 0.1 was reached. This is in accordance with previous work where G(Na+, e;) = 2 was recorded.8 0 0.2 0.4 0 . 6 0 . 8 1 . [NAHl/mol dm-3 FIG. 5.-Yield of trans-stilbene radical-anion as a function of NAH concentration, [t-St] = 5 x mol dm-3. The yield of radical-anions in NAH/THF was shown to be independent of initial trans-stilbene concentration by both gamma and pulse radiolysis. The results are depicted in fig. 6, which also includes some values obtained using LAH/THF. The higher values of G found in LAH/THF are ascribed to the greater solubility and dissociation constant of the lithium salt in THF and are in agreement with other results.1t8 The difference between the values found in gamma and pulse radiolysis is probably due to the higher dose rate of the latter technique which enhances the second-order removal of radicals.For solutions with [t-St] < 5 x lop3 mol dm-3 and using 50 ns pulses and sweep speeds of 100 ns per division it proved possible to observe the broad absorption band of (Na+, e;). Monitoring the decay of this absorption at 880 nm provided a convenient means of following the kinetics of the reaction between (Na+, e;) and trans-stilbene, reaction (1 2) (Na+, e;) + t-St + (Na+, St *-). (12) The decay of (Na+, e;) was matched by the concomitant formation of stilbene radical-anions at 490 nm [see fig. 3(c) and (41. For trans-stilbene concentrations in the range 2.5 x lop4 to 2.2 x lop3 mol dm-3 the decay of absorption at 880 nm was found to obey a first-order rate law and a plot of the observed rate constants against trans-stilbene concentration yielded k,, = (1.29 0.04) x 1O1O dm3 mol-1 s-l.The growth of absorption at 490 nm also obeyed first-order kinetics and yielded3652 PULSE RADIOLYSIS OF t!YLEFZS-STILBENE 3*0* 0 -5 -4 -3 -2 log [trans-stilbene I FIG. 6.-Yields of stilbene radical-anion obtained in gamma and pulse radiolysis of trans-stilbene solutions : 0, pulse radiolysis of NAH/THF solution; x , y-irradiation of NAH/THF solution; 0, y-irradiation of LAH/THF solution; A, pulse radiolysis of LAH/THF solution. k,, = (1.30 & 0.08) x 1O1O dm3 mol-l s-l, in good agreement with the value obtained from the decay of (Na+, e;). As was noted [fig. 3(a) and (b)], there was very little decay of the radical-anion following the pulse.However, this seems to result from the high dose used. On a microsecond timescale, using a 10 Gy pulse, two distinct phases are apparent in the behaviour of the absorption after the pulse. Initially, over ca. 5 ps, there is an increase in the absorption followed by a decay lasting for several hundred microseconds. These features are illustrated in fig. 3(e) and (f). Three spectra were recorded, at the end-of-pulse, at 5 ps and at 600ps after the pulse, respectively. The maximum absorbance for all three occurred at 490 nm in NAH/THF and 500 nm in LAH/THF. At the longer sweep speeds the decay at 570 nm appeared to be offset by the formation of the complex referred to previously. The growth in absorbance at Amax after the pulse gave linear plots when fitted to a first-order rate equation and some results are given in table 2.The subsequent decay obeyed neither first- nor second-order kinetics and was greatly influenced by the dose. TABLE 2.-RATE CONSTANTS FOR THE POST-PULSE GROWTH OF ABSORBANCE [ trans-stilben-el solvent /mol dm-3 dose/Gy k'/s-' % slow growth ~ ~~ ~~ ~~ NAH/THF 1.13 x 15 4.07 x 107 24 6.02 x 10-3 7 8.63 x 107 15 LAH/THF 9:23 x 10-3 12 7.8ox 107 25 8.85 x 10-3 12 6.66 x lo7 __J. R. LANGAN AND G. A. SALMON 3653 Similar observations have been made elsewherel in the pulse radiolysis of pyrene in LAH/THF. The kinetic evidence, taken together with the fact that at high dose rates neither effect is observed, strongly suggests that the reactions involve radicals.Pulse-radiolysis studies of amine solvents containing alkali cations have described post-pulse growths of absorption lasting several microsecond~.'~ Further proof of the stability of the ion-pair formed in the pulse irradiated solutions was provided by e.s.r. The width of the signal and the line splitting were consistent with the published In addition the solutions were able to initiate the polymerization of butadiene. The resulting polymer was found not to have incorporated stilbene showing that initiation was by electron transfer rather than addition. EFFECT OF ADDED CROWN ETHER Addition of dicyclohexyl-l8-crown-6 (DCH) to a solution of NAH in THF is expected to result in the complexation of the sodium cations and hence in the in- creased dissociation of the salt.Pulse radiolysis of a solution containing NAH (6.1 1 x mol dm-3) and DCH (1.03 x lo-* mol dmP3) yielded a spectrum which was similar to that of the solvated electron2T8 with GE = 2.5 x lo4 at 2000 nm.? The absorption band of (Na+, e;) with 3,,,, = 890 nm which is observed in NAH/THF solutions was absent. For kinetic reasons (see below) it is believed that the spectrum resembling that of e; is not due to e;, but is the spectrum of (Na+ C, e;), i.e. an ion-pair between e; and the sodium cation complexed with crown ether. The absorption decayed completely over a period of ca. 10 p s [see fig. 7 (a)] by first-order kinetics with an observed rate constant of 2.51 x lo5 s-l. This value, measured at a dose of 43 Gy, is similar to that for e; at the same dose.8 Pulse radiolysis of a solution of trans-stilbene in NAH/DCH/THF is expected to yield the spectrum of the solvent separated ion-pair of stilbene anion, and this was found to be so.To ensure that the appearance of LmaX at 500 nm was not merely due to the effect of a lower salt concentration a spectrum was also recorded using only NAH at a similar concentration. The absorption maximum was then observed to be at 485 nm. The behaviour of St*- in this system followed the same pattern as in the NAH/THF system. The post-pulse changes were still present [see fig. 7(b)-(d)] and again were greatly reduced by the use of large doses per pulse. The decay observed at lower dose per pulse [fig. 7 ( d ) ] was found to fit a second-order plot with k/E = 1 x lo5 cm s-1 under the conditions used.The kinetics of formation of St*- were followed using the decay of the species at 2000 nm and the growth in absorbance at 500 nm. For concentrations of trans-stilbene from 5 x mol dm-3 the decay of absorption at 2000 nm and the growth at 500nm obeyed first-order rate laws and the observed rate constants increased linearly with trans-stilbene concentration, from which we deduce k,, = (3.86k0.07) x lo1" dm3 mol-1 s-l to 2.2 x (Na+C, e;) + t-St -+ (Na+C, St * -). (13) This rate constant is significantly different from that for the reaction between e; and trans-stilbene. Thus, although the spectrum of e; is not modified by the presence of Na+C, this latter ion is exerting a significant influence on the reactivity of e; and it seems appropriate to ascribe this to the formation of the ion-pair (Na+C, e;).The effect of NAH concentration of the yield of St*- is given in fig. 8. This shows t Gc is the product of the radiation chemical yield in units of molecules (100 eV)-' and the molar decadic extinction coefticient in units of dm3 mol-1 cm-'.3654 PULSE RADIOLYSIS OF tTaYtS-STILBENE la I l CI FIG. 7.-Oscilloscope traces recorded in the pulse radiolysis of NAH/DCH/THF [(b)-(d) with ca. mol dmV3 trans-stilbene added]. ordinate: % absorption abscissa : pulse-length/ps dose/Gy i/nm per division ps per division (a) 0.2 43 2100 2.3 1 (b) 0.2 35 500 7.2 1 (c) 0.6 140 500 17.8 100 (d) 0.2 35 500 7.3 100 2 .c h I c-' F", u 1 . 0 I 1 1 I I 1 2 3 4 ! [ NAHJ / 1 O-* mol dm-3 FIG. 8.-Yield of stilbene radical-anion as a function of NAH concentration in NAH/DCH/THF; 0.2 ,us pulses, [t-st] = 5 x lop3 mol drnp3.J .R. LANGAN A N D G. A. SALMON 3655 that G(St*-) = 2, which corresponds to the maximum value in NAH/THF solutions, is reached at a much lower salt concentration than in the absence of crown ether. This implies that the greater dissociation of the salt has enhanced the scavenging of the counter-cation [THF(H)+] by the AlH, ions. Using the dissociation constant of NAH in THF,8 the concentration of AlH, in the NAH/THF system was estimated to be 3 x mol dm-3. If it is assumed that the salt is completely dissociated in the presence of an excess of crown ether then the evidence suggests that undissociated NAH is also able to scavenge the positive ion of THF.In an experiment where sodium tetrahydridoborate was substituted for NAH, G(St'-) was found to be 1.65. However, no stabilization of St*- resulted and it seems that, as in the case of sodium tetraphenylboron,8 the reaction of BH, with the cation of THF gives a radical which can react with St*-. THERMAL GENERATION OF ANIONS During the course of this work it was noticed that when solutions of trans-stilbene in THF containing the tetrahydridoaluminate salts were left to stand at room temperature a red colour developed. Similar behaviour in pyrene solutions has been reported16 where the presence of the anion was ascertained by e.s.r. A solution of trans-stilbene (ca. l OP4 mol dmP3) in NAH/THF was allowed to stand in the dark until it had developed red colour, approximately one week, after which it was found to exhibit the absorption and e.s.r.spectra of the stilbene anion. The rate of coloration quickened when the solutions contained crown ether. DISCUSSION The spectrum of the radical-anion of stilbene recorded in the pulse radiolysis of solutions of trans-stilbene in neat THF may readily be assigned to that of the free ion. In the presence of tetrahydridoaluminate salts the anion will probably be associated with the metal cation. In the case of the sodium salt the effect of this can be seen in the occurrence of a hypsochromic shift of Amax. The absence of such a shift in solutions containing Li+ suggests that solvent-separated ion-pairs are formed in this system. The difference in the behaviour of the two cations is explicable in terms of the larger solvation energy of the small Li+ ion.For larger cations, interaction with an anion as a contact ion-pair is more favourable than interaction with solvent molecules. The most important property of a solvent in promoting the formation of solvent separated ion-pairs is its ability to coordinate alkali-metal cations, which is measured by its donicity rather than by bulk properties such as permitti~ity.'~~ The solvation of ions and ion-pairs and its effects on anionic polymerization have been described comprehensively elsew here. Computer simulations of the pulse radiolysis of solutions of trans-stilbene in NAH/THF showed that lO-l5% of the electrons were scavenged directly by the hydrocarbon. Any reaction of St*- formed in this way with Na+ in the solution to give ion-pairs would be expected to proceed with a rate constant 2 loll dm3 mol-1 s-l and hence be unobservable under the conditions of the experiment.The absorption spectrum recorded in NAH/THF will be due to the species in the following equilibria (Na+, St*-)tighte(Na+, St'-)loose e N a + + S t ' - . The small difference in A, seen here compared with that of the anion prepared by sodium-metal reduction4 may be due to the presence of some free ions or loose ion- pairs. It is known that radical-anions of hydrocarbons in THF at concentrations < ca. mol dmP3 may exist as free ions17 l8 and the concentration of anions3656 PULSE RADIOLYSIS OF trCInS-STILRENE produced by pulse radiolysis will not exceed this figure. Solvent-separated pairs are not distinguishable optically from the free ions.However, in flash-photolysis studies of extremely dilute solutions of sodium pyrenide in THF, uiz. (2-4) x mol dm-3, the dissociation of ion-pairs into free ions was suppressed by the presence of excess sodium tetraphenylboron.18 Therefore, it seems likely that the concentration of free ions in our solutions should be negligible. The reduced interaction of the solvated electron with the Na+ cation when it is complexed by crown ether is indicated by the effect of crown ether on the transient absorption spectrum and, as may be expected from this result, the spectrum of the stilbene anion in a solution containing crown ether is that of the loose ion-pair. The rate constants for the formation of radical-anions and ion-pairs of trans-stilbene are given in table 3 and are similar to those reported for other aromatic hydrocarbons.2* 7 7 l8 The rate of reaction (1 3) is consistent with the spectral data, which shows that the interaction between the sodium cation and the solvated electron is reduced by the presence of crown ether.TABLE 3.-RATE CONSTANTS FOR THE FORMATION OF STILBENE RADICAL-ANIONS reacting species rate constant/lO1° dm3 mol-’ s-’ e, 1 1.6 0.03 (Na+C, e;) 3.86 0.07 “a+, e,) 1.3 k 0.08 The variation in the observed yield of stilbene radical-anions with the concentration of trans-stilbene in THF followed the pattern established in previous pulse-radiolysis studies.l In solutions containing tetrahydridoaluminate salts the total reducing power depends on the concentration of the salt and consequently this factor governs the yield of radical-anions formed in the radiolysis.The salt increased the yield of ion-pairs by scavenging THF cations within the spur, thus making electrons available in the form of (Na+, e;). The subsequent stability of the radical-anion is also due to the AlH; ion scavenging the counter-ion as described earlier. Complexation of the Na+ cation by crown ether led to the expected increase in scavenging of the counter-ions by the AlH; ions. It seems from comparison of yields of St-- in NAH/THF with those in NAH/DCH/THF that the undissociated salt also acts as a scavenging agent. Alternatively, it may be that in NAH/DCH/THF some of the Na+ cations and AlH, anions are associated in crown-separated pairs whose reactivity lies between that of the free ions and the undissociated salt.Computer simulations of the pulse radiolysis of trans-stilbene in NAH/THF indicated that over 99.904 of the electrons reacted with trans-stilbene as either e; or (Na+, e;). This is a reflection of the extreme rapidity of these reactions, which thus prevent significant decay of the negative species by other possible pathways. The values of GE for (Na+, St.-) found in this work (see fig. 5 ) are consistent with the reported values of G(Na+, e;)* and the extinction coefficient for the radical-anion of ~tilbene.~ Hence the value G(anion) = 2.0 0.1 may be assumed for any hydrocarbon of similar electron affinity to trans-stilbene whose reaction with (Na+, e;) proceeds at a similar rate.Thus, the absorption spectra and extinction coefficients of radical-anions which are not sufficiently long lived to be observed when prepared by reduction by alkali metals may be determined and the results of such work will be reported in a subsequent paper.J . R . L A N G A N A N D G. A. SALMON 3657 As indicated earlier, the fast formation of St*- in solution containing the tetra- hydridoaluminate salts can be attributed to reactions ( 5 ) and (1 2). However, the slow formation of St*-- observed on microsecond timescales is more difficult to interpret. A similar delayed formation of e; has been observed in the pulse radiolysis of amines containing bases.14 Since the effect is reduced when larger doses per pulse are used it is tempting to speculate that radiolytically produced radicals are involved in the formation of e; by reaction with A1H; ions [reaction (14)] +AIH, + (THF-) -+ e; + AIH, (14) 3T-T 0 but further experimentation is required to establish the mechanism of this slow formation of St*-.A further difficulty concerns the effect of dose on the partial decay of St. in the NAH/THF solutions. Again theeffect ofdose suggests the involvement of radiolytically produced radicals [reaction (1 5)] R* +St*- -+ St+product. (1 5 ) However, it is thought that several cationic species are formed on the radiolysis of THFl9 and it is possible that the slow reaction with St. -- involves reaction with a cation derived from THF which is not readily scavenged by NAH. Again further experimentation is required to resolve this problem.We thank the S.E.R.C. for the award of a CASE Studentship to J. R. L., the Ministry of Defence for laboratory facilities at Waltham Abbey and Dr D. H. Richards for his interest in the work and for fruitful discussions. I J. H. Haxeniiale, D. Beaumond and M. A. J. Rodgers, Trans. Faraday Soc., 1970, 66, 1996. '? F. Y. Joii ;id L. M. Dorfman, J . Them. Phys., 1973, 58, 4715. E. A. Shaede. H. Kurihara aiid L. M. Dorfman, Int. J . Radiut. Phys. Chem., 1974, 6, 47. E. R. Zabolotny and J. F. Garst. J . .4m. C!zrm. Soc., 1964, 86, 1645. E. .4. Robinson and G. A. Salmon, J. Phys. C'hen7.. 1975, 82, 382. G. A. Salmon and W. A. Seddon, Chem. Phys. Z,ett., 1974, 24, 366. B. Bockrath and L. M. Dorfman, J . Phys. Chem.. 1973, 77, 1002. G. A. Salmon, W. A. Seddon and J. W. Fletcher, Cun. J . Chem., 1974, 52, 3295. L. J. Giling, J. G. Kloosterboer, R. P. H. Rettschnick and J. D. W. van Voorst, C'hem. Phys. Lett., 10 11 12 1 :1 14 15 18 1 7 18 19 1971, 8, 457. F. S. Dainton, E. A. Robinson and G. A. Salmon, J . Phys. Chem., 1972, 76, 3897. T. J. Kemp, J. P. Roberts, G. A. Salmon and G. F. Thompson, J . Phys. Chem., 1968, 72, 1464. D. H. Ellison, G. A. Salmon and F. Wilkinson, Proc. R. Soc. London, Ser. A, 1972, 328, 23. G. V. Buxton, J. Kroh and G. A. Salmon, J . Phys. Chem., 1981, 85, 2021. J . A. Delaire, M. 0. Delcourt and J. Belloni, Radiat. Phys. Chem., 1980, 15, 255. R. Chang and C. S. Johnson, J . C'hem. Phys., 1964, 41, 3273. D. H. Paskovich, A. H. Reddoch and D. F. Williams, J . Chem. Soc., Chem. Commun., 1972, 1195. (a) Ions and Zon Pairs in Organic Reactions, ed. M. Szwarc (John Wiley and Sons, New York, 1972), vol I. (h) M. Szwarc, Carbanions, Living Polymers and Electron Transfer Processes (Interscience, New York, 1968). M. Fisher, G . Ramme, S. Claesson and M. Szwarc, Chem. Phys. Lett., 1971, 9, 306. J. H. Baxendale, D. Beaumond and M. A. J. Rodgers, Int. J . Radiat. Phys. Chem., 1970, 2, 39. (PAPER 2/58])

