摘要:

J. Chem. Soc., Faraday Trans. I, 1984,80, 1517-1524 Binary Systems of 1, I ,2,2-Tetrachloroethane with Benzene, Toluene, p-Xylene, Acetone and Cyclohexane Part 2.-Dielectric Properties at 308.15 K BY JAGAN NATH* AND A. D. TRIPATHI Chemistry Department, Gorakhpur University, Gorakhpur 273001, India Received 12th September, 1983 Measurements of dielectric constants, E, have been made for binary liquid mixtures of 1,1,2,2-tetrachloroethane (TCE) with benzene, toluene, p-xylene, cyclohexane and acetone at 308.15 K. Measurements of refractive indices, n, have also been made for binary liquid mixtures of CHCl,CHCl, with acetone at 308.15 K. The values of the quantity A&, which refers to the deviations of the dielectric constants of these mixtures from values for the ideal case, have been calculated.The positive values of A& obtained for TCE+acetone mixtures are attributed to the formation of a molecular complex between acetone and CHCl,CHCl,, whereas the negative values of A& for the other systems may be explained as being due to a decrease in the degree of alignment of the molecular dipoles with changing composition of the mixture. The values of the equilibrium constant, K,, for the formation of a 1 : 1 intermolecular complex between acetone and TCE have been calculated, using the dielectric-constant data. In addition, values of the apparent dipole moments, papp, of TCE at various mole fractions in the solvents cyclohexane, benzene and p-xylene have also been calculated. The values of pipp show that the molecules of TCE are self-associated, and this self-association is more favoured in cyclohexane solution than in solutions of benzene andp-xylene; this is attributed to the existence of specific interactions between TCE and the aromatic hydrocarbons.The values of pipp also show that the strength of specific interactions of TCE with p-xylene is greater than in the case of benzene. Binary systems of 1,l72,2-tetrachloroethane (TCE) with aromatic hydrocarbons, acetone and cyclohexane are of interest because there exists a donor-acceptor interaction between the components. This is caused by the presence of four C1 atoms and two H atoms in TCE, which can thus act as a a-acceptor toward, and be involved in the formation of hydrogen bonds with, aromatics and acetone. The latter act as n-donors and an n-donor, respectively.The system cyclohexane + TCE, in which only dispersion, dipolar and induction forces are believed to be present between the components, can be used as a reference system. Extensive studies of the interactions between components in such systems have not been made. However, recently' measurements of ultrasonic velocities, adiabatic compressibilities and excess volumes in the case of binary liquid mixtures of TCE with benzene, toluene, p-xylene, acetone and c-C,H,, were carried out at 298.15 and 308.15 K; from these studies there appears to exist a specific interaction between TCE and the other component. Dielectric- constant measurements of binary mixt~res~-~ have also shed light on the existence of a specific interaction between the components. Furthermore, the values of the apparent dipole moments, pap?, of the various components in ~ o l u t i o n ~ - ~ have also provided evidence concerning interactions between the components.The values of paPp can be obtained from measurements of the dielectric constants of binary liquid mixtures. Hence, in order to obtain conclusive evidence of the existance of a specific 15171518 DIELECTRIC PROPERTIES OF BINARY SYSTEMS interaction between TCE and aromatics or acetone we have made measurements of the dielectric constants for binary liquid mixtures of TCE with benzene, toluene, p-xylene, acetone and c-C,H12 at 308.15 K. Since the refractive indices of acetone, TCE and binary mixtures of TCE with acetone were needed to calculate both papp values of TCE in various solvents and the molar polarisations for mixtures of TCE with acetone, measurements of refractive indices have also been made at 308.15 K.The results of the present measurements are interpreted in this paper. EXPERIMENTAL MATERIALS Benzene, toluene,p-xylene, acetone, cyclohexane and 1,1,2,2-tetrachloroethane were purified, and their purity was checked as described earlier.' METHODS The dielectric-constant ( E ) measurements were made at 308.15 K and at a frequency of 1.8 MHz with a dekameter (type DK,,, Wissenschaftlich-Technische, Werkstatten, Germany), using one cell (MFL 1/S, no. 2078) for