ISSN:0300-9599    

DOI:10.1039/F19827803645

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3659-3669 Photosensitised Dissociation of Water using Dispersed Suspensions of n-Type Semiconductors BY ANDREW MILLS* AND GEORGE PORTER Davy Faraday Research Laboratory of the Royal Institution, 21 Albemarle Street, London WlX 4BS Received 6th April, 1982 n-Type semiconductor powders (TiO, and SrTiO,) were used to sensitise the photochemical cleavage of water using U.V. light (A < 400 nm) under conditions of room temperature and pressure. Although several methods of piatinising these powders were used, one method in particular (involving precipitation of a platinum sol, by addition of an inert electrolyte, in the presence of the semiconductor powder) was found to produce an efficient photocatalyst for water reduction. Many different photocatalysts were tested for water reduction activity, using EDTA as an electron donor, and for water oxidation activity, using Fe3+ as an electron acceptor.In the absence of EDTA and Fe3+, U.V. irradiation of these photocatalysts liberated H, but 0, evolution was not observed. Reasons for these observations are discussed. In recent years numerous studies have been made on the use of semiconductors in photosynthetic and photocatalytic reacti0ns.l In particular, since the work of Fujishima and Honda2 a great deal of attention has focused on the use of semiconductors to cleave water photo~hemically,~~ and of the many availabL semiconductors it is the n-type oxides such as strontium titanate (SrTiO,) and titanium dioxide (TiO,) which have been most widely used, because of their high stabilities towards photo- corrosion and their favourable band energie~.~ Initially water splitting was achieved photoelectrochemically, using these n-type oxide semiconductors in the form of electrodes;, however, it now appears possible to use semiconductors in a particle form4 to carry out many of the reactions previously associated with semiconductor electrodes. In a particle system a ' short-circuited' photoelectrochemical cell (p.e.c.) may be constructed6 by depositing some platinum (Pt) onto a particle of a semicorl- ductor, the overall desired reaction occurring by electron and hole transfer at the two sites (i.e.Pt and semiconductor) on the particle. Such particulate systems appear, in general, much simpler and less expensive to use and construct than their p.e.c.counterparts. Also, materials may be used which are not available as single-crystal (or even polycrystalline) semiconductor electrodes, for reasons of high resistivity or difficulty in fabrication. Platinised semiconductor powders, such as TiO,, have been used to photo-oxidise water,4 acetic acid,lU alcohol^,^ hydrocar60ns,8 carbohydrates,'" active carbonld and biomass,lc and simultaneously reduce water. In the photochemical cleavage of water with semiconductor powders both oxidative and reductive reactions are believed to occur on the same particle, resulting in the simultaneous production of H, and 0,.9 On a particle the spatial separation between oxidative and reductive sites will be very smalllo in contrast to a p.e.c., thereby increasing the likelihood of back reaction.In order to reduce this experimental conditions are usually chosen to favour desorption of H, and 0, from the particle In general these conditions take the form of elevated temperatures and/or reduced pres~ures.~~-f Indeed Lehn et working on metallised SrTiO, powders, found that the photochemical cleavage of water only took place at reduced pressures 36593660 PHOTOSENSITISED DISSOCIATION OF WATER ( I 5 mmHgT) and was enhanced at elevated temperatures. Absence of water photo- decomposition, under conditions of room temperature and pressure, has also been reported for platinised TiO, powders by Pichat et ~ 2 1 . ~ In contrast to this, several workers have claimed water photodecomposition using semiconductor powders, under ambient conditions of temperature and pressure.For example, Bulatov and Khideke14a have claimed a quantum yield for H, production [O(H,)] of 6% for this reaction, using platinised rutile TiO, in 0.5 mol dm-3 H,SO,. However, this appears unlikely, as electrochemical studies have shown that the rutile form of TiO, does not have a sufficiently negative flat-band potential ( Vfb) to reduce water5 and, in addition, Jaeger and Bardll using e.s.r. found no evidence for the production of radicals usually associated with water decomposition (such as OH' and OH;) on U.V. irradiation of rutile (platinised and unplatinised)/H,O suspensions. Other claims for the photo- chemical cleavage of water, under conditions of room temperature and pressure, have been made by Kawai and Sakata,lr using a mixture of RuO,, TiO, and Pt powders [@(H,) = 0.0273, and by Gratzel and coworkers,12 using colloids of TiO, coupled with RuO, and/or Pt [@(H,) = 30+ 10x1. In view of these varying reports it was decided to look more closely at the photochemical cleavage of water, under ambient conditions, using platinised TiO, and SrTiO,, with the overall aim of producing a reproducible and efficient photocatalyst for this reaction.EXPERIMENTAL MATERIALS Powdered anatase titanium dioxide (TiO,) was obtained from B.D.H. (9973, Aldrich (99.9+%) and Laporte (Tiona G, 98%; Nb,O,, 0.15%). Other powders used included rutile TiO, (Tiona 010,98%; Nb,O,, 0. IS%, Laporte), strontium titanate (SrTiO,, 2j4 99.5%, Alpha Inorganics) and cadmium sulphide (CdS, 99.999 %, Koch-Lite).Ethylenediaminetetra-acetic acid (EDTA), sodium citrate, ferric chloride (Fe C1;6H20) and chloroplatinic acid (5% w/v H,Pt Cl6.6H,O) were purchased from B.D.H. (A.R. grade). A RuO,/TiO, powder was prepared, using 50 mg of ruthenium tetroxide and 5 gm of TiO, (B.D.H., anatase) in 50 cm3 of water. The reaction mixture was placed in a 100 cm3 stoppered glass flask, stirred for 5 days, filtered and the grey powder residue dried in air. This powder has been found to catalyse the reduction of cerric ions in aqueous s01ution.l~ METHODS Methods used for platinisation of the semiconductor powders were as follows : (A) An N,-purged suspension was prepared containing 1 gm of the semiconductor powder, 15 mg of chloroplatinic acid and 10 cm3 of a 40% formaldehyde solution, stabilised by methanol (B.D.H.).This suspension was subsequently stirred continuously on irradiation for 8 h in a quartz vessel with a 250 W Hg medium-pressure lamp (this method is analogous to that reported by Gratzel et a1.).12 (B) A Pt/citrate sol was prepared by refluxing, for 4 h, a solution containing 30 mg of chloroplatinic acid, 30 cm3 of a 1 % sodium citrate solution and 120 cm3 of water. A third (50 cm3) of the resultant Pt sol was stirred with 1 gm of the semiconductor powder and 5.8 gm of sodium chloride added. The destabilisation of the sol, followed by Pt precipitation, appeared to be complete within seconds of the addition. (C) A Pt sol was prepared as described in (B), in the presence of 1 gm of the semiconductor. (D) A suspension was prepared, containing 1 gm of the semiconductor powder in 50 cm3 of a Pt sol [described in (B)] and the solvent subsequently removed by rotary evaporation.(E) A suspension was prepared containing 1 gm of the semiconductor powder and 15 mg of chloroplatinic acid in 10 cm3 of acetic acid (1 mol dm-,). Saturation of this suspension with H, brought about the reduction of the Pt salt to the metal. For all the methods described above, t 1 mmHg = 13.5951 x 980.665 x 10F Pa.A. MILLS AND G. PORTER 3661 the final platinised powder suspensions were filtered and repeatedly washed with distilled water before being dried in air. All steady-state experiments were performed with an Applied Photophysics clinical reactor, using a 900 W Xe lamp and a cold-water infrared filter.A cut-off filter was used to restrict the light to wavelengths > 400 nm for some measurements, and a high-radiance monochromator used for action-spectra measurements. The concentration of evolved oxygen or hydrogen was measured with a Clark membrane electrode, purchased from Rank Bros. A detailed description of the experimental arrangement, sensitivity and calibration of this instrument is given elsewhere.l4 The solutions were stirred and thermostatted at 25 0.5 "C and H, and 0, formation confirmed by gas chromatography. Before each experiment the solutions were sonicated for ca. 1 min in a Dawe ultrasonic bath to disperse the powder. The solutions were then transferred to the photochemical cell and purged with oxygen-free nitrogen prior to illumination.The initial rate of H, (or 0,) evolution [or R(02J) was determined from the concentration of H, (or 0,) evolved over the first 4 min of irradiation. Photoacoustic spectra were recorded at the City University with the help of Mr C . Morrison. RESULTS AND DISCUSSION SACRIFICIAL SYSTEMS The photoreduction of water to H,, using just semiconductor powders, has been achieved in several systems but with low efficiencie~.*~?l~ This may be improved by depositing Pt onto the semiconductor powder particles, thus lowering the overpotential for water reduction.1° There are two major methods used to deposit Pt onto semiconductor powders. Developed by Gratzel et a1.l29 l6 the first method involves the preparation of a Pt/citrate sol, followed by removal of the protective citrate (using an ion-exchange resin), addition of the semiconductor powder and sonication of the resulting mixture.However, using this procedure we have encountered problems, including Pt adhering to the resin and incomplete precipitation of the non-protected Pt sol onto the semiconductor powder. Preferred12b over the citrate reduction method (because it is easier to perform and gives results of excellent reproducibility) is the second method of platinisation, i.e. photoplatinisation. Although there are several variations of this method, they all involve the U.V. irradiation of a semiconductor powder suspended in a solution containing a Pt salt and, usually, an electron donor (e.g. acetic acid,17 forrnaldehydel2 or ethanolle). In general this procedure is suitable only for photostable semiconductors because long, high-intensity irradiations are involved.With this in mind, several alternative [(B)-(E)], as well as established (A), platinisation techniques were used for producing photocatalysts, in order to determine a facile routine method of depositing Pt onto a support, resulting in reproducible batches of catalyst without elaborate techniques. The photocatalysts produced by these methods were compared1* for water reduction activity using a test system incorporating an electron donor (such as EDTA) to scavenge, efficiently and irreversibly, the ' hole' of the 'electron-hole pair' produced on irradiation of the semiconductor with band-gap light. The remaining conductance- band electron should then be able to migrate to a Pt site, where water reduction can occur.In a typical experiment a sonicated suspension of a platinised semiconductor powder (37 mg) in EDTA solution (37 cm3, lo-, mol dm-3) was placed into the Pyrex irradiation cell of a H,-detecting Clark membrane electrode, purged with N,, and the initial rate of H, production [R(H,)] determined on irradiation. The results of this work are shown in table 1 , from which it appears that only method (B) produces photocatalysts of similar (if not greater) activity to those produced by a photo- platinisation technique [i.e. method (A)]. Using method (A), other metals (Rh, Ru,3662 PHOTOSENSITISED DISSOCIATION OF WATER TABLE 1 .-COMPARISON OF THE ACTIVITIES OF DIFFERENT PHOTOCATALYSTS TOWARDS WATER REDUCTION AND OXIDATION R ( H 2 ) a in R(oz)a in type of semiconductor method of lo-, mol dm-, R(H2)a in lo-, mol dmP3 support platinisation EDTA (%) H,O (%) FeC1, (%) SrTiO, SrTiO, SrTiO, SrTiO, TiO, (anatase, B.D.H.) TiO, (anatase, B.D.H.) TiO, (anatase, B.D.H.) TiO, (anatase, B.D.H.) TiO, (anatase, B.D.H.) TiO, (anatase, B.D.H.) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Aldrich) TiO, (anatase, Laporte) TiO, (anatase, Laporte) TiO, (rutile, Laporte) TiO, (rutile, Laporte) TiO, (anatase, B.D.H.)/RuO, TiO, (anatase, B.D. H .)/RuO, TiO, (anatase, B.D.H.)/RuO, CdS CdS A1203 A1203 none B none B none 5 3 4 0 28 46 11 22 14 0 96 100 13 63 2 5 8 2 0 48 0 21 0 40 24 4 136 3 0 0 0 0 0 0 2.0 2.6 0 1.6 0.7 0 6.0 10.8 1.6 4.1 0 0 0 0 0 3.7 0 0 0 0 0.3 0 0.5 0 0 0 8.0 5.9 8.8 9.6 9.8 5.7 1 .o 5.2 10.3 21 .o 15.5 13.7 13.4 16.8 15.8 17.0 25.8 23.2 8.5 5.4 4.0 0 0 a A relative rate of H, production [R(,,J or oxygen production [R(,,,] of 100% corresponds to 1.2 x irradiation of a platinised [method (B)] TiO, powder (Aldrich, 37 mg), in the presence of a lo-, mol dm-, EDTA solution (37 cm3), with light of 360 k 20 nm and intensity 10l6 photon s-l (as determined using a calibrated thermopile) resulted in H, production with a formal quantum yield15 of 3 & 1 %.mol (gas) min-l; Ir, Co) were deposited on TiO, from their respective salts4h but only Rh showed any appreciable photocatalytic activity (see table 1).Surprisingly, SrTiO, appears a less active photocatalyst for water reduction than anatase TiO,, even though its flat-band potential (&) is more red~cing.~ Also, R(H2) values for anatase TiO, appear dependent on the commercial source of this material. One possible explanation of these results may lie in the surface chemistry dependence of n-type oxides on their method of preparation and subsequent chemical history.lg (This may also be one of the reasonsA. MILLS AND G . PORTER 3663 for the varying reports concerning photochemical water cleavage using these semi- conductors.) Interestingly, water reduction was also achieved in EDTA solution using a platinised rutile TiO, powder, even though this reaction is thermodynamically unfavourable (AG > 0), for reasons stated previously. However, niobium oxide (Nb,O,) is known to shift V,, for rutile cathodically,12b720 and this oxide is present as an impurity (0.15 %) in the rutile TiO, powder used.Accumulation of negative charge (or conductance-band electrons) at a Pt site owing to the irreversible nature of the EDTA oxidation by valence-band holes may also contribute towards the shift in Gb, increasing the likelihood of water reduction. This latter effect would not be expected in the absence of EDTA and, indeed, no H, evolution was observed with rutile under such conditions. It is known from electrochemical studies that RuO, electrodes exhibit overvoltages for H, evolution similar to those for Pt,,l and recently Amouyal et ~ 1 . ~ ~ reported RuO, to be an effective redox catalyst for H, generation from a Ru(bipy)t+/methyl viologen/EDTA sacrificial system.However, Gratzel and have reported evidence that, when bound to TiO,, RuO, and Pt specifically catalyse water oxidation and reduction, respectively. We have found RuOJTiO, powders on irradiation with U.V. light, in either the presence or absence of EDTA, showed little ability to reduce water. In agreement with the results of Darwent and P ~ r t e r , ~ ~ ~ ~ irradiation (2 > 400 nm) of a CdS powder (1 mg ~ m - ~ ) , suspended in a EDTA solution (37 cm3), resulted in H, evolution, the rate of which was enhanced greatly (ca. x 45) by platinising the CdS powder using method (B). Harbour et al.15" have recently reported no H, production from CdS powders except in the presence of both Pt and EDTA, but it is worth noting that method (B) represents a quick, easy and efficient method of platinisation of powders which, in contrast to established photoplatinisation techniques, is amenable to platinisation of powders not photostable to U.V.light (e.g. CdS, CdSe and dye-coated semiconductors) or which do not absorb light (e.g. Al,03).17a7 24 The rate of water reduction, sensitised by the photocatalysts, in the presence of X/nm FIG. 1.