mixtures of TCE with benzene, toluene, p-xylene and cyclohexane, and another (MFL 2/S, no. 2084) for mixtures of TCE with acetone. The cells were thermostatted through the outer jacket using a water bath whose temperature was controlled to within kO.01 K.The two cells were first calibrated using liquids of known dielectric constants,a and dielectric-constant measurements were then made for the pure liquids and binary mixtures studied in the present programme. The precision of the dielectric-constant measurements was better than 0.0004 unit for dilute solutions of TCE in benzene, p-xylene and cyclohexane, and ca. 0.001 unit for mixtures of TCE with toluene and acetone and for mixtures having higher concentrations of TCE in benzene, p-xylene and cyclohexane. Refractive-index (n) measurements accurate to within f 0.0002 were carried out using a thermostatted Abbe refractometer at 308.15 K. The values of n were obtained for sodium-D light.RESULTS AND DISCUSSION The experimental values of e for the pure liquids benzene, toluene, p-xylene, acetone, cyclohexane and TCE, and for binary mixtures of TCE with benzene, toluene, p-xylene, acetone and cyclohexane at 308.15 K are given in table 1, where xA refers to the mole fraction of TCE. The values (see table 1) of E obtained for benzene, toluene, p-xylene, cyclohexane and acetone are 2.2540, 2.3552, 2.2454, 1.9992 and 19.749, respectively, in excellent agreement with the literature values8 of 2.2540, 2.3547, 2.2460, 1.9990 and 19.745, respectively. The value of E for TCE was found to be 7.096, which is in good agreement with the value of 7.090 which is based on data available in the literat~re.~ The present experimental values of the refractive indices, n, for acetone and TCE are 1.3520 and 1.4865, respectively, which can be compared with literature valueslO of 1.351 57 and 1.48658, respectively.The values of the refractive indices, n12, for TCE+acetone mixtures have been fitted by the method of least squares to the equation (1) where xA refers to the mole fraction of TCE. Values of Ae, the deviation in the dielectric constants of the mixtures from values obtained from the ideal volume-fraction-mixture law, have been calculated from the relation where el, is the dielectric constant of the mixture and el and E , are the dielectric constants of the pure liquids 1 and 2, for which volume fractions in the mixture are n12 = 1.35 1 60 + 0.195 26x, - 0.060 62xi (2) A& = & 1 2 - 4 1 & 1 - 4 2 & 2J.NATH AND A. D. TRIPATHI 1519 Table 1. Dielectric constants for TCE in various mixtures at 308.15 K TCE + benzene TCE + toluene TCE +p-xylene & & X A X A X A & 0.0000 0.0028 0.0035 0.0052 0.0071 0.0 164 0.0350 0.0454 0.05 19 0.0662 0.0759 0.0886 0.2087 0.3252 0.5357 0.6 150 0.6559 0.6593 0.7016 0.7599 0.97 17 1 .oooo 2.2540 2.2636 2.2658 2.2718 2.2786 2.3106 2.378 2.410 2.436 2.493 2.525 2.563 2.994 3.436 4.332 4.696 4.893 4.897 5.109 5.432 6.89 1 7.096 0.0000 0.0982 0.2067 0.2459 0.3894 0.4528 0.5239 0.5696 0.6434 0.657 1 0.6636 0.7093 0.8106 0.9524 1.0000 2.3552 2.638 2.959 3.109 3.593 3.821 4.152 4.352 4.667 4.770 4.790 4.980 5.618 6.628 7.096 0.0000 0.00 13 0.0030 0.0060 0.0074 0.0084 0.0275 0.0437 0.0542 0.0626 0.0730 0.0808 0.1967 0.2074 0.3393 0.4539 0.5880 0.6002 0.8807 0.9592 1 .oooo 2.2454 2.2487 2.2528 2.2598 2.2642 2.2659 2.314 2.359 2.386 2.416 2.447 2.467 2.780 2.8 16 3.203 3.663 4.190 4.23 1 5.966 6.638 7.096 TCE + cyclohexane TCE +acetone X A X A & & 0.0000 0.0005 0.0009 0.0025 0.005 1 0.0082 0.0096 0.0372 0.0460 0.0470 0.0760 0.0862 0.1 123 0.1650 0.2225 0.497 1 0.5178 0.6952 0.7053 0.921 1 1 .oooo 1.9992 2.0008 2.0020 2.0067 2.0 150 2.0249 2.0285 2.1 12 2.138 2.139 2.209 2.239 2.306 2.491 2.708 3.804 3.910 4.886 4.938 6.476 7.096 0.0000 0.0778 0.1016 0.2004 0.3019 0.4049 0.4904 0.5532 0.5996 0.67 14 0.7002 0.7795 0.8793 1 .oooo 19.749 18.817 18.601 18.102 17.161 16.540 15.474 14.42 1 13.741 12.634 12.21 1 1 1.032 9.340 7.0961520 DIELECTRIC PROPERTIES OF BINARY SYSTEMS m m *.