-Action spectrum of platinised [method (B)] anatase TiO, (B.D.H., 37 mg) in EDTA solution (0.1 mol drnp3, 37 cm3).3664 PHOTOSENSITISED DISSOCIATION OF WATER EDTA showed a strong wavelength dependence, as illustrated for platinised anatase TiO, (fig. 1). Hydrogen was only produced by light of 3, -c 400 nm since, above this wavelength, no light is absorbed by anatase TiO,, which has a band-gap of 3.2 eVll (corresponding to 388 nm).This action spectrum (fig. 1) matched the photoacoustic spectrum recorded for the photocatalyst which, in turn, was identified from the literaturez5 as anatase TiO,. Work with Al,O, (Ebg > 7 eV24) indicated that the oxide support must be light-absorbing for H, evolution to occur, and that light (A > 300 nm) absorbed by EDTA does not lead to detectable H, production. The photocatalysts were also tested for water oxidation activity by using an electron acceptor (FeC1,) to scavenge the ‘electron’ of the ‘electron-hole pair’ which is produced on irradiation of the semiconductor with band-gap light. This allows the remaining valence-band ‘hole’ to migrate to the semiconductor surface where water oxidation can occur.Using a procedure similar to that described for the determination of R(H,) [repiacing EDTA with FeCl, (37 cm3, mol dm-, H,SO,)] the initial rate of 0, production [R(o,,] was determined for each catalyst (37 mg) with an 0, Clark membrane electrode, and the results are also shown in table 1. From these results it would appear that platinisation reduces the abilities of the semiconductors to oxidise water. This may partly be the result of decreased light absorption by the semiconductor, due to the Pt deposited onto its surface. Also, thermodynamically, there is the possibility of Pt competing with water for oxidation. However, although Pt oxidation is claimed by Kawai and Sakata26 to occur in a system containing Pt, TiO, and active carbon, Sat0 and Whiteld found no evidence for this but rather that platinum oxide is reduced by TiO, on U.V.irradiation. Although the valence-band ‘ hole ’ is an extremely oxidising species, having an overpotential for water oxidation of ca. 1.9 eV, some workers have claimedlc. improved yields by coupling or mixing the TiO, particles with RuO, (one of the best electrode materials for water oxidation).lc! 27 However, a RuO,/TiO, powder (see table l), upon illumination in the presence of FeCl,, showed no enhancement of R(OP) compared with the TiO, support (similar results were found for W03).,* From table 1, Co deposited onto TiO, appears to enhance R(02), although Sat0 and WhiteLd claim that Co (along with Ni) is readily photo-oxidised by TiO,.Interestingly, Co2+ ions have r e ~ e n t l y ~ ~ ~ ~ ~ been used to catalyse water oxidation by Ru3+, Os3+ and Fe3+ tris-bipyridyls. Work with Al,O, indicated that the oxide support must be light absorbing for 0, evolution to occur and, although Fe3+ ions are known31 to oxidise water on irradiation with U.V. light, no evidence for this was found under the experimental conditions of pH and concentration used. mol dm-, in 5 x NON-SACRIFICIAL SYSTEMS From the above work it is clear that platinised TiO, and SrTiO, can photochemically reduce or oxidise water in the presence, respectively, of an electron donor or acceptor. However, in the photochemical cleavage of water both processes would have to occur simultaneously, i.e.(1) hv semiconductor .+ e- + h+ electron-hole pair 2e-+ 2H+ 2 H, 4h+ + 2H,O -, 0, + 4H+A. MILLS AND G . PORTER 3665 4 I 0 time/min FIG. 2.-Typical H, concentration, measured on a Clark membrane electrode, against time profiles on U.V. irradiation of several aqueous anatase TiO, (B.D.H.) suspensions (37 mg of semiconductor dispersed in 37 cm3 of H,O), platinised by the following methods: (a) B; (b) A; (c) E and ( d ) none. The arrow denotes the start of U.V. irradiation. leading to H, and 0, evolution from the same particle. The rates of H, evolution in water (i.e. no EDTA) for each of the photocatalysts were determined (see table 1) and, although reduced by ca. 10 times, appeared to follow similar trends to those found for water reduction in EDTA (an example of some typical H, concentration against time profiles is given in fig.2). By varying the amount of chloroplatinic acid used in method (A), or the amount of Pt colloid used in method (B), the estimated Pt concentration (expressed as a percentage of Pt) deposited onto the semiconductor powder may be varied. The results of this work, using platinised TiO,/water suspensions, are shown in fig. 3 and indicate, for both procedures, an optimum Pt concentration of ca. 0.5% (ca. 20 times lower than those used by other workers).l29 1 7 b 9 24 Also, from fig. 3(a), at Pt levels above the optimum, R,H2, gradually falls (presumably due to the Pt screening the TiO, surface, thereby reducing the actual light intensity ‘seen’ by the semiconductor). This would indicate that Pt deposit levels > 0.50,; are not only unnecessary but may also be detrimental towards the rate of H, production.The rate of water reduction using a semiconductor photocatalyst was found also to be a function of semiconductor concentration ([SC]), as shown in fig. 4. The against [SC] profile is very similar to that reported by Rao et aL3, for ZnO and may be explained as follows. Initially, by increasing [SC] the amount of light absorbed by the semiconductor increases and therefore increases (A-B). Eventually a point is reached at which all the incident light is absorbed and R(H,) can increase no further (B). A further increase in [SC] only reduces the penetration depth of the incident light. However, this may well increase the likelihood of losing scattered light to the exterior and, in turn, account for the reduction of R(H,) with further increases in [SC]. Reduction of the penetration depth to such a level that the light loss due to scattering is almost constant may then account for the levelling-off of [i.e. (C-D)] at higher with pH for a platinised TiO, powder [SCI - Fig.5 illustrates the variation of3666 PHOTOSENSITISED DISSOCIATION OF WATER I I I 2 Pt (%) ( b ) 100- * 1 I 0. I 0.2 0.3 Pt (%) FIG. 3.-Relative rate of H, production [R!H?)] on U.V. irradiation of aqueous suspensions (1 mg ~ m - ~ , 37 cm3) of anatase TiO, (Aldrich) platinised by (a) method (A) [R(,2, = 100 corresponds to 1.3 x mol dm-3 (H,) min-l] or (b) method (B) [R&) = 100 corresponds to 9 x lo-' mol dm-3 (H,) min-'1. dispersed in water.From electrochemical the conductance and valence bands of TiOz electrodes remain constant relative to the reduction and oxidation potentials of water regardless of pH, and therefore (assuming TiO, particles also show this behaviour) should not produce any variation of with pH. In contrast, the surface nature of TiO, is pH-dependent, and for anatase TiO, two acid-base equilibria are known:34 -Ti-OH + H++ -Ti-OHi pKa,-4.98 -Ti-0-+ H++ -Ti-OH p Ka2- 7 .a (4) resulting in a point of zero zeta potential35 (P.z.z.P.) of 6.39. Work by Bard et aZ.3s on anatase TiO, particulate systems indicates that platinisation shifts the p.z.z.p. toA. MILLS AND G . PORTER 3667 I 6 FIG. 0 5 0 100 I50 [ SC I Img 4.-Relative rate of H, production [R(&)] on U.V. irradiation of varying amounts of semiconductor (in this case, platinised [method (B)] anatase TiO, (B.D.H.)} dispersed in 37 cm3 of H,O.I 1 I 0 I 2 3 4 5 6 7 8 9 I011 PH FIG. 5.-Relative rate of H, production [R(H2)] as a function of pH, for a platinised [method (B)] anatase TiO, (B.D.H.) powder (37 mg) dispersed in 37 cm3 of solution. Variation of pH was achieved by adding H,SO, or NaOH. lower pH values (ca. pH 4), probably due to anion (Cl-) adsorption. This p.z.z.p. (i.e. pH 4) corresponds roughly to the position of maximum R(H2) (see fig. 5), indicating that water reduction occurs best on a neutrally charged surface. A positively charged surface and, to a lesser extent, a negatively charged surface (produced by lowering or raising the pH, respectively) appear less favourable for water reduction.However, other factors may also be involved, for example TiO, shows ageing effects in H,SO, 37 and is soluble in this and alkaline In addition, it was noticed that the ease of dispersion of the photocatalyst decreased with decreasing pH.3668 PHOTOSENSITISED DISSOCIATION OF WATER So far we have described only the evolution of H, on U.V. irradiation of our photocatalysts in the presence of water. However, under identical irradiation conditions as those used above, no 0, evolution was observed. Although measurement of small concentrations of 0, in the presence of H, with a Clark electrode did prove difficult,14 this result was confirmed by gas chromatography. Results similar to these have been reported by Gratzel et al.,g712 where several hours of irradiation lapse before 0, evolution is observed on U.V.irradiation of Pt/RuO,/TiO, colloids. This effect has been attributed12b to the time taken for O,, which is readily photoadsorbed by TiO, as Oi-,39 to occupy all the vacant adsorption sites and has been suggestedg as a means of separating H, and 0, produced from such systems. Blank experiments involving U.V. irradiation of air-saturated aqueous suspensions of TiO, (and Pt/TiO,) confirmed that oxygen is readily photoadsorbed onto these powders (ca. 2 x lop6 mol 0, for 40 mg semiconductor dispersed in 40 cm3 water). The platinised forms of TiO, appeared to photoadsorb 0, faster ( t = 10 min) than TiO, ( t = 20 min). Attempts to displace the photoadsorbed 0, (produced on irradiation of a Pt/TiO, powder suspended in water) by addition of sodium phosphateg (Na,PO,, 5 x lo-, mol dm-3) proved ineffective.Also, irradiation of a Pt/TiO, powder suspen- sion (40 cm3, 1 mg ern-,) in the presence of 0.1 mol dm-, Na,PO, solution not only failed to liberate 0, but prevented water reduction as well. Similar inhibiting effects for Na,PO, have been found by Malati and Seager40 in the photo-oxidation of primary alcohols by TiO, powders and were attributed to the adsorption of phosphate ions41 onto the surface sites, rendering them less reactive. Although most of the irradiations performed on the photocatalysts were of short duration (ca. 30 min), initial work on prolonged irradiations ( t > 3 h) of Pt/TiO, (anatase) suspensions indicated that no 0, is evolved and H, evolution ceased under such conditions.This may well be the result of 0, (and Oi-) accumulation onto the TiO, particle’s surface, to such an extent that reduction of 0, and oxidation of Oi-- become themajor photoprocesses involved. However, regeneration ofthe photocatalyst appears possible by saturating the suspension with N,, although the reasons for this remain unclear. The adsorption of one or more of the photochemically produced species on to a semiconductor particle’s surface may prove an important limiting factor when considering such materials for use in practical solar-energy devices. few methods exist for probing the details of the reaction mechanisms or the catalyst properties in situ. Photoacoustic or reflectance spectroscopic techniques allow us to determine the energy of the band of the powders themselves, and spin trapping has been employed to identify intermediate radicals formed in these processes.ll The fact that the particles scatter and reflect as well as absorb light makes evaluation of anything but formal quantum yield^^^^,^ difficult. In a recent development Bard and employed photoelectrophoretic and electrochemical techniques to characterise a particulate TiO, photocatalyst.Work is currently in progress in this laboratory using techniques such as these to achieve a better understanding of the semiconductor particle systems described above. As pointed out by Bard and We thank the S.E.R.C., the E.E.C. and G.E. (Schenectady) for financial support of this work. We are indebted to Miss M. L. Zeeman for her assistance.We thank Mr C. Morrison and the City University for the photoacoustic spectroscopy and Dr A. Harriman and Dr J. Darwent for many helpful discussions.A. MILLS AND G. PORTER 3669 1 (a) B. Kraeulter and A. J. Bard, J. Am. Chem. SOC., 1978, 100, 2239, 5985; 1977, 99, 7729. (b) S. N. Frank and A. J. Bard, J. Phys. Chem., 1977, 81, 1484. (c) T. Sakata and T. Kawai, Nouv. J. Chim., 1981, 5, 279; Nature (London), 1980, 286, 474; 1979, 282, 283. ( d ) S. Sat0 and F. M. White, J. Phys. Chem., 1981,85, 336. (e) J. C. Hemminger, R. Carr and G. A. Somarjai, Chem. Phys. Lett., 1978, 57, 100. cf) S. Sato, J. Chem. SOC., Chem. Commun., 1982, 26. A. Fujishima and K. Honda, Nature (London), 1972, 238, 37; Bull. Chem. SOC. Jpn, 1971, 44, 1148. (a) M.S. Wrighton, D. S. Ginley, P. T. Wolczanski, A. B. Ellis, D. L. Morse and A. Lintz, Proc. Natl Acad. Sci. USA, 1975, 72, 1518. (b) M. S. Wrighton, A. B. Ellis, P. T. Wolczanski, D. L. Morse, H. B. Abrahamson and D. S. Ginley, J. Am. Chem. SOC., 1976,98,2774. (c)T. Watanabe, A. Fujishima and K. Honda, Bull. Chem. SOC. Jpn, 1976, 49, 355. * (a) A. V. Bulatov and M. L. Khidekel, Izu. Akad. Nauk SSSR, Ser. Khim., 1976, 1902. (6) G. N. Schrauzer and T. D. Guth, J. Am. Chem. Soc., 1977, 99, 7189. (c) H. van Damme and W. K. Hall, J. Am. Chem. SOC., 1979, 101, 4373. ( d ) S. Sat0 and J. M. White, Chem. Phys. Lett., 1980, 72, 83. (P) T. Kawai and T. Sakata, Chem. Phys. Lett., 1980, 72, 87. (f) K. Domen, S. Naito, M. Soma, T. Onishi and K. Tamaru, J. Chem. SOC., Chem. Commun., 1980, 543.Cg) F. T. Wagner and G. A. Somorjai, Nature (London), 1980,285, 559. (h) J. M. Lehn, J. P. Sauvage and R. Ziessel, Nouu. J. Chim., 1980, 4, 623. H. P. Maruska and A. K. Ghosh, Solar Energy, 1978, 20, 443. P. Pichat, J-M. Herrmann, J. Disdier. H. Courbon and M-N. Mozzanega, Nouv. J. Chim., 1981, 5, 627. S. Sato and J. M. White, Chem. Phys. Lett., 1980, 70, 131. M. Gratzel, Acc. Chem. Rex, 1981, 14, 376. P. C. Jaeger and A. J. Bard, J. Phys. Chem., 1979, 83, 3146. J. Kiwi, E. Pelizetti, M. Visca and M. Gratzel, J. Am. Chem. SOC., 1981, 103, 6324. Soc., Dalton Trans., 1982, 1213. ti A. J. Bard, J. Protochem., 1979, 10, 50. lo A. J. Bard, J. Phys. Chem., 1982, 86, 172. l 2 (a)D. Duonghong, E. Borgarelloand M. Gratzel, J . Am. Chem. Soc., 1981,103,4685.(b) E. Borgarello, l3 (a) A. Mills and M. L. Zeeman, J. Chem. SOC., Chem. Commun., 1981, 948. (b) A. Mills, J . Chem. l4 A. Mills, A. Harriman and G. Porter, Anal. Chem., 1981, 53, 1254. l5 (a) J. R. Darwent and G. Porter, J . Chem. Soc., Chem. Commun., 1981, 145. (b) J. R. Darwent, J. Chem. SOC., Faraday Truns. 2, 1981, 77, 1703. (c) J. R. Harbour, R. Wolkow and M. L. Hair, J. Phys. Chem., 1981, 85, 4026. P. A. Brugger, P. Cuendet and M. Gratzel, J. Am. Chem. SOC., 1981, 103, 2923. A. J. Bard, J. Am. Chem. SOC., 1978, 100, 4317. A. Mills, J. Chem. SOC., Chem. Commun., 1982, 367. l9 G. D. Parfitt, Prog. Surf. Membr. Sci., 1976, 11, 181. 2o P. Salvador, Sol. Energy Mater., 1980, 2, 413. 21 D. Galizzioli, F. Tantardini and S. Trasatti, J. Appl. Electrochem., 1974, 4, 57. 22 E. Amouyal, P. Keller, A. Moradpour, J. Chem. Soc., Chem. Commun., 1980, 1019. 23 R. Humphry-Baker, J. Lilie and M. Gratzel, J. Am. Chem. Soc., 1982, 104, 422. 24 W. W. Dunn and A. J. Bard, Nouv. J. Chim., 1981, 5, 651. 25 M. J. Adams, B. C. Beadle, A. A. King and G. F. Kirkbright, Analyst (London), 1976, 101, 553. 26 T. Kawai and T. Sakata, J. Chem. SOC., Chem. Commun., 1979, 1047. *: M. Morita, C. Iwakura and H. Tamura, Electrochim. Acta, 1978, 23, 331. 28 J. R. Darwent and A. Mills, J. Chem. Soc., Faraday Trans. 2, 1982, 78, 359. 29 V. Ya. Shafirovich, N. K. Khannov and V. V. Strelets, Nouu. J. Chim., 1980, 4, 81. 30 A. Harriman, G. Porter and P. Walters, J. Chem. Soc., Faraday Trans. 2, 1981, 77, 2373. 31 A. A. Krasnovsky and G. P. Brin, Dokl. Akad. Nauk SSSR, 1962, 147, 656. 32 M. V. Rao, K. Rajeshwar, V. R. Pal Verneker and J. Du Bow, J. Phys. Chem., 1980, 84, 1987. 33 E. C. Dutoit, F. Cardon and W. P. Gomes, Ber. Bunsenges. Phys. Chem., 1976,80, 475. 34 P. W. Schindler and H. Gamsjager, Discuss. Faraday SOC., 1971, 52, 286. 3s B. Reichman and C. E. Byvik, J. Electrochem. SOC., 1981, 128, 2601. 37 L. A. Harris and R. H. Wilson, J. Electrochem. SOC., 1976, 123, 1010. 3H Handbook of Chemistry and Physics, ed. R. C. Weast (CRC Press, Cleveland, Ohio, 61st edn, 1980). 3y (a) A. H . Boonstra and C. A. H. A. Mutsaers, J. Phys. Chem., 1975, 79, 1694. (6) G. Munuera, V. Rives-Amau and A. Saucedo, J. Chem. SOC., Faraday Trans. 1, 1979,75,736. (c) A. R. Gonzalez- Elipe, G. Munuera and J. Sona, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 748. l7 (a) A. J. Bard, W. W. Dunn and B. Kraeulter, U S . Patent, 1981, no. 4 264 421. (6) B. Kraeutler and W. W. Dunn, Y. Aikawa and A. J. Bard, J . Am. Chem. SOC., 1981, 103, 3456. 40 M. A. Malati and N. J. Seager, J . Oil Colour Chem. Assoc., 1981, 64, 231. 41 H. P. Boehm, Discuss. Faraday SOC., 1971, 52, 264. (PAPER 2/585)