+ 2.0 ' 8 8 w +'*O* 0.0 0.0 - 0.4 1 % a "$ A I d -0.8t I ! I , 0.0 0.2 0.4 0.6 0.8 I 0 Fig. 1. Plot of A& against the mole fraction of TCE, xA, at 308.15 K: a, TCE+benzene; V, TCE + toluene; A, TCE +p-xylene; 0, TCE + cyclohexane; ., TCE + acetone. and 42, respectively. It has been pointed out that dielectric constants of polar mixtures can be represented as linear functions of the volume fraction." Fig. 1 , however, shows that the values of A& are negative for the systems TCE+benzene, TCE + toluene, TCE +p-xylene and TCE + cyclohexane, and positive for TCE + acetone. Comparison on a volume-fraction basis largely compensates for the ' dipole- dilution' effect, as it has been termed by Franks and Ives.12 The negative values of A&, which are of about the same order of magnitude (see fig.1) for the systems TCE + benzene, TCE + toluene, TCE +p-xylene and TCE + cyclohexane, can be attributed to a decrease in the degree of alignment of the molecular dipoles with changing composition of the mixture. A& is found to be positive for systems where strong specific interaction is believed to be present between the components.2 Measurements of excess volumes'. l3 indicate the existence of a specific interaction leading to the formation of complexes between TCE and the aromatic hydrocarbons. The negative values of A& for the systems involving TCE and the aromatic hydrocarbons may be explained by the predominance of contributions to A& from dipolar forces over those from specific interactions. The highly positive values (see fig.I ) of A& for the system TCE + acetone show that acetone forms a strong molecular complex with TCE in the liquid state; this is also true in the case of the pure binary system pyridine + chloroform, in which case a 1 : 1 complex is believed to be formed2 because of hydrogen bonding between chloroform and pyridine. In order to obtain further evidence to support the fact that acetone forms complexes with TCE we have calculated values of the total molar polarisation, P, for acetone, TCE and binary mixtures of TCE with acetone at 308.15 K, using the Kirkwood-Frohlich equation :14 (3) ( E - n2) (2e + n2) V 9& P =J. NATH AND A. D. TRIPATHI 1521 620 540 - I - E 460- E e 1 420- 380 - 340 - I I I I I 1 0.0 0.2 0.4 0.6 0.8 .1.0 *D Fig. 2. Plot of the apparent molar polarisation, PD, against mole fraction, xD, of acetone for the system TCE + acetone at 308.15 K. where Vdenotes the molar volume. The values of n used to calculate P for the mixtures were obtained from eqn (1). The molar volumes of acetone and TCE were obtained from the available density data,lo whereas those for the mixtures were estimated from the molar volumes of the pure liquids and measurements of the excess volumes.1 The total molar polarisations of the mixtures were then used to calculate values of the apparent molar polarisation, PD, of acetone at different compositions of its mixtures with TCE, in a manner similar to that described by Rastogi and Nath.15 The values of PD so obtained are plotted as a function of mole fraction of acetone, xD, in fig.2, which shows that the value of P,, increases sharply with a decrease in the mole fraction of acetone. This confirms that there exists a strong specific interaction between acetone and TCE, leading to the formation of molecular complexes between the two species in the liquid state. Similarly, values of the apparent molar polarisation, PA, of TCE at different compositions of its mixtures with acetone were calculated. Considering a 1 : 1 complex DA to be formed between acetone (D) and TCE (A), the value of the total molar polarisation, PDA, of the complex formed was obtained in a manner similar to that described by Earp and Glasstone.ls The value of PDA was found to be 733.4 cm3 mol-l. Furthermore, the values of the equilibrium constant, K,, for the formation of the complex DA, were also calculated as described by Earp and Glasstone.16 The results show a significant variation in Kf with the composition of the mixture.Rivail and Thiebaut2 have also observed that in the case of the system pyridine +chloroform, the values of Kf estimated from dielectric-constant data exhibit a significant variation with the composition of the liquid. As pointed out by Rivail and Thiebaut,2 a theory1' based upon electrostatic interactions of the solute with the liquid predicts a linear variation of the logarithm of Kf with the quantity1522 0.4 C 0 .I Y z 0.0- i? e, - v 1 -0.4 - M - -0.8 DIELECTRIC PROPERTIES OF BINARY SYSTEMS - - - I I I I I 0.12 0.13 0.14 0.15 0.16 f (€1 Fig. 3. Plot of log [&/(mole fraction)-’] againstflc) for the system TCE + acetone at 308.15 K.where E, refers to the infinite-frequency dielectric constant of the mixture. For calculations offTE) we have taken E, = n2. In fig. 3 are plotted values of log Kf as a function off(&). There is a linear variation of log Kf withf(E) for TCE+acetone, thus suggesting that the values of Kf calculated using the simple approach of