ISSN:0300-9599    

DOI:10.1039/F19827803659

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 3671-3677 Study of Calcium Ion Binding to D-Ribose in Aqueous Solutions using Hydroxy-proton Resonance Shifts? BY MARTYN C. R. SYMONS,* JOHN A. BENBOW AND HEATHER PELMORE Department of Chemistry, The University, Leicester LE 1 7RH Received 6th April, 1982 Resolved hydroxy-proton resonance features for various forms of D-ribose in dilute aqueous solutions shift markedly on the addition of calcium chloride. Down-field shifts are assigned to OH groups which bind to Ca2+ ions. These shifts level off in the 1.0-1.5 mol dm-3 region when complexing is complete. Resonances for OH groups which are not involved in complex formation shift linearly to high-field, the shift being greater than that expected for the sum of general Ca2+ and C1- ionic effects.This is because the Ca2+ ions are removed by complex formation, the up-field shift being largely due to the effect of chloride ions. When chloride is replaced by perchlorate, this up-field shift is greatly enhanced. Approximate equilibrium constants for complex formation have been derived, and the influence of conformational changes is discussed. It is commonly supposed that aqueous solutions of sugars exhibit only a single hydroxy-proton resonance because of rapid proton exchange with water. We have recently shown that by cooling, and carefully controlling the pH of the solutions, it is possible to slow down this exchange and obtain well resolved OH-proton resonance features, some of which even exhibit coupling to C-H In an attempt to exploit this discovery we have studied the effect of added electrolytes on these resonances.Results for D-ribose with calcium ions were particularly well defined, and have been singled out for detailed discussion. This system has been studied previously using alkyl-proton resonance shift^,^ 9 which are sensitive to conformational changes induced by complex formation. EXPERIMENTAL Sugars were of the highest grades available and were used as supplied; further purification had no effect on the n.m.r. spectra. The pH values of the solutions were measured at room temperature or at 0 OC with a Pye Dynacap H05 Eo 2 meter modified to take Eo 7 electrodes. The pH values were controlled by NaH malate/maleic acid buffer (1 mmol dm-3), and optimum conditions were sought to give the narrowest signals. N.m.r.spectra were recorded on a Jeol PS-100 spectrometer over a range of temperatures (measured with a Comark thermocouple). In all cases, the possible interference of spinning sidebands from the strong peak for water was checked by altering the rate of sample spinning. RESULTS AND DISCUSSION The spectrum in the 0-H region for D-ribose is shown in fig. 1. This shows features for all four of the common forms of D-ribose, the a- and B-pyranose forms (I and II), and the a- and /3-furanose forms (I11 and IV). The resonances for the anomeric OH protons (0, H) are, as u~ual,l-~ shifted markedly to low-field of the remaining resonances and hence were easier to assign and monitor. In addition to these four t Taken as Solvation Spectra, Part 71. 367 13672 CALCIUM ION BINDING TO D-RIBOSE I I I I 4 3 2 1 6 (PPm) FIG.1.-N.m.r. spectrum for a solution of D-ribose (1.4 mol dm-3) in water at pH 5.75. TPS reference is present at ca. 1.2 mmol dm-3. (i) /3-p, (ii) p-f, (iii) or-p, (iv) a-f, all anomeric hydroxys; (v) non-anomeric hydroxys; (vi) sideband. OH OH HOHh 0 OH w OH OH 111 N forms, which equilibrate slowly, the two six-membered rings both exist as two rapidly equilibrating forms (I, II,,). These give weighted mean resonance features in the n.m.r. spectrum. On the addition of calcium chloride (fig. 2) or calcium perchlorate (fig. 3) all OH resonances underwent marked shifts (table 1). Those which initially move down-field display minima in the 1.0 mol dm-3 region, whilst those which move up-field do so I, and 11,M.C. R. SYMONS, J. A. BENBOW A N D H. PELMORE 3673 h E a v 2.0 ;k 0.1 0.5 1.0 1.5 [CaCI, I /mol dm-3 FIG. 2.-Shifts (in ppm) in the OH-proton resonance features for D-ribose in water as a function of the concentration of calcium chloride. (i+ii) CaCl, water shift (i) with D-ribose (1.4 mol dm-3) present and (ii) without D-ribose present; (iii) 02, 3, 4H pyranose, not specifically assigned; (iv) O,, 3, 4H furanose, not specifically assigned; (v) [O,H @-p)+O,H @-f)]; (vi) [O,H (a-p)]; (vii) [O,H (a-f)]. monotonically. In addition, features assigned to the a-forms increased whilst those for the p-forms decreased in intensity as the salt concentrations were increased. There are two possible causes of these spectral shifts, direct ion effects and changes in conformation induced by ion-binding. The latter can only apply to the six-membered ring systems, and since the resonances for the five-membered ring molecules shift as strongly as the others, we start by attempting to assign all shifts to direct ion effects, and then consider possible additional shifts caused by conformational changes.The shifts induced by calcium chloride and calcium perchlorate on water protons in the presence and absence of D-ribose are indicated in fig. 2 and 3. These are averaged shifts caused by both cations and anions. We have previously shown that for methanolic6 and aqueous' systems, it is possible to derive separate shifts due to the cations and anions. The results show that C10, ions cause a rapid up-field shift, C1- ions have only a minor effect whilst Ca2+ ions cause a down-field shift.We expect that the OH protons of sugars will be affected similarly by these ions so the shifts would resemble those for water protons in the absence of any preferential interactions. Clearly the results are not statistical. The down-field shifts must be due to preferential complexing to Ca2+ ions, as has been inferred previously.*~ The fact that these trends3674 n 5 a W CALCIUM ION BINDING TO D-RIBOSE r 4-01 I I I I 0.1 0.5 1.0 1.5 [Ca(ClO,), l/mol dm-3 FIG. 3.--Shifts (in ppm) in the OH-proton resonance features for D-ribose in water as a function of the concentration of calcium perchlorate. (i) Ca(ClO,), water shift with D-ribose (1.4 mol dm-3) present; (ii) Ca(ClO,), water shift with no D-ribose present; (iii) 0,.3. ,H pyranose, not specifically assigned; (iv) 02, 3, ,H furanose, not specifically assigned; (v) [O,H (B-p)+O,H (B-f)]; (vi) [O,H (a-p)]; (vii) [O,H (a-f)]. are reversed at high salt concentrations suggests that, in our solutions, complexing is almost complete in the 1.0-1.5 mol dmP3 region, and hence shifts caused by Ca2+ ions in the complexes can be calculated. Since these shifts are almost equal for the chloride and perchlorate salts, it seems that, initially, there is almost no anion effect. This surprising conclusion is discussed below. The shifts estimated are in the region of 0.8 ppm. These are larger than the cation-induced shifts for the C-H pro tons,*^ but are in good accord with values calculated from the molar shifts for Ca2+ in water.Thus an average value for Ca2+ in a hexahydrate is ca. 0.6 ppm. Similarly, for methanolic solutions8 we estimate a shift of ca. 0.8 ppm for the hexasolvate. This agreement reveals satisfactory internal consistency with the theories involved in obtaining these values. Since Ca2+ ions are being effectively scavenged by certain forms of D-ribose, they must be less available for general interactions with other protons, and in particular they are unlikely to interact with those OH protons in the Ca2+ complexes which are not involved in binding the ions. Hence these protons will only be susceptible to the effects of the anions. If these were purely random, the trends should be comparable with those deduced for pure water, indicated in fig. 2 and 3. In fact, those for the anomeric OH protons are greater than these predictions. This suggests some preferential interaction, but absence of curvature in these trends shows that this mustTABLE l.-lH N.M.R.DATA FOR HYDROXY PROTONS OF D-RIBOSE (1.42 mol dmP3 IN WATER) AS A FUNCTION OF ADDED CALCIUM CHLORIDE' anomeric hydroxy" non-anomeric hydroxyC waterC CaCl,/ % com- TPSd mol dm-3 anomer positionb pH range 6 A v ~ pH 6 Av; pH 6 Av, pH 6 ref. 0.00 0.072 0.196 0.496 1 .082e 1.591 P-P a- f a-p a-f 8-P E-P a-f P-P P-f 0c-P a-f P-P a-f a-p a-f P-f } P- f a-P a-f P- f 0r-P a-f P-P} - - - - ca. 57 ca. 30 ca. 13 - - - - - - - - 51 36 13 45 40 15 5.2-6.8 3.82 15 5.8 ca. 3.62 - 5.6-5.9 3.44 < 15 5.6-5.9 ca. 3.12 - 5.6-5.9 5.1-6.6 3.77 20 5.5 - - - 3.50 < 22 5.45 3.26 sh ca.5.5 5.0-6.4 3.74 20.5 5.5 - - - 3.59 < 27 5.35 3.45 sh ca. 5.4 4.9-6.3 3.64 18 5.3 ca. 3.46 sh < 5.3 ca. 3.6 sh < 5.3 - - - 4.7-6.1 3.44 18.5 4.95 - - - 3.68 16 4.9 3.90 14.5 ca. 4.75 4.3-5.6 3.23 16 4.75 ca. 3.1 - - 3.60 14 4.65 3.90 11 4.65 1.95 s 2.09 2.08 2.09 s 2.18 2.42 2.52 2.90 2.50 2.72 2.52 2.90 2.24 2.72 31.5 sh 31 34 33.5 sh sh 23.5 21 sh sh 15 23 13 12 6.4-6.6 6.4-6.6 6.4-6.6 6.3 ca. 5.8 6.05 5.75 5.75 ca. 5.6 ca. 5.6 5.0 5.5 ca. 5.0 ca. 5.1 4.85 5.25 4.75 4.95 1.26 5.4 5.6-6.7 3.90 3 P ? 1.25 6 5.5-6.5 3.92 VJ e x 0 2, ? VJ 1.23 5.4 5.2-6.2 3.94 " 9 W m z 1.19 5.7 5.1-6.2 3.92 W 1.13 6 4.8-5.7 3.92 " 3 cd m r z 1.05 6 4.3-5.5 3.95 0 ~ m a Buffered solution [see ref. (3) for details]. Approximate values calculated from integrations. 6 = chemical shift (in ppm) measured from tallest alkyl feature of P-D-ribopyranose; Av4 = the half-height width (in Hz) at optimum pH; pH = the optimum pH for narrowest resonances; sh = shoulder; s = distinct single peak.Internal TPS to test stability of alkyl reference peak, i.e. chemical shift of alkyl reference w 4 wl o\ peak from TPS. The peak selected was shown to be insensitive to added salt. % Composition determined at ca. - 16 OC.3676 CALCIUM I O N B I N D I N G TO D-RIBOSE be relatively minor compared with the calcium binding. In fact, comparing the observed shifts for Ca(ClO,), with that estimated for perchlorate ions (1.6 ppm for a tetrasolvate) shows that there is still a long way to go before the limit is reached. However, for chloride systems the limiting shift in water is estimated as 0.7 ppm for the hexahydrate.The maximum shift observed is 0.55 ppm for the anomeric 0,H proton in the a-pyranose molecule, which is close to this limit. Since no curvature was observed, this suggests that the limiting shift for anomeric OH protons is considerably greater than that for water protons. This is not unexpected in view of the greater acidity of these OH proton^.^ In our view this higher acidity may also explain the apparent preferential interaction between anomeric OH protons and the anions. These groups will form stronger hydrogen bonds than water or the other OH protons, so at least a small preference might be expected. So far as we know, however, there is no precedent for such preferential interaction with anions.This preferential interaction with anomeric protons does not seem to be extended to those involved in Ca2+ binding. We suggest that this is largely a steric effect. We conclude that the general form of the shifts can be understood purely in terms of cation and anion effects. We now consider, in the light of previous work,,. 5 9 why different OH proton resonances are affected in such different ways. Angyalg has found that the ions most firmly bound to sugar molecules are Ca2+, Sr2+, Ba2+ and La3+, and that these strongly prefer ax-eq-ax sites, as illustrated in insert V. Sites with only two well disposed OH groups are relatively ineffective. This result is not surprising when one considers that the sugars are in direct competition with water molecules. These, especially in view of the high concentration of free lone-pair units,1° are freely available for cation solvation, and the sugars only compete at greater than statistical levels when there is an efficient chelating effect.If we accept the ax-eq-ax rule, then we expect that the ,C, form of a-pyranose (I,) will form a good complex, and so also will the 'C, form (Ib). The former should cause the 0 , H proton resonance to shift strongly to low field. The latter leaves this OH group unaffected, so that anions are expected to cause an up-field shift for this form (Ib). Since these two forms are in rapid equilibrium, the nett effect should be a compromise. The observed trend is to low fields, but the shift is small relative to the a-furanose form. This dominance of the 4C, (I,) form is in agreement with the suggestion that the concentration of this form is ca.twice that of the 'C, term., For the a-pyranose conformations only the lC, form (IIb), which is in ca. 25% abundance,, has an ax-eq-ax unit, so that the anomeric OH-proton resonance should be unaffected by Ca2+ ions. The rapid up-field shift accords well with this. The other OH-proton shifts for the lC, (11,) form should show a small initial down-field shift, but because of the equilibrium this will be modified by the anion effect on 11,. Some of the pyranose resonances are indeed only weakly affected, but unfortunately overlapping is so severe that it proved impossible to unravel the individual shifts. The five-membered ring systems are not in two forms, and hence the a-form (111) having an ax-eq-ax unit should bind strongly to Ca2+, whereas the a-form (IV) should not.This is supported by our results. The anomeric resonance for I11 shows a veryM. C. R. SYMONS, J. A. BENBOW AND H. PELMORE 3677 TABLE 2.-ASSOCIATION CONSTANTS FOR THE BINDING OF CALCIUM IONS TO D-RIBOSE ~~ ~ 52 2.2 1.4 0.7" 27 5.3 3.4 O.sa - 9 4.6 2.6 0. 5b a Ref. (5); present work. marked down-field shift, which we take as a measure of the real limiting shift for Ca2+ ions. The poorly resolved features for the other OH protons also shift strongly, and to about the same extent (fig. 2 and 3). For the p-furanose structure (IV) the shifts are dominated by the anions. Thus our results fit in well with those of others4$ On a quantitative basis agreement is less satisfactory, but this is probably not a real problem, since our results are less accurate than those of Lenkinski and R e ~ b e n .~ We have used their method, involving resonance shifts, to obtain Scatchard plots and hence approximate formation constants. Our results are given in table 2 together with those of Lenkinski and Reuben. The trends are very similar, but our data, obtained at -9.0 OC, are close to theirs at room temperature, whereas they should be appreciably larger. We now consider the effects, if any, of changes in conformational equilibria. For this, we use the observation3 that the resonance for a given axial OH proton is on the high-field side of that for the corresponding equatorial OH proton. Hence, for the P-pyranose forms, complexation with Ca2+ ions should shift the equilibrium in favour of 11,.This should cause the anomeric OH-proton resonance to shift to high-fields. This is indeed observed, and the rate of shift is slightly greater than that for the p-furanose form, which is unaffected by such equilibria. Probably this shift in conformation is responsible for this, but we cannot draw any firm conclusions. For the a-pyranose structures both can form complexes, and we have no evidence for any major change in the equilibrium on complexation. Finally, we should mention the linewidth changes indicated in table 1 . We have already stressed that there are very large pH effects on the exchange rates, and widths probably still reflect the exchange process. However, there is a general narrowing of the resonances at high salt concentrations, suggesting that this slows down exchange. In some cases a broadening was detected initially. This may reflect changes in the *C1 lC, equilibria, but in view of the errors involved in these measurements we prefer not to attempt any detailed interpretation. We thank Dr W. Derbyshire for helpful discussions and the S.E.R.C. for a grant to J.A. B. J. M. Harvey, R. J. Naftalin and M. C . R. Symons, Nature (London), 1976, 261, 435. M. C . R. Symons, J. A. Benbow and J. M. Harvey, Carbohydr. Res., 1980, 83, 9. S. J. Angyal, Aust. J . Chem., 1972, 25, 1957; Pure Appl. Chem., 1973, 35, 131. R. E. Lenkinski and J. Reuben, J . Am. Chem. SOC., 1976, 98, 3089. R. N. Butler and M. C. R. Symons, Trans. Faraday SOC., 1969, 65, 945; 2559. J. Davies, S. Ormondroyd and M. C . R. Symons, Trans. Faraday SOC., 1971, 67, 3465. S. J. Angyal, Chem. SOC. Retl., 1980, 9, 415. * J. M. Harvey and M. C. R. Symons, J . Solution Chem., 1978, 7, 571. a M. C. R. Symons and J. Davies, J. Chem. SOC., Faraday Trans. 2, 1975, 71, 1037. lo M. C . R. Symons, Acc. Chem. Res., 1981, 14, 179. (PAPER 2/595) 119 FAR 1