Earp and Glasstonels are in accord with the theory2. l7 based upon the electrostatic interactions of the solute with the liquid. McClellan and Nicksicl* have shown that the molecules of TCE are self-associated through hydrogen bonding, and Campbell et aL6 and Stokes and Marsh5 have shown that the values of the apparent dipole moment, papp, of polar solutes in non-polar solvents furnish useful information about both self-association of the solute molecules and association of the solute molecules with the molecules of the solvent.Hence, to determine whether the present study gives evidence concerning the self-association of TCE molecules, and to find out if a specific interaction exists between TCE and the aromatic hydrocarbons, we have calculated the values of the apparent dipole moment, papp, of TCE at various mole fractions in the non-polar solvents benzene, p-xylene and cyclohexane using the equation5* l9 where xA is the mole fraction of the polar solute, E is the dielectric constant of the solution, E~ is the dielectric constant of the non-polar solvent (benzene, p-xylene or cyclohexane) and E; is the internal dielectric constant of the solute.Vm, 4 and & are the molar volumes of the solution, the non-polar solvent and the polar solute, respectively. N is Avogadro’s constant and k is Boltzmann’s constant, p s , o is the moment of the isolated polar molecule and g is the Kirkwood correlation parameter.2o As mentioned by Stokes and M a r ~ h , ~ we have taken (gp$-,)i to be equal to the apparent dipole moment, paPp, of the solute. The values of 4, V, and Vm used to calculate (gp:, ,); from eqn (5) were ascertained as described earlier in this paper. The value of E;, which was obtained from the refractive index of TCE as described by Stokes and M a r ~ h , ~ was 2.387. Fig. 4 shows the concentration dependence of pipp for TCE in (i) cyclohexane, (ii) p-xylene and (iii) benzene, with a logarithmic scale along the abscissa.The dielectric behaviour of TCE in benzene is similar to that in p-xylene, but different from that in cyclohexane. The concentration dependence of pipp for TCE in benzene and p-xyleneJ. NATH AND A. D. TRIPATHI 1523 I I I I I I - 3.0 -2 .o -1 .o 0 .o I .o log (C/mol dm-3) Fig. 4. Plot of pgPp against the logarithm of solute concentration for TCE in (i) cyclohexane, (ii) p-xylene and (iii) benzene. solutions is similar to that of octanols in benzene s~lution,~ whereas the concentration dependence of pipp for TCE in cyclohexane solution is similar to that of octanols in cyclohexane solution.' The concentration dependence of &pp in fig. 4 gives information about the association behaviour of TCE.The initial rise in p:pp is indicative of the formation of the first, high-dipole-moment species. Fig. 4 therefore shows that the formation of this species takes place at lower concentrations in cyclohexane solution and that it is delayed in benzene and p-xylene solutions. This shows that the TCE monomer is stabilised little by interaction with the non-polarisable solvent (cyclo- hexane), but is stabilised more by association with the solvents benzene and p-xylene. The self-association of TCE is therefore inhibited by solute-solvent interactions in the solvents benzene and p-xylene. Woolley and Hepler,21 using thermodynamic data, have found similarly that phenol is more self-associated in cyclohexane than in benzene. The association of TCE with benzene and p-xylene can be attributed to the existence of a specific interaction between TCE and the aromatic hydrocarbons, which may be due to the formation of a weak hydrogen bond through the interaction of the hydrogen of TCE with the n-electrons of the aromatic ring.However, there is also a possibility that TCE is involved in the formation of a charge-transfer complex with the aromatic hydrocarbon through the interaction of the chlorine atoms with the aromatic n-electrons. Fig. 4 also shows that a maximum in pipp occurs at low concentrations of TCE in benzene and that this maximum is more pronounced in p-xylene, whereas in cyclohexane such a maximum does not appear. This observation, that the maximum inp:,, at low concentrations of TCE is more pronounced inp-xylene than in benzene, shows that the strength of the specific interaction of TCE withp-xylene is greater than that with benzene.This can be attributed to the fact that the n-electron density of the aromatic ring is increased in p-xylene due to the presence of two CH, groups. CONCLUSIONS In conclusion, we note that the dielectric-constant data show that molecules of TCE are self-associated, and that this self-association is more favoured in cyclohexane than in benzene and p-xylene, thus indicating that the TCE monomer is stabilized little by1524 DIELECTRIC PROPERTIES OF BINARY SYSTEMS interaction with cyclohexane in comparison with benzene and p-xylene. This shows that the self-association of TCE is inhibited by solute-solvent interactions in the solvents benzene and p-xylene, a fact which confirms the existence of a specific interaction between TCE and the aromatic hydrocarbons.The plots of the values of ptpp in fig. 4 show that the specific interaction of TCE is stronger with p-xylene than with benzene, a fact which has been attributed to the increased n-electron density of the aromatic ring of p-xylene. The dielectric-constant data also show that acetone forms strong complexes with TCE in the liquid state. The presence of a specific interaction between TCE and the aromatic hydrocarbons can be explained as due to the formation of a weak hydrogen bond between the hydrogen atoms of TCE and the n-electrons of the aromatic ring. There is, however, also a possibility that TCE may form a charge-transfer complex with the aromatic hydrocarbons, via chlorine- atom-n-electron interactions.On the other hand, the complexation between acetone and TCE can be attributed to the formation of strong hydrogen bonds between the hydrogen atom of TCE and the lone-pair electrons on the oxygen atom of acetone. We are grateful to Prof. R. P. Rastogi, Head of the Chemistry Department, Gorakhpur University, Gorakhpur, for his encouragement during the course of these investigations. Thanks are also due to the C.S.I.R. and U.G.C, New Delhi, for financial support. J. Nath and A. D. Tripathi, J. Chem. Eng. Data, 1983, 28, 263. J. L. Rivail and J. M. Thiebaut, J. Chem. Soc., Faraday Trans. 2, 1974,70, 430. J. Nath and S. N. Dubey, J. Phys. Chem., 1980,84, 2166. J. Nath and S . S . Das, Indian J. Pure Appl. Phys., 1981, 19, 343. R. H. Stokes and K. N. Marsh, J. Chem. Thermodyn., 1976, 8, 709. C. Campbell, G. Brink and L. Glasser, J. Phys. Chem., 1975, 79, 660. C. Campbell, G. Brink and L. Glasser, J. Phys. Chem., 1976, 80, 686. N. A. Lange, Lunge’s Handbook of Chemistry (McGraw-Hill, New York, 1973). International Critical Tables of Numerical Data: Physics, Chemistry and Technology (McGraw-Hill, New York, 1929), vol. VI, p. 84. lo J. Timmermans, Physico-Chemical Constants of Pure Organic Compounds (Elsevier, Amsterdam, 1950). I1 T. B. Hoover, J. Phys. Chem., 1969, 73, 57. l2 F. Franks and D. J. G. Ives, Q. Rev. Chem. Soc., 1966, 20, 1. l3 M. Gracia, S. 0 t h and C. G. Losa, J. Chem. Thermodyn., 1975, 7, 293. l5 R. P. Rastogi and J. Nath, Indian J. Chem., 1967, 5, 249. l6 D. P. Earp and S . Glasstone, J. Chem. Soc., 1935, 1709. l8 A. L. McClellan and S. W. Nicksic, J. Phys. Chem., 1965, 69, 446. 2o J. G. Kirkwood, J. Chem. Phys., 1939,7,911. C. Moreau and G. Douhkret, J. Chem. Thermodyn., 1976, 8, 403. J. Barrio1 and A. Weisbecker, C.R. Acad. Sci., Ser. C, 1967, 265, 1372. H. Frohlich, Trans. Faraday Soc., 1948, 44, 238. E. M. Woolley and L. G. Hepler, J. Phys. Chem., 1972, 76, 3008. (PAPER 31 1603)

ISSN:0300-9599    

DOI:10.1039/F19848001517

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1984, 80, 1525-1537 Adsorption of Ions at the Cellulose/Aqueous Electrolyte Interface Part 1 .--Charge/pH Isotherms BY THELMA M. HERRINGTON* AND BRIAN R. MIDMORE Department of Chemistry, University of Reading, Reading RG6 2AD Received 14th September, 1983 Charge/pH isotherms are reported for various cellulose fibres : cotton linters, bleached sulphate pulp and unbleached sulphate pulp. This charge was determined as a function of pH for 1.0, 0.1, 0.01 and 0.001 mol dm-3 NaC1. The effect of various cations on the charge on bleached sulphate pulp was also investigated. The results for cotton linters and bleached sulphate pulp are analysed in terms of polyelectrolyte theory. The coagulation of cellulose fibres at the mesh of a paper machine is a complex process with many variables.Many workersly have determined the zeta potential of cellulose fines as a function of pH and added electrolyte. This study was undertaken to find out the charge at the cellulose surface rather than at the plane of shear. Previous workers investigating the surface charge of cellulose have treated the system as a polyacid and have investigated the