ISSN:0300-9599    

DOI:10.1039/F19827803671

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 78, 3679-3689 CH(A 2A-X "n) and OH(A T+-X ") Chemiluminescent Radiation from q3P) + C6H6 Discharge-Fast-flow Mixtures BY FRANCISCO TABARES, VINCENTE SAEZ RABANOS AND ANGEL GONZ~LEZ UREGA* Departamento de Quimica Fisica, Facultad de Ciencias Quimicas, Universidad Complutense de Madrid, Ciudad Universitaria, Madrid-3, Spain Received 14th April, 1982 Using a discharge-fast-flow apparatus a study of the CH* and OH* chemiluminescent radiation produced in O(3P) + C,H, reactive mixtures has been carried out. CH* and OH* emission intensities centred at 43 15 and 3090 8, have been assigned to the CH(A2A-X2H) and OH(A2C+-X2H) electronic transitions, respectively, and monitored as a function of reactant concentrations for several temperatures. The following empirical relations for their emission intensities have been obtained : Z(CH*) = a,[B] exp (- b,[B]) and I(OH*) = a2[BI2 exp (- b,[B]) where [B] is the concentration of benzene and Q, and b, (i = 1,2) are empirical factors independent of [B].An absolute kinetic determination of the CH* chemiluminescent reaction, using the NO+O glow as a standard, has been determined, leading to an estimate of the quantum yield of d(CH*) = 0.06. 1. INTRODUCTION The major contribution of fast-flow discharge systems has been the measurement to high precision of the rate constants of elementary gas-phase reactions where atoms and radicals produced by means of radiofrequency or microwave discharge at low pressure are involved. Excellent reviews of fast-flow chemical reactions have been reported1* covering most of the literature.On the other hand modulation techniques3 constitute one of the simplest and most elegant methods of studying fast reactions. Its main use has been in the investigation of photoinduced additive polymerization (the rotating-sector method), although recently several a u t h o r ~ ~ - ~ have extended its use to atom-polyatom gas-phase reactions with organic compounds. Thus Hunziker4 has made a number of studies of the absorption spectra of triplet species formed by energy transfer from Hg(3P1) or Hg(3P0) by using rapidly modulated mercury resonance radiation. A similar modulation system has been used by Atkinson and Cvetanovic5 to study reactions of ground-state oxygen atoms produced by the 253.7 nm mercury- sensitized N20 decomposition in a flow system.In this study we describe a modulated-discharge-fast-flow reactor suitable for kinetic data determination of elementary reactions involved in more complicated processes, such as combustion, atmospheric and photochemical reactions, and we report the CH(A 2A-X "n) and OH(A 2C+-X "n) emissions observed in O(3P) + benzene mixtures. The kinetics and dynamics of reactions between oxygen atoms and hydrocarbons have been studied widely, and excellent reviews have been published.6 In particular, benzene as the hydrocarbon is of great interest in combustion processes and atmospheric chemistry, and this system has been studied by a variety of techniques. Thus, relative reaction rates for the O(3P) + benzene reaction have been determined using static 3679 119-23680 q 3 P ) + C6H6 DISCHARGE-FLOW MIXTURES photolysis techniques. 7 9 Also, absolute rate-constant measurements have been reported using pulse radioly~is,~ discharge-flow,1°-12 modulation phase-shifP? l3 and flash-photolysis-N02-chemiluminescence14 techniques.Recently the molecular-beam methodl5* l6 (and particularly the crossed molecular- beams investigation of Lee and coworkers16) has identified the first two major reaction pathways as (1) (2) the first channel being hydrogen elimination and the second oxygen addition leading to the formation of highly excited phenol. Except for the above-mentioned molecular-beam study, the identification of products in the O(3P)+C6H, system is complicated by the presence of viscous, non-volatile reaction products (probably polymeric in character).At the present, phenol, CO, C5H6 and C,H,O (but not CH) have been reported1'* 1 5 9 l6 among others as O+C,H, reaction products. CH emission from the A2A-X211 system has been observed mainly in oxyacetylene flames. On the other hand ultraviolet OH emission corresponding to the electronic A 2X+-X 211 transition has been observed in a variety of systems, such as hydrogen,17 hydrocarbon flames18 (particularly oxya~etylene,'~ shock tubes,20 in the reaction between ground-state 0 and H atoms at room temperature21 and also as a component in the spectra of discharge systems containing water vapour22 even in trace amounts. In the present paper we focus attention on trying to obtain maximum information on these chemiluminiscent processes in terms of band identification, dependence of the emission upon the reactant concentrations and temperature, and the absolute kinetic determination for these chemiluminiscent reactions using the NO + 0 glow as a These results are presented and discussed, with our main goal being to report such findings rather than to provide a detailed mechanism which would perhaps require more work, particularly directed toward identifying intermediates or precursors, because of the complexity of reactions occurring in the system.q 3 P j ) + C6H6 + C6H50 + H AHg8 = - 15.9 kcal m01-l q 3 P j ) + C,H, + C6H50H AHg8 = - 102.4 kcal mol-l 2. EXPERIMENTAL DISC H AR GE-G AS-FL 0 W RE ACTOR The apparatus2* is shown schematically in fig. 1. A Pyrex reactor (2 cm diameter, 50 cm length) is pumped by means of a strong mechanical pump producing a fast-flow velocity of several thousand cm3 s-l.The flow system is a conventional one with gas flows controlled by needle valves and using a calibrated 'floating ball' flow meter (Goring S.A.) to measure the organic compound. A typical benzene flow range for the present experiment is from 2.2 to 1 1.2 pmol s-l. The main gas flow (0.7-1.3 mmol s-l of Ar and 70-140 pmol s-l of 0,) passes through a modulated discharge cavity powered by a microwave unit (Electromedical Supplies, Microtom 200 MK2) operated at 100 W, in which the O(3P) atoms are produced before they enter the flow tube. The Pyrex reactor wall is coated with boric acid to minimize oxygen-atom recombination. Since the oxygen atoms in the present experiment are produced by modulated discharge any emission of products would appear modulated at the same frequency as the discharge.As the reactants are introduced into the reactor inlet the reaction between oxygen and the organic compound takes place, and any chemiluminescent product formed is detected by measuring its emission intensity. As shown in fig. 1 we monitored the signal via a monochromator (Jobin Ybon H20) coupled to a photomultiplier (RCA C3 1034) connected to a preamplifier/lock-in amplifier system (Keithley models 427 and 840, respectively) and then to an oscilloscope (Tetronix 5403) or strip-chart recorder (Leeds & Northrup, Spedomax W). After each emission band was identified we recorded its intensity as a function of reactant concentration and temperature, the latter being varied by a furnace allocated to the reactorF.TABARES, v. SAEZ RABANOS AND A. GONZALEZ URERA 368 1 MONOCROMATOR PHOTOMULTIPLIER BENZENE FLOWMETER - FIG. 1 .-Schematic diagram of the discharge-fast-flow reactor and monitored by several thermocouples placed inside the flow tube. Therefore, despite the fast-flow velocity (see below) of our gaseous mixture care was taken to insure both temperature homogeneity and that the observed temperature was that of the actual reactive mixture. FLO W-VE LOC I TY ME A S UREME N T S In the present experiment a time-of-flight determination of the flow velocity was adopted. In this method we produced the O(3P) atoms by a short (ca. 1 ms) discharge pulse and measured the appearance time of the NO,* fluorescence intensity produced by the reaction O+NO+M+NO,*+M when NO was introduced in a particular reactor inlet (dl).The same procedure was repeated further downstream (d,). Thus under our experimental conditions the difference in the appearance time A? can only be attributed to the gas velocity, 2, = Ad/At. Fig. 2 shows typical time-of-flight data giving a value tl = 4000 & 100 cm s-l. Apart from the experiments described in section 3.1, carried out 1.0 ms into the reaction, all the present experiments were carried out at 7.0 ms. t/ms FIG. 2.-Typical time-of-flight data for flow-velocity determination. MATERIALS Benzene, toluene, cyclohexane, cyclohexene and n-hexane (Merck, purity 99.5 %) were purified by low-temperature distillation and subsequent degassing at 77 K.All the gases were from S.E.O. and were of the following purity: Ar > 99.9%; 0, > 99.98%; NO > 99.9% and N,O > 99.9%. They were used directly.3682 o(3~) + C,H, D IS c HA R G E--F LO w M I x T u R ES 3. RESULTS AND DISCUSSION 3.1. OBSERVED EMISSIONS During the present experiments reactive mixtures of O(3P) +benzene were prepared and several emission bands, illustrated in table 1, were identified and recorded. Typical experimental conditions are also reported. For the benzene reaction a low-resolution spectrum of the most important bands is shown in fig. 3. The relative1 intense band CH(A 2A-X211) electronic transition, as reported in ref. (19c) and (19d). Although from 4240 to 4380 A, centred at 4315 A, was identified as the 4315 K band for the TABLE 1 .-EXPERIMENTAL CONDITIONS AND RELATIVE BAND INTENSITIES IN THE q 3 P ) + C,H, REACTION experimental conditions gas flow/pmol s-l total pressure/Torr temperature / K discharge frequency/Hz discharge power/W reaction time/ms photomultiplier voltage/V typical signal/pVa typical noise/pVa monocromator resolution/nm Ar 1300 0 2 140 'GH6 11.2 2.0 42.5 7.0 333 100 1800 ( 2 x 10-3)- 10-2 10-4 2 band intensities CH A2A-X211 4315 1 .ooo CH B2A-X211 3895 0.062 OH A2E-X211 3090 0.765 a The lower and higher values are shown, depending on the benzene concentration range. our resolution was not adequate to reveal the rotational distribution of the electronically excited CH(A 2A) species, the low shoulder observed around 4335 A, in agreement with more precise observations from the chemiluminescence of CH in the 0 + C2H, flame,lSd can be attributed to a P-branch with a rotational distribution characterized by Got = lOOO+ 100 K (estimated from a wavelength shift of AII x 20 A with respect to the main peak, i.e.the Q-branch). The 3060-3130 A band centred at 3090 A was ascribed to the OH(A 2Z+-X 211) transition as reported by Krishramachi and Broida in their study on the emission spectra of atomic-oxygen-acetylene flames.lgC A very low intense band around 3890 A observed in the 0(3P)+benzene reaction was attributed to the CH (B2A - X 2 n ) transition, as described e1~ewhere.l~~ The observed band intensities, corrected for the instrument's response, were normalized, with the intensity of the 43 15 A band of CH(A 2A-X ") taken as unity.These emissions (but with changes in intensity) were observed in the following cases: (a) when the 0 atoms were produced either by Ar/N,O discharge or by the N+NO + N 2 + 0 reaction, indicating that for the observed emissions no 0, is necessary; (b) when toluene was used instead of benzene, under the same experimental conditions.F. TABARES, V. SAEZ RABANOS AND A. G O N Z ~ L E Z U R E ~ A 3683 OH(A ’ 1 - X *n) 2 L I I 1 I I I I 1 300 305 310 315 ” 420 425 430 435 440 h/nm FIG. 3.-Low-resolution emission spectrum obtained in the O(3P) +C,H, reactive mixture under the experimental conditions of table 1 . No emission bands at 4315 and 3895 A, i.e. the A and B electronic states of CH, were observed when benzene was replaced by n-hexane, cyclohexane or cyclohexene.This could indicate that triply bonded hydrocarbons, formed only in the benzene reaction (see section 3.4), are needed as CH precursors in the complex reaction. 3.2. OH* AND CH* EMISSION INTENSITIES AS A FUNCTION OF BENZENE CONCENTRATION AT A FIXED INITIAL CONCENTRATION OF OXYGEN ATOMS FOR SEVERAL TEMPERATURES Fig. 4 and 5 show the benzene concentration dependence of both CH* and OH* emission intensities at a fixed initial concentration of oxygen atoms for several temperatures. These data are well represented by the following empirical relations : I(CH*) = a,[B] exp (- b,[B]) I(OH*) = a,[BI2 exp (- b2[B]) (2) where [B] is the benzene concentration and ai and bi ( i = 1,2) are parameters dependent on temperature but not benzene concentration. In fact this dependence, at least for the b parameters, shows Arrhenius-like behaviour (see fig.6), i.e. bi x Ai exp (- Ci/RT). From the slopes in fig. 6 the following Ci values were obtained: C, (CH* emission) = 4.4f0.3 kcal mol-l and Cz (OH* emission) = 2.1 kO.2 kcal mol-l. The overall temperature range was 328-474 K. Note that the present temperature depen- dence of the CH* emission is similar to the activation energy1’ for the corresponding initiation reaction of 0 + C,H, given by Ea = 4.4 f 0.5 kcal mol-l; this indicates that even though complex secondary reactions of 0 plus intermediates are occurring at least the CH* chemiluminescent radiation is limited by the primary reaction. on the CN(B2X - X 2 n ) emission intensity in mixtures of active Previous3684 q3P) + C6H6 D ISCH ARGE-FLO W MIXTURES I I 1 5 10 15 [ B ] /lo-'* rnol cm-3 FIG.4.-Logarithmic I/[B], i.e. CH* emission intensity divided by benzene concentration, [B], as a function of [B] for several temperatures: 0, T = 445 K; 0, T = 350 K; 0, T = 417 K; A, T = 383 K ; solid lines are smooth lines through the data. nitrogen and carbon tetrachloride have shown that empirical relationships similar to that of CH emission, i.e eqn (l), can reasonably be described by a complex mechanism where the initiation reaction is the rate-limiting step. All the present experiments were carried out at an initial concentration of oxygen atoms of [O], < 2 x mol ~ m - ~ . By adding a constant flow of NO to the 0 +benzene reactive mixture through the reactor inlet located in front of the monocromator/photomultiplier zone we measured the NO,* emission intensity due to the residual concentration of 0 atoms.Fig. 6 shows the CH* to NO,* emission intensity ratio as a function of benzene concentration. These data are quite well represented by a straight line with zero intercept, from which (under the present conditions) one may conclude that the CH* emission intensity is directly proportional to the product of benzene and oxygen concentrations, i.e. Z(CH*) oc [B] [O]. The same data representation (not shown in the figure) gives a more complicated dependence for the OH* intensity. This is perhaps an indication of a different reaction mechanism, as one might expect from the different value observed for its temperature dependence.3.3. KINETIC DATA DETERMINATION OF THE CH* A N D OH* CHEMILUMINESCENT REACTIONS. ESTIMATION OF QUANTUM YIELDS The kinetic rate constants of the chemiluminescent processes 0 + C6H6 + OH* + CH* +products + ow + CH + hv +productsF. TABARES, v. SAEZ RABANOS AND A. GONZALEZ UREGA 3685 0 [Bl/lO-'O mol ~ r n - ~ FIG. 5.-As fig. 4 but for OH* emission in the benzene reaction; the ordinate scale is now log I/[BI2: 0, T = 333 K; 0, T = 378 K; 0, T = 408 K; A, T = 438 K; ., T = 463 K. emission. Solid3686 q 3 P ) + C6H6 D I S C H A R G E-F L 0 W M I X T U R E S were estimated by using the NO+O glow as a standard reaction.23 First, and as suggested by Fontijn et aZ.,23 the unknown CH* glow and the NO+O glow were viewed under the same conditions (i.e.using the same apparatus, physical location, detector, etc.) and their respective intensities I NO,*) and I(CH*) were recorded over displayed in fig. 3. Next the ‘specific’ intensity I ( X ) was defined as the same wavelength range from 4240 to 4380 8, , corresponding to the CH* spectrum where [XI is the concentration of NO or C6H6 and [o] is the oxygen-atom concentration. To insure the vality of the present calculation care was taken to correct for small changes in the oxygen-atom concentration in both measurements. This was accom- plished by measuring the NO,* fluorescence intensity at 3, = 5500 A with and without C,H, and including this correction factor in our calculation. Then the rate constant of the two light-emitting reactions over the wavelength region A1-A2 (as indicated by Fontijn et al.) was calculated via the equation The value of k(NO,*) used here was the reported value of 6.4 x lo-’’ cm3 molecule-’ s-l, corrected by considerin the fractional area of the wavelength interval of interest (i.e.from 4240 to 4380 R ) over the whole NO,* glow spectrum extended from 3875 to ca. 14000 A. By using this procedure a k(CH*) value of 7.45 x lo9 cm3 mol-l s-l was obtained. Since the OH* glow lies outside the NO,* glow wavelength range we have normalized our OH* spectra to those of CH* under the same experimental conditions, their relative emission intensities being those shown in fig. 3. This simple scaling procedure shows that typically I(CH) > I(OH*), and could indicate (see fig. 3) a smaller quantum yield, i.e.&OH*) < q5(CH*), but care should be taken with this comparison since the OH* emission cannot be reduced (see below) to the ‘bimolecular’ scheme, i.e. I(OH*) z [B] [O], as the CH* emission does, and therefore it cannot be compared with the standard NO+O glow. In other words the respective quantum yield for the OH* emission is a complex function [see eqn (2)] of the concentrations of both benzene and atomic oxygen. Thus under our experimental conditions only q$(CH*) can be properly defined as #(CH*) = k(CH*)/k, where k is the rate constant for the primary reaction, i.e. O+C,H, --+ products. Now if one uses for k the mass-spectrometric value of Bonano et al.,ll obtained by monitoring the benzene concentration and given by 1.22 x loll cm3 mol-l s-l at T = 383 K, one obtains a fluorescence quantum yield of b(CH*) = 0.06, i.e.for every 100 benzene molecules 6 give fluorescence emission via the CH(A 2A-X 211) electronic transition. 3.4. POSSIBLE SOURCES FOR THE FORMATION OF CH* AND OH* In spite of the fact that our main goal, as mentioned before, is not to give a detailed mechanism for either OH* or CH* formation, the present experimental results together with previous studies2,? 27 provoke some comments on the most likely pathways for the formation of these radicals. These are as follows. CH* FORMATION A molecular-beam studyl59 l6 of the same reaction (0 + C6H6) has identified both CO and C5H6 (cyclopentadiene or 3-pent-1-yne) as products of the excited decom- position of phenol described in eqn (1). Indeed, the cyclopentadiene + oxygen-atom reaction has been shown2, to originate C,H, isomers, one of the major products beingF.TABARES, v. SAEZ RABANOS AND A. GONZALEZ UREAA 0.6 n E v) *-' .3 0.4 .f' 0" 5 v h 4 u 4' 0.2- 3687 - - I 1 [Bl/lO-'O mol cm-3 FIG. 7.-Z(CH*)/Z(NO,*) plotted against [B]. Both intensities are monitored simultaneously at 7 ms reaction time and at 340 K. Note the zero (within experimental error) intercept. but- 1 -yne, which upon reaction with oxygen atoms27 produces CH-C-CH, and CHECH. On the other hand if one assumes 3-pent-l-yne, rather than cyclopentadiene, as the C,H6 product of the C6H6 + 0 reaction, it also seems likely that acetylene is formed.27 Finally, evidence for the formation of CH* in O+C,H, mixtures has been given by Broida and Since all these intermediate reactions have rate constants faster than that of the primary reaction (0 + C6H6) we believe that under the present experimental conditions CH* could be formed via the sequence unsaturated hydro- carbons -+ acetylene -+ CH*; however, we reiterate that this can only be considered as a reasonable suggestion rather than a complete mechanism.OH* FORMATION OH* emission has also been observed in the O+C,H, reaction,lg and path CH*+O, -+ CO+OH* has been proposed as a major source of OH*. In the present study the OH* mechanism seems to differ from that of CH*, as shown by the following facts: (i) OH* emission is also observed when 0, is excluded and 0 produced from N,O; (ii) the empirical law found for its intensity evolution shows a different dependence on benzene concentration than in the case of CH*.4. CONCLUDING REMARKS The present study has focussed attention on the CH(A 2A-X ") and OH(A 2Z-X ") emissions observed in the O(3P) + benzene reactive mixtures produced in our modulated microwave-discharge-fast-flow reactor. We have drawn attention to the identification of bands, the dependence of the emission upon temperature and reactant concentrations and to absolute kinetic determinations for these chemiluminescent reactions by using the NO+O glow as standard. The main results of the present study are as follows. (a) In the discharge-flow q 3 P ) + C6H6 reaction, and under the present conditions, chemiluminescent radiation centred at 43 15 and 3090 A has been observed associated with the CH(A ,A-X ") and OH(A 2C+-X "n) electronic transitions.A very low intense band around 3890A observed in the benzene reaction is attributed to the CH(B2A-X211) transition. Both CH and OH emissions are present when the O(3P)3688 q3P) + C6H6 DISC H A RG E-F LO W MIXTURES atoms are produced either by an Ar/N,O discharge or by the N + NO -+ N, + O(3P) reaction, indicating that 0, for the observed emission is not essential. No CH* emission bands were observed when benzene was replaced by n-hexane, cyclohexane or cyclohexene. This is thought to indicate that unsaturated triple-bonded hydrocarbons are required as CH(A ,A) precursors. Also low-resolution data analysis has shown that the CH(A) state is formed with a rotational temperature of Tot = 1000 100 K. (b) Among the main findings are the empirical laws [eqn (1) and (2)] for the dependence of the CH* and OH* emission intensities upon temperature and reactant concentration.It seems that two different mechanisms are responsible for each chemiluminescent process. Whereas CH* formation appears to be limited by the initiation reaction (despite the presence of complex secondary and tertiary reactions) and shows the same temperature dependence as does the primary reaction, OH* formation seems to be a more complicated process in which subsequent steps with benzene participation could be equally important and be required to account for the observed emission behaviour. (c) A comparison of the chemiluminescent radiation with the standard NO + 0 glow discharge gave an absolute rate constant value of k(CH*) = 2.45 x lo9 cm3 mol-1 s-l, from which a quantum yield of @ = 0.06 was estimated, i.e. 6% of the benzene molecules give, upon reaction with atomic oxygen, fluorescence emission via the CH(A ,A-X ,lT) electronic transition; we consider this to be quite considerable since there are so many reactions occurring in our system. Although the OH* emission could not be reduced and compared with the NO+O glow discharge some indication was reported that, at least, its quantum yield (a complex function of both benzene and oxygen-atom concentrations) was not higher than that of CH*.Finally, we realize the complexity of the reaction system (under gas-phase conditions) being studied, and we reiterate that our main purpose has been to report several items of information concerning these chemiluminescence processes.More experimental and theoretical work is needed, in particular some direct identification of the assumed OH* and CH* precursors via (perhaps) laser-induced fluorescence. This type of study (to include modelling) is in progress in our laboratory. We thank the staff of the Universidad Complutense mechanics, glass and electronics workshop for their help. F.T. and V.S. acknowledge F.P.I. and I.N.A.P.E. fellowships. We also acknowledge the constructive criticism and valuable suggestions of the referees. J. Wolfrum, Atom Reactions, in Physical Chemistry: an Advanced Treatise, ed. H. Eyring (Academic Press, New York, 1975), vol. VI B. (a) A. A. Westemberg, Annu. Rev. Phys. Chem., 1973,24,77 (b) J. T. Herron and R.E. Huie, J. Phys. Chem. Ref. Data, 1973, 2, 467. L. F. Phillips, Prog. React. Kinet., 1975, 7, 83. H. E. Humzinker, IBM J. Res. Dev., 1971, 15, 10. (a) R. Atkinson and R. J. Cvetanovic, J. Chem. Phys., 1971, 55, 659; 1972, 56, 432; see also (6) R. Atkinson and J. N. Pitts, J . Phys. Chem., 1974, 78, 1780; 1975, 79, 295. (a) R. E. Huie and J. T. Herron, Prog. React. Kinet., 1975,8, 1 ; (b) M. C. Lin, Dynamics of Oxygen Atoms Reactions in Potential Energy Surfaces, ed. K. P. Lawley (Wiley, New York, 1980); (c) I. W. M. Smith, The Excited State in Chemical Physics (Wiley, New York, 1975). E. Grovenstein Jr and A. J. Mosher, J . Am. Chem. SOC., 1970, 92, 3810. I. Mani and M. C, Sauer Jr, Adv. Chem. Ser., 1968, 82, 142. 1, 28. 7 G. Bocock and R. J. Cvetanovic, Can.J. Chem., 1961, 39, 2436. lo L. I. Avramenko, R. V. Kolesnikova and G. I. Savinova, Isv. Akad. Nauk SSSR, Ser. Khim., 1965, l1 R. A. Bonamo, P. Kin, J. H. Lee and R. B. Timmons, J. Chem. Phys., 1972, 57, 1377.F. TABARES, v. SAEZ RABANOS AND A. GONZALEZ U R E ~ ~ A 3689 l2 D. Furuyama and N. Ebara, Znt. J. Chem. Kinet., 1975, 7, 289. l3 A. J. Colussi, D. L. Singleton, R. S. Irwin and R. J. Cvetanovic, J. Phys. Chem., 1975, 79, 1900. l4 R. Atkinson and J. N. Pitts Jr, Chem. Phys. Lett., 1979, 63, 485. l5 T. M. Sloane, J. Chem. Phys., 1977, 67, 2267. l6 S. J. Sibener, R. J. Buss, P. Casavecchia, T. Hirooka and Y. T. Lee, J. Chem. Phys., 1980, 72,4341. l 7 (a) M. Charton and A. G. Gaydon, Proc. R. SOC. London, Ser. A , 1958, 245, 84; ( b ) W. E. Kaskan, J . Chem. Phys., 1959, 31, 944; (c) H. P. Broida and K. E. Shuler, J. Chem. Phys., 1952, 20, 168. T. Carrington, J. Chem. Phys., 1957, 30, 887. l9 (a) A. G. Gaydon and H. G. Wolfhard Proc. R. SOC. London, Ser. A , 1951,208,63; ( b ) T. Carrington, J . Chem. Phys., 1959,30, 1087; (c) S. I. N. G. Krishnamachari and H. P. Broida, J. Chem. Phys., 1961, 34, 1709; ( d ) K. A. Quickert, J. Phys. Chem., 1972, 76, 825. 2o F. E. Belles and M. R. Lauver, J. Chem. Phys., 1964, 40, 415. 21 S. Ticktin, G. Spindler and H. 1. Schiff, Discuss. Faraday SOC., 1967, 44, 218. 22 A. Michel, 2. Naturforsch., Teil A , 1957, 12, 887. 23 A. Fontijn, C. B. Meyer and H. I. Schiff, J. Chem. Phys., 1964, 40, 64. 24 A preliminary description of the apparatus can be found in A. Gonzalez Ureiia, F. Tabares and 25 V. Saez Rabanos; F. Tabares and A. Gonzalez Ureiia, J. Photochem., 1982, 78, 301. 26 Kazumoto Nakamura and Sehchiro Koda, Int. J. Chem. Kinet., 1979, 9 67. 27 P. Herbrechtsmeier und H. Gg. Wagner, Ber. Bunsenges. Phys. Chem., 1975, 79, 461. V. Saez Rabanos, Physico-Chemical Hydrodynamics, Europ. Conf. Abs., 1980, 3F, 114. (PAPER 2/624)