ion-exchange reaction between acidic protons and metal cations cellulose-H + M+ + cellulose-M+ + H+. This reaction in aqueous solution is equivalent to the adsorption of hydroxide ions, as the exclusion of water from the cellulose molecule is not analytically detectable: OH- / cellulose-H +OH- --+ cellulose, H The acidity of cellulose has generally been attributed within the cellulose structure. -H,O --+ cellulose-.to occasional carboxy groups Neale3 titrated oxycellulose against alkali ; only the ‘end-point ’ of the titration was determined, which gave the carboxy content. He then dispersed the oxycellulose in sodium chloride solution and measured the pH. This gave him a value for the degree of dissociation, a, and he estimated a pK of 4.2 for the acid group in oxycellulose. Neale in effect determined a single point on a pH/alkali adsorption isotherm at ca. pH 7. The work of Heymann4 which determined the quantity of acid released after dispersing cotton linters in neutral electrolyte solution also gave a single point on this isotherm. The first to find more points on this isotherm was David~on.~ He worked on oxycellulose prepared by the action of alkaline hypobromite on cotton using the adsorption of methylene blue, which has a high affinity for the carboxy group.However, the ionic strength was not held constant. None of these workers attempted 15251526 THE CELLULOSE/ AQUEOUS ELECTROLYTE INTERFACE to determine the surface area. The first work attempted on cellulose using the titration technique with comparison to a blank to determine the absorption of hydroxide ions was performed by Edelson.6 He titrated microcrystalline cellulose both in water and in 0.1 mol dm-3 KCI, varying the pH between 2 and 11. Only in the latter case was the ionic strength kept effectively constant. Two acid groups were detected, having pKvalues of 4.0 and 9.2, respectively. The latter was concluded to be ammonia present on the surface as an impurity, while the former was confirmed to be a c6 carboxy group.This was concluded as oxycellulose prepared by reaction with nitrogen dioxide gave a pK of 4.1. Using geometrical approximations, a charge density of 0.0 16 C m-2 can be estimated for the crystallites at the neutralization point. From the work of Davidson5 and Edelson6 a pH value of CQ. 2 for the P.Z.C. of cellulose seems reasonable. As far as we can ascertain this is the first work to determine the charge/pH isotherms on various types of cellulose fibres and also to determine the surface area of these fibres by different techniques. EXPERIMENTAL MATERIALS All salts were AnalaR grade and were used without further purification. The standard sodium hydroxide solution prepared from stock? was checked to be free of carbonate by titration; a concentration of only 0.1 % of carbonate gave two end-points.All water used was double distilled and had a conductivity of < 5 x The cotton linters were Hercules Chemical Cotton (Holden Vale Manufacturing Co., Haslingden, Lancashire), which is prepared from raw cotton linters by wet and dry mechanical cleaning, pressure digestion in caustic soda, bleaching, washing and drying. The degree of polymerization was 1040 and the carboxy content of the sample (determined by Hercules) was 1.5 x lop2 mol kg-l. The bleached sulphate pulp (STORA32) and unbleached sulphate pulp were Scots Pine supplied by Pira (Leatherhead, Surrey) and had never been dried. Both pulps were used in their wet form (16%, w/w, solids) and were stored at 4-6 "C.The cotton linters came in dry sheet form; a blender was used to disperse the fibres in water allowing a standard time of 90 s to produce the slurry, which was filtered on a sintered glass funnel to give a wet pulp. The fibres were treated with acid to remove unwanted cations and to produce the acid form of acidic groups present: 20 g of the fibres were soaked in 2 dm3 of 0.1 mol dmp3 HC1 for 20 h, washed and then soaked in conductivity water for a further 20 h to remove the last traces of acid; the supernatant was checked for the absence of chloride. After further washing the fibres were sucked as dry as possible. 