ISSN:0300-9599    

DOI:10.1039/F19827803679

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1982,78, 3691-3692 Solid-state Studies Part 26.-Raman Spectroscopic Evidence for a Phase 11-like Intermediate during the Course of the IV-I11 Phase Transition in Ammonium Nitrate BY GORDON J. KEARLEY AND SIDNEY F. A. KETTLE* School of Chemical Sciences, University of East Anglia, Norwich NR4 7TJ Received 23rd April, 1982 Raman evidence is presented which supports the view that the intermediate occurring in the IV + I11 phase transition of ammonium nitrate has a structure similar to that of phase 11. The room-temperature, stable form of ammonium nitrate, phase IV, converts into phase I11 at ca. 32 OC. This transformation is of technical importance because if a sample is repeatedly cycled through the transition, crystalline samples crumble, giving rise to a variety of problems.The crystal structure of both phase IV1 and phase 1112 are known and bear few similarities to each other. For instance, the nitrate-ion molecular planes are parallel in phase IV but in phase I11 these ions lie in planes at ca. 90'. In addition, on proceeding from phase IV to phase I11 there is an increase in volume which is calculated to be 4.6% from unit-cell dimensions1*2 but found to be 3.6% by density determination^.^ This discrepancy may be due to expansion of the lattice into voids which are commonly found in crystallites of phase IV or alternatively may reflect local compression and lattice distortions in phase I11 arising from the phase transition. It has recently been reported that a transient intermediate phase exists in the IV-I11 tran~formation.~* X-ray data have been recorded for this transient4 and it has been suggested that it has a structure akin to that of phase I1 (a phase stable between 82 and 125 OC).Since this report we have obtained photographic evidence for the existence of this inte~nediate;~ in our hands it has proved to be stable for a maximum of ca. 1 min. There may be considerable hysteresis on the IV-I11 transition; the higher the temperature at which it occurs the longer the period of stability of the intermediate. Thus, the ca. 1 min stability period was obtained for a sample for which the transition occurred at ca. 45 OC. This limited temporal stability is such as to preclude complete spectral coverage with the conventional apparatus available to us.We have therefore studied the Raman spectrum of the intermediate in the 720 cm-l region, a region in which phase 11 has a very different spectrum to phase IV. Although complete Fourier-transform infrared studies are possible and superficially attractive, many infrared bands are broad and not useful for characterisation of the intermediate. Although the infrared band at ca. 720 cm-l could be used for this purpose, it has considerable intensity and structure in phases IV and 111; very strangely it is of negligible intensity in phase 11. We now report spectroscopic evidence that the intermediate occurring between phases IV and I11 resembles phase 11. In fig. 1 we show Raman spectra recorded over a range of ca. 30 cm-l of the nitrate-ion in-plane deformation mode, v4, recorded as a sample passed through the 111-11 transition, the spectrometer scan direction being reversed at the points indicated by arrows.We have recorded similar data for the other transitions and for the IV-I11 transition, on rare occasions, have obtained a spectrum 369 It > E e: 4- .C U .- !6Y2 IV-111 PHASE TRANSITION I N AMMONIUM NITRATE I crn-' m 71 7 I Lp 71 6 100 7 5 50 46 45 44 T/'C T/'C FIG. 1. FIG. 2. FIG. 1 .-Consecutively recorded Raman spectra of the 720 cm-' region of poiycrystalline ammonium nitrate illustrating the spectral consequences of a 111-11 transition. Whilst the sample was heated from 50 to 100°C the spectrometer was scanned over a ca. 50cm-' range eight times. The spectrometer scan direction was reversed at the points designated by arrows so that a wavelength decreasing scan is followed by one with wavelength increasing.Temperature calibration is approximate only. FIG. 2.-Spectra recorded as in fig. 1 but at a lower temperature illustrating a two-step IV-I11 transition occurring at ca. 45 O C with the intermediate giving rise to a doubled v4 feature analogous to that obtained for phase I1 (see fig. 1). Arrows indicate reversal of spectrometer scan direction. which must be that of the intermediate. The best example is shown in fig. 2, unfortunately run at a different chart speed to fig. 1 but over a similar frequency range, where the single v4 features in phases IV and I11 are to be contrasted with the double feature of the intermediate. In both frequency and intensity pattern the spectrum of the intermediate resembles that of phase I1 (fig.l), consistent with the two phases having very similar structures. Because of the very different nitrate-ion orientations in phases IV and I11 it seems probable that in any intermediate they would be free to reorientate from the one orientation to the other; note that in phase I1 the nitrate ions are disordered.6 EXPERIMENTAL Raman spectra were recorded using a Spex 1401 double monochromator, an Ortec photon counting system 5C1 and a Spectra Physics Ar+/Kr- laser. The exciting line used was 20490 cm-l and the laser power, measured at the sample, was ca. 30 mW. Polycrystalline samples of AnalaR grade ammonium nitrate were sealed in capillary tubes and temperature control was achieved using a modified microscope hot-stage arrangement. G. J. K. thanks the S.E.R.C. and Fisons Ltd for financial support. C. S. Choi, J. E. Mapes and E. Prince, Acta Crystallogr., Sect. B, 1972, 28, 1357. B. W. Lucas, M. Ahtee and A. W. Hewat, Acta Crystallogr., Sect. B, 1980, 36, 2005. S. B. Hendncks, E. Posnjak and F. C. Kracek, J. Am. Chem. SOC., 1932, 54, 2766. W. Engel and P. Charbit, J. Therm. Anal., 1978, 13, 275. G. J. Kearley and S. F. A. Kettle, J. Chem. SOC., Faraday Trans. I , 1982, 78, 1817. B. W. Lucas, M. Ahtee and A. E. Hewat, Acta Crystallogr., Sect. B, 1979, 35, 1038. (PAPER 2/676)

ISSN:0300-9599    

DOI:10.1039/F19827803691

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1982, 78, 3693-3102 Volumetric, Dielectric and Transport Properties of Some Liquid Tri-n-alkylamines BY CVETO KLOFUTAR,* SPELA PALJK AND RADISLAVA MALNERSIC ' J. Stefan' Institute, ' E. Kardelj' University of Ljubljana, 61000 Ljubljana, Yugoslavia Received 26th May, 1982 Densities (293.15 d T/K d 333.13, electric permittivities (293.15 < T/K < 323.15) and viscosities (293.15 d T/K d 333.15) of some tertiary n-alkylamines were measured. The values of the density and electric permittivity at 293.15 K and the coefficients of thermal expansion, a, and the temperature dependence of the electric permittivity, y , are given. The dependence of a (or y ) on the number of carbon atoms in the molecules of the liquids investigated are presented. The Onsager and Kirkwood relations were used to calculate the dipole moments. It was found that within the temperature range studied, the dipole moments are constant and decrease with increasing number of carbon atoms in the alkylamine chains.The values obtained for the Kirkwood correlation coefficient are slightly higher than unity, showing a tendency to increase from triethylamine to tri-n-dodecylamine. To describe the viscosity data, transition-state theory as well as free-volume theory were applied. The values of the energy of activation for viscous flow at constant pressure and the expansion energy of tertiary n-alkylamines were calculated. In addition, triethylamine was classified as a volume-restrained liquid, while the other amines investigated behaved as energy-restrained liquids.Tertiary n-alkylamines are important organic bases. With the exception of tri- methylamine, they are either liquids or low-melting-point solids. Their solubilities depend on the length of the alkyl chain. The lower members of a homologous series are highly soluble in polar and non-polar solvents, while the higher members are soluble only in non-polar solvents. As a consequence of steric hindrance, the basicity of tertiary n-alkylamines decreases with increasing number of carbon atoms per chain.' Liquid tertiary n-alkylamines are weakly polar, non-associated, strong proton acceptors. The molecular properties of these liquids are not known extensively and the information available on their dielectric, volumetric and transport properties is limited. In this paper we report a study of some liquid tertiary n-alkylamines with up to twelve carbon atoms per alkyl chain and show how some bulk properties depend on the molecular complexity and size.EXPERIMENTAL Triethylamine (TEA, Fluka A.G.), tri-n-butylamine (TBA, Fluka A.G.), tri-n-hexylamine (THxA, Fluka A.G.), tri-n-octylamine (TOA, Fluka A.G.), tri-n-decylamine (TDA, Eastman Kodak) and tri-n-dodecylamine (TLA, RhGne-Poulenc) were purified and analysed as in ref. (2). The density measurements in the temperature range studied were performed as in ref. (2). Electric permittivity measurements were carried out using a WTW dipole meter DM 0 1 with a DFL 1 cell at a constant frequency of 2 MHz. The cell was thermostatted at 293.15, 298.15, 303.15, 313.15 and 323.15 K by a water bath to within f0.05 K.The calibration of the cell was performed as in ref. (3) using carbon tetrachloride, benzene and cyclohexane as standard^.^ Viscosity measurements were performed with various viscometers of the Cannon-Fenske type 36933694 PROPERTIES OF LIQUID TRI-n-ALKYLAMINES in the temperature range from 293.15 to 333.15 K. In the case of TEA and TBA, the absolute viscosity, (kg m-' s-l), was calculated by means of the equation5 Ed 'I = Cdt -- t 2 in which the kinetic energy correction is considered and where t is the flow time in s and C and E are constants characteristic of the viscometer. The values of C = 4.446 x lo-' (k 7 x m2 s were obtained from the calibration of the viscometer with water, using the data for the absolute viscosity and density of water from ref.(6) and (7), respectively. The absolute viscosities of THxA, TOA, TDA and TLA were calculated through the equation dt m2 s - ~ and E = 1.41 x ( + 4 x (2) ' I = f l 0 - where qo and do are the absolute viscosity and density of a standard solution. The absolute viscosities of THxA and TOA were determined using a viscometer calibrated with a 40 wt % aqueous sucrose solution. In the case of TDA and TLA, a viscometer was used. which was calibrated with a 60 wt % aqueous sucrose solution. The data for the absolute viscosity of sucrose solutions in the temperature range studied were taken from ref. (8). do 43 RESULTS AND DISCUSSION The values of the density, d, of the tri-n-alkylamines investigated as a function of (3) temperature are given by d = do(l -aAT) where AT = T - To, To is the reference temperature (in our case T, = 293.15 K), do is the density of the liquid at T, and a is the coefficient of thermal expansion defined as a = -(l/do)(c'd/c'7)p.In eqn (3) it is assumed that a is temperature independent within the temperature range studied. The values of a and do for tri-n-alkylamines, calculated from eqn (3) by the method of least squares, together with the standard error of the estimate, s, are given in table 1. It can be seen that the values of a increase TABLE VALUES OF THE DENSITY AT 293.15 K AND THE COEFFICIENTS OF THERMAL EXPANSION FOR THE TRI-n-ALKYLAMINES INVESTIGATED, WITH THE STANDARD ERROR OF EQN (3) compound do/kg dm-3 a/ 10-3 ~ - 1 S X 104 TEA 0.7288 f 0.0002 1.54 f 0.01 2.4 TBA 0.7774 f 0.0002 1.07 f 0.01 3.6 THxA 0.7980 0.0002 0.94 f 0.01 3.1 TOA 0.81 10f0.0002 0.90 f 0.01 3.3 TDA 0.8 190 & 0.0001 0.88 f 0.01 1.8 TLA 0.8241 f 0.0002 0.83 fO.01 3.3 with decreasing number of carbon atoms in the tri-n-alkylamines.The values of a obtained, from 0.8 x K-' (TEA), are characteristic of organic liquid^.^ The temperature dependence of the densities of the amines investigated is presented in fig. 1. The literature density data for TEAlO and TBAlO at 293.15 K, for TBA,12 THxA,~ TOA,2* l4 TDA2 and TLA2$ l5 at 298.15 K and for TEAlO and TBAlO at 313.1 5 K, and the thermal expansion coefficients for TEA and TBA, determined from density data at 293.15 and 313.15 K,1° are close to our values. K-l (TLA) to 1.5 xc. KLOFUTAR, 3.PALJK AND R. MALNERSIC 3695 08000 - - m I E 0 7500 - - Y . -Q -u r l 0 7000 - - u I I I I I t 0.6XIO I 1 I I I I I 0.00 10.00 20.00 30.00 4 0.00 ATlK FIG. 1.-Densities of TEA (O), TBA (m), THxA (A), TOA (A), TDA (0) and TLA (@) as a function of (7'- 293.15) K. In fig. 2 the coefficients of thermal expansion as a function of the number of carbon atoms, n, in tri-n-alkylamine molecules are given. The curve in fig. 2 is drawn on the basis of the equation b a =a+- n (4) with the regression coefficients a = 0.000 684 K-l and b = 0.005 039 K-l. The standard error of the estimate of eqn (4) amounts to s = 3 x lop5. The values of the molar volume, V (dm3 mol-l), of tri-n-alkylamines at 298.15 K as a function of the number of carbon atoms is given by V = (0.0412+0.0004)+(0.01653_+0.00002)~.( 5 ) The value of the slope of a plot of V against n is close to that for liquid n-alkanes up to n-dodecane.' This indicates that the methylene group in tri-n-alkylamines and n-alkanes occupies practically the same volume. The variation of electric permittivity, E , with temperature for the tri-n-alkylamines investigated is given by (6) E = ~ o ( 1 -yAT) where E, is the electric permittivity at and y is the average coefficient of temperature dependence of the electric permittivity, defined as y = - ( I / E , ) ( C ~ E / ~ T ) ~ . The values of E , and y, obtained by the method of least squares, together with the standard error of the estimate, s, are compiled in table 2. In fig. 3 the electric permittivity is given as a function of temperature.The lines in fig. 3 are drawn by means of eqn (6) using the values of E , and y from table 2. It is evident that the electric permittivities increase uniformly with decreasing temperature, albeit at different rates for different amines. The decrease of electric permittivity with increasing length of the alkyl chain in a tri-n-alkylamine at a definite temperature is due to the decrease in the number of dipoles per unit volume, nd. The plot of E against nd at 298.15 K, shown in fig. 4, is linear with an intercept value of 2.131 characteristic3696 PROPERTIES OF LIQUID TR I-n-A L K Y L A MINE S I I 1 I 1 16.0 7 12.0 2 U- M d 1 ?-- 8 .O 1.0 8 16 21 32 n FIG. 2.-a (-) and y (---) coefficients of TEA (a), TBA (a), THxA (A), TOA (A), TDA (0) and TLA (0) as a function of the number of carbon atoms in the molecule, n.TABLE 2.-vALUES OF THE ELECTRIC PERMITTIVITY AT 293.15 K AND THE COEFFICIENT OF THE TEMPERATURE DEPENDENCE OF THE ELECTRIC PERMITTIVITY FOR TRI-ll-ALKYLAMINES, TOGETHER WITH THE STANDARD ERROR OF EQN (6) compounds EO y / 10-3 ~ - 1 S X 103 TEA 2.440 f 0.001 1.71 kO.01 1.21 TBA 2.328 0.002 1.16 0.04 2.2 THxA 2.256 f 0.001 0.85 f 0.02 1 .o TOA 2.238 & 0.001 0.79 k 0.03 1.4 TLA 2.1 97 f 0.00 1 0.66 & 0.02 1.3 of the E value of liquid n-alkanes.' The electric permittivities of TBA and TLA at 298.15 K are close to those in ref. (16) and (17), respectively. The dependence of y on n for the tri-n-alkylamines investigated, presented in fig. 2, is given by B y = A + - n (7) where A and B are empirical coefficients.The value of A of 0.000466 K-l, of B of 0.007578 K-l and of s of 5 x were obtained by the method of least squares. For the amines investigated it follows that y increases with decreasing number of carbon atoms in the tri-n-alkylamine molecule.2.100 2.300 f 2.200 2100 c. KLOFUTAR, S. PALJK AND R. MALNERSIC 3697 I I I I I I I I 0.00 10.00 20.00 30.00 ATIK FIG. 3.-Electric permittivities of TEA (a), TBA (m), THxA (A), TOA (A) and TLA (0) as a function of (T- 293.15) K. 2.100 2.300 E 2.200 2.100 I I I I I I I I 9 i a 2 1 36 nd / 1 OZ3 dm-3 FIG. 4.-Electric permittivities of TEA (O), TBA (m), THxA (A), TOA (A), TDA (0) and TLA (0) as a function of the number of dipoles per dm3 at 298.15 K.3698 PROPERTIES OF LIQUID TRI-n-ALKYLAMINES The values of the molecular dipole moment, po, of the tri-n-alkylamines studied were calculated by means of the Onsager equation18 9kTM(2&+~,)(~+2) 4nN d 3 ~ ( ~ , + 2 ) p i = - - where k is the Boltzmann constant, T is the absolute temperature, N is the Avogadro constant, M is the molecular mass and E , is the optical permittivity.In calculating dipole moments of tri-n-alkylamines, the contributions of atomic polarization are neglected and therefore E, = rzk was used. The respective values of the refractive index, n,, are calculated from the Eykman constant2 and density data for the tri-n-alkylamines at definite temperatures. The values calculated for the dipole moments at the highest and lowest temperature studied differ by < 3.34 x C m for each of tri-n-alkylamines investigated. In table 3, therefore, the mean values of dipole moments, obtained by eqn (8), are given.TABLE 3.-MEAN VALUES OF THE DIPOLE MOMENT, /lo, CALCULATED FROM EQN (8) AND VALUES OF &/lo CALCULATED FROM EQN (9) AND ( 10) FOR TRI-n-ALKYLAMINES IN THE TEMPERATURE RANGE FROM 293.15 TO 323.15 K compound p,,/ 1 0-30 C m $po/l C m TEA TBA THxA TOA TDA TLA 2.54 2.54 2.37 2.37 2.24 2.07 2.87 3.17 3.30 3.54 3.67 3.81 In ref. (19) it was suggested that the dipole moments of tertiary n-alkylamines lie within the range 2.67 x C m. From table 3 it is evident that our values of the Onsager dipole moment, po, are lower and show a tendency to increase with decreasing number of carbon atoms in the tri-n-alkylamine molecule. However, the value of the dipole moment for tri-ethylamine is equal to that in ref.(20). On the other hand, the dipole moment of 9.25 x C m for TOA in benzenez1 is substantially lower than our value. As can be seen from ref. (22), the solvent effect raises the dipole moment of TEA by 1.67 x C m. Therefore, the observed difference of the dipole moment of TOA in benzene may be expected from the uncertainty inherent in analysing permittivity data in cases where the solute and solvent have nearly the same permittivities. The Kirkwood equation for the dipole moment, gip, taking into account the hindrance of molecular orientation by neighbouring molecules, is given by18 to 3.01 x gp2 = - - 4nN d (9) where p is the actual dipole moment and g is the correlation parameter. To determine the value of g the structure of the liquid and the short-range intermolecular forces are required.The values of gip for tri-n-alkylamines, calculated from eqn (9), are transformed to the values of gipo using18 3 ( 2 ~ + E,) = (E,+2)(2E+1)pc. KLOFUTAR, S. PALJK AND R. MALNERSIC 3699 and are summarized in table 3. From table 3 it can be seen that these values show a tendency to increase with increasing number of carbon atoms in the molecule. The comparison of Onsager, po, and Kirkwood, gipo, values indicates that the values of g are higher than unity and increase with increasing chain length. Following transition-state theory, the dependence of absolute viscosity, q, at constant pressure on temperature is given by23 (1 1) where q* is a constant and E i is the energy of activation for viscous flow at constant pressure. From fig.5 it can be seen that for tri-n-alkylamines In q is a linear function of 1 / T in the temperature range studied. The values of In 'I* and E;, together with the standard error CJf the estimate, s, obtained by the method of least squares, are given in table 4. The values of absolute viscosity determined previously for TEA at 293.1 5,1° 298.15,'~ 13-24 303.157 and 313.15 KIO and for TBA at 293.15 and 313.15 Kl0 agree very well with our values. In fig. 6 the dependence of energy of activation for viscous flow at constant pressure as a function of the number of carbon atoms in the tri-n-alkylamine molecule is given. In v = In 'I*+-- E2, RT - 4.000 - 5 000 h - I m - I E 5 -6000 OD c C v M - 7 000 -8.000 I I I I I 3.0 3.1 3.2 3.3 3 4 103 KIT FIG.5.-Viscosities of TEA (O), TBA (m), THxA (A), TOA (A), TDA (0) and TLA (0) as a function of 1jT. From the data on the energy of activation for viscous flow at constant pressure for TEA and its energy of vaporization, AEv, calculated from the enthalpy of vaporization at 293.15 K,7 a ratio of AEV/Ei z 4 is obtained. This value is typical of non-spherical molecules and is in accordance with the postulated pyramidal structure of TEA.I* On the basis of free-volume theory, the isobaric viscosity related to the density is (12) presented as25 v-1 = b(d-1 -&I) where b and d , are constants.3 700 PRO PER TIES OF LIQUID TR I-n-A L K Y LA MINES TABLE 4.-vALUES OF In v* AND ENERGY OF ACTIVATION FOR VISCOUS FLOW AT CONSTANT PRESSURE, TOGETHER WITH THE STANDARD ERROR, FOR TRI-n-ALKYLAMINES IN THE TEMPERATURE RANGE STUDIED compounds Ei/J mol-l -In f ‘ s x lo2 TEA 8 148+58 1 1.25 & 0.02 0.3 TBA 12 945+ 108 1 1.89 f 0.04 0.5 THxA 18 640 k 225 13.09 k 0.09 1 .1 TOA 20 843 & 125 13.26 & 0.05 0.6 TDA 23 379f 125 13.62 k 0.07 0.7 TLA 25 467k 190 14.02 f 0.07 0.9 21000 18000 - I - E c, \ ++a 4 12000 6000 I I I I I I I I 6 16 2 1 32 FIG. 6.-Energies of activation for viscous flow at constant pressure of TEA (a), TBA (m), THxA (A), TOA (A), TDA (0) and TLA (e) as a function of the number of carbon atoms in the molecule, n. n It was shown that the energy of activation at low pressure can be given by26$27 E$ = E$+nAVt (13) where EL is the energy of activation at constant volume, A f l is the activation volume, RT[a In 41/aP],, and n is the internal pressure of the liquid, defined as n = aT/B, where The dependence of free volume on temperature can be given by the van’t Hoff relation in the form2* (14) is the compressibility coefficient.Ed In ( d - l - d ~ l ) = a-- RTc. KLOFUTAR, S. PALJK AND R. MALNERSIC 370 1 where a is a constant and Ed is the expansion energy, i.e. the work done against the internal pressure to create the holes in the liquid. At low pressure Ed is equal to nA V l . A volume-restrained liquid is characterized by nAT/f % and thus E& gnAV1 = E d . For such a liquid eqn (1 1 ) and (12) equally well describe the viscosity data. On the other hand, for an energy-restrained liquid nA V i is only a small fraction of Therefore, the behayiour of such a liquid depends on the energy of activation at constant volume, Pv.As a consequence, for such a liquid both viscosity equations cannot describe the experimental data. The values of the expansion energy for the liquids investigated are calculated from their density data via eqn (14) in the following fashion. A value of d, is selected and the corresponding a and E(d) parameters calculated by a linear least-squares fit of In (d-l - d ~ l ) against 1 / T. Then the d* value is changed and the procedure repeated until a minimum in the standard error is obtained. In table 5 the best-fit parameters are listed for the tri-n-alkylamines studied. In addition, it is interesting to note that the dependence of density on temperature is equally well described by eqn (3) and (1 4).TABLE 5.-vALUES OF PARAMETERS OF EQN (14), TOGETHER WITH THE STANDARD ERROR, FOR TRI-n-ALKYLAMINES IN THE TEMPERATURE RANGE STUDIED compound 4 a E,/J mol-l s x lo4 TEA 0.8500 1.521 k0.003 7680 3.6 TBA 0.8700 0.932 & 0.003 7120 3.5 THxA 0.8800 0.768 5 0.004 71 I0 4.4 TOA 0.9000 0.564+0.002 6500 3.0 TDA 0.9000 0.626 0.003 6910 3.6 TLA 0.9100 0.417 k0.003 6300 3.7 From table 5 it is evident that with increasing number of carbon atoms in the tri-n-alkylamine molecule, the values of d* increase slightly while a and E d decrease. However, the values of Ed differ by ca. 10% from TBA to TLA. From this it follows that the work required to create the holes does no change drastically with increasing number of carbon atoms in the tri-n-alkylamine molecules.On the other hand, the values of the ratio of the energy of activation for viscous flow at constant pressure to the expansion energy, E&/E,, increase from 1.06 for TEA to 4.04 for TLA. From the energies required to create the holes, it may be suggested that with increasing chain lengths in tri-n-alkylamines the units of flow are no longer whole molecules but are smaller; hence, the greater molecules move in sections rather than in single units. The non-linear relationship between E$ and n (see fig. 6) confirms the adopted mechanism of viscous flow. From the fact that the viscosity data at the temperatures studied obey both viscosity eqn (1 1) and (12) for TEA only, and from the values of Ea/E, it may be concluded that of the tri-n-alkylamines investigated only TEA belongs to the class of volume-restrained liquids, while the others belong to the group of energy- restrained liquids.On the basis of the dielectric properties of tri-n-alkylamines, it is obvious that they are weakly polar liquids. Their bulk properties, in general, resemble those of liquid n-alkanes. On the other hand, the Kirkwood correlation coefficient g shows that tri-n-alkylamines are non-associated liquids. In addition, from the values of coefficient g it can be assumed that with increasing number of carbon atoms the tendency to parallel orientation of neighbouring molecules increases.3702 PROPERTIES OF LIQUID TRI-n-ALKYLAMINES We thank Mrs J. Burger for her skilful technical assistance. We also thank the Slovene Research Community for financial support.A. Rieure, M. Pumeau and B. Tremillon, Bull. SOC. Chim. Fr., 1964, 1053. C . Klofutar, s. Paijk and D. Kremser, J. Inorg. Nucl. Chem., 1975, 37, 1729. F. Oehme and H. Wirth, Die Bestimmung des molekularen Dipolmomentes (Wissenschaftlich- Technischen Werkstatten GMBH, Weilheim Oberbayern, 4/70). R. Finsy and R. van Loon, J. Phys. Chem., 1976, 80, 2783. M. R. Cannon, R. E. Manning and J. D. Bell, Anal. Chem., 1963, 32, 355. L. Korson, W. Drost-Hansen and F. J. Millero, J . Phys. Chem., 1969, 73, 34. J. A. Riddick and W. B. Bunger, Organic Solvents, in Techniques of Chemistry, ed. A. Weissberger Wiley, New York, 1970), vol. 11, pp. 66, 71-107 and 439. Handbook of Chemistry and Physics, ed. C. D. Hodgman, R. C. Weast and S. M. Selby (Chemical Rubber Company, Cleveland, 1957-1 958), p.2048. Y. Marcus, Introduction to Liquid State Chemistry (Wiley, London, 1977), p. 1 1. Chem., 1981, 85, 2520. lo F. Kohler, H. Atrops, H. Kalali, E. Liebermann, E. Wilhelm, F. Ratkovics and T. Salamon, J. Phys. l1 T. M. Letcher and J. W. Bayles, J. Chem. Eng. Data, 1971, 16, 266. l e T. M. Letcher, J. Chem. Thermodyn., 1972, 4, 159. l3 A. Kumar, 0. Prakash and S . Prakash, J. Chem. Eng. Data, 1981, 26, 64. l4 A. S. Kertes and F. Grauber, J. Phys. Chem., 1973, 77, 3107. l5 F. Grauber and A. S. Kertes, J. Chem. Eng. Data, 1973, 18, 405. l6 J. Liszi, I. Domonkos and I. Szabo, Acta Chim. Acad. Sci. Hung., 1973, 79, 323. l7 D. Gourisse, French Report, CEA-R-3005 (1966), in Solvent Extraction Chemistry, ed. D. Dyrssen, l8 S . P. Smyth, Dielectric Behaviour and Structure (McGraw-Hill, New York, 1955), pp. 24, 30-34 and I9 Y. Marcus and A. S. Kertes, Ion Exchange and Solvent Extraction of h e t a l Complexes (Wiley, 2o K. Higasi, Sci. Pap. Inst. Phys. Chem. Res. (Jpn), 1937, 31, 311. 21 W. J. McDowell and K. A. Allen, J. Phys. Chem., 1959, 63, 747. 22 B. Krishna and A. N. Snvastava, Aust. J . Chem., 1966, 19, 1847. 23 S. Glasstone, K. J. Laidler and H. Eyring, The Theory of Rate Processes (McGraw-Hill, New York, 24 N. F. Lange, Handbook of Chemistry (Handbook Publishers, Sandusky, 1956), p. 1657. 25 J . H . Hildebrand, Viscosity and Dzflusivity (Wiley, New York, 1977), p. 1. 26 B. Cleaver, Discuss. Faraday SOC., 1967, 43, 241. 27 F. B. Macedo and T . A. Litovitz, J . Chem. Phys., 1965, 42, 245. 28 R. S . Stairs, Can. J. Chem., 1981, 59, 2170. J. 0. Liljenzin and J. Rydberg (North-Holland, Amsterdam, 1967), p. 373. 309. London, 1969), p. 740. 1941), p. 477. (PAPER 2/880)

ISSN:0300-9599    

DOI:10.1039/F19827803693

出版商:RSC    

年代:1982

数据来源: RSC