0-l cm-l. DETERMINATION OF CHARGE/PH ISOTHERMS The basic technique consists of the potentiometric titration of a fibre suspension in aqueous electrolyte solution of varying ionic strength with H+ or OH- ions using a glass plus a silver, silver chloride electrode for pH determination.A sample of electrolyte solution of the same volume and concentration as the fibre suspension is then titrated with the same acid or alkali. The difference between the amounts of H+ or OH- ions that produce a given pH in the fibre suspension and the corresponding pH in the blank sample of the electrolyte gives the amount of H+ or OH- ions adsorbed by the surface. The surface charge is defined as where rH+ and r o H - are the surface excess of H+ and OH- ions. Silver, silver chloride electrodes were prepared ele~trolytically.~ The bias potentials of such electrodes did not exceed 0.2 mV. The potentials were measured at 25k0.05 "C using a Metrohm E636 titroprocessor.The electrodes were standardized using 6 standard buffers.' The assumption that ycl- = YNsCl was made; this is reasonable as for most pH values (4-10), mHCl + mNaCl. Y~~~~ was obtained from tables.8 Thus pH is defined as pH = ( E - p ) eIRT+ log (YNaclmC1-)*T. M. HERRINGTON AND B. R. MIDMORE 1527 A polypropylene guard was fitted to prevent tangling of the fires on the electrodes. The titration was carried out in a gas-tight vessel under a current of pure nitrogen to avoid contamination with CO, The greatest problem in extending this titration technique, which has been used for silica gel9 and metal oxides,1° to a suspension of cellulose fibres is the attainment of equilibrium. Because cellulose fibres are highly porous, there is the problem of obtaining uniform distribution of the electrolyte solution within the fibre.Coupled with this is the problem of stirring and ensuring adequate mixing in the bulk solution. To complete the titration in a reasonable period of time, i.e. ca. 4 h, a maximum of 1.8 g of fibre could be added to 450 cm3 of solution. This weight limitation proved to be a problem in the case of the cotton linters because of their very low charge density per gram. More rapid titrations of ca. 2 h in duration led to erroneous peaks and troughs in the differential plot of e.m.f. against volume, indicating non-attainment of equilibrium. There are two dilution effects, one caused by the water in the wet fibre and the other by the addition of acid or alkali.These two effects were corrected for by determining the expected e.m.f. of the solution if the fibres were dry and the acid or alkali of infinite molarity. Allowance was also made for the fact that certain areas of the fibre which are accessible to water are inaccessible to electrolyte. Thus a certain amount of water is effectively removed from the system, increasing the concentration of electrolyte present. As the amount removed i s constant, the absolute increase in electrolyte concentration is proportional to the total electrolyte concentration. Therefore, as we are only interested in variations in hydrogen-ion concentration, this effect only assumes significance below pH 2. The volume of water so absorbed by the fibre was determined in connexion with the negative adsorption experiments (Part 2 of this series).The continuous titration method indirectly determines the hydrogen-ion concentration by measurement of e.m.f. The precision of the titration technique between the pH values of 4 and 10 is governed by errors in the concentration of titrant, volume added and weight of fibre, which amount to 0.2%. The reproducibility of the results between batches of fibre was very high, the error being k 1 %. Even relatively large errors in the e.m.f., e.g. 0.3 pV, at pH 4 generate only small errors of kO.3 pmol. However, below pH 3 the precision rapidly deteriorates and small errors in e.m.f. generate large errors in ArOH-. Below pH 3 the continuous titration technique was not employed and instead the single-point method was used in which the acid concentration in the supernatant was determined by direct titration against alkali.HYSTERESIS EXPERIMENT A hysteresis experiment was performed on the bleached sulphate pulp to determine the reversibility of the system. It was performed in the same way as the normal continuous titration, using 0.1 mol dmP3 NaCl as the electrolyte. After 4 h of titration with alkali to pH 10.5, the titrant was changed to hydrochloric acid and the pH lowered to its initial value over a further period of 4 h. From fig. 1 it can be seen that the degree of hysteresis is small, amounting to 14% of the maximum charge at its greatest at pH 9.5. The amount of hysteresis continued to drop to a value of 8 % of the maximum charge, which suggests that some at least of this hysteresis is not permanent and is time-dependent.In view of the very high pH, 10.5, to which the suspension was raised, this degree of hysteresis seems very small, SURFACE AREA Extensive determinations of the surface area of cellulose fibres were made both by B.E.T. and negative-adsorption techniques. The effect of beating and of oxidation on the surface area were investigated. These measurements are discussed in detail in Part 2 of this series. In this paper, the surface charge for the adsorption isotherms is recorded in C g-' of fibre, as this is the quantity which is directly determined. SODIUM-ION ADSORPTION MEASUREMENTS The adsorption of sodium ion was determined at a few points on the isotherm to check the charge balance. The sodium adsorbed together with the negative adsorption of chloride (this is in fact negligible: see Part 2 of this series) should exactly balance the surface charge.The adsorption of sodium was determined using the principle of the single-point technique but1528 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE 1c 9 8 PH 7 6 5 4 I I I amount of OH- added/pmol amount of H" added/pmol I I I I I I i 200 I I 1 I I00 200 0 Fig. 1. Hysteresis curve for bleached sulphate pulp in 0.1 mol dm-3 NaCl: ., forwards; V, backwards. Table 1. Sodium adsorption results for bleached sulphate pulp PH st7+/c g-1 - So& g-' 0.001 rnol dm-3 NaCl 10.89 3.50 9.39 2.84 4.91 1.41 3.10 2.90 1.35 0.01 mol dm-3 NaCl 10.82 2.92 3.10 analysing for sodium using a flame photometer. The experiment was performed in and mol dm-3 NaC1.In table 1 it can be seen that the sodium adsorption agrees well with the hydroxide adsorption. This indicates that it is the hydroxide ions that are responsible for the charge on the cellulose surface, which is balanced by sodium-ion adsorption in the liquid side of the electrical double layer. RESULTS Charge/pH isotherms at 25 "C were obtained for various cellulose fibres using the continuous titration and single-point techniques. The fibres chosen were cotton linters, a bleached sulphate commercial pulp and an unbleached sulphate commercial pulp. The isotherms were constructed over a pH range 1.3-10.3 and sodium chloride was used as the indifferent electrolyte. Four concentrations of indifferent electrolyte were employed, namely 1.0, 0.1, 0.01 and 0.001 mol dm-3.The isotherms are given in fig.T. M. HERRINGTON AND B. R. MIDMORE 1529 Fig. 2. Charge/pH isotherm at 25 "C for cotton linters. [NaCl]/mol dmP3 as follows: 0, 10-3; 8, 10-2; 0, 10-1; 0, 100. 3 2 - I 00 u . 0 cs: I 1 C Fig !. Charge/pH isotherm at 25 "C for bleached sulphate pulp. [NaCl]/mol dm-3 as foiiows: 0; 10-3; 0, 10-2; a, 10-1; 0, 100.1530 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Fig. 4. Charge/pH isotherm at 25 "C for unbleached sulphate pulp. [NaCl]/mol dmP3 as foiiows: 0, 10-3; 8, 10-2; 0, 10-1; 0, 100. Fig. 5. Effect of various cations on the charge/pH isotherm at 25 "C for bleached sulphate pulp. A, 0.05 mol dm-3 BaC1,; 0, 0.05 mol dm-3 CaC1,; 0, 0.1 mol dm-3 NaCl; A, 0.1 mol dm-3 KC1; ., 0.1 mol dm-3 (C,H,),NCl.T.M. HERRINGTON AND B. R. MIDMORE 1531 2-4. The reader is referred to the original thesis’ for the experimental data from which these plots are drawn. For the bleached sulphate pulp, the effects on the isotherm of various cations were investigated ; solutions of calcium, barium, potassium and tetraethylammonium chloride were used at 0.1 mol dm-3 with respect to the chloride ion. The resulting charge/pH isotherms are shown in fig. 5 . DETERMINATION OF a AND CALCULATION OF pK In the case of a simple weak acid, such as acetic, the ionization constant is given by pK, = pH +log [( 1 - a)/a]. (3) For a polymer carrying a large number of ionizable groups, the apparent ionization constant of an average ionizable group is defined by pK = pH + log [( 1 - a)/al (4) where K will vary with the degree of ionization, since the charged polymer will interact with the hydrogen ions.If the required electrostatic Gibbs energy for the removal of an equivalent of protons at a given degree of ionization is AGel(a), then pK = pK, + AGel (a) log,, e/RT where KO is characteristic of the ionizing group when electrostatic interactions with other ionizing groups are absent. Thus pH = pK, -log [( 1 - a)/a] + AGel (a) log,, e/RT. (6) pK, can be determined from a plot of pK against a by extrapolating a to zero where AGel = 0. If conformational changes occur as the titration proceeds, then there is a non-electric contribution to the Gibbs energy, AGno,.el, and then The presence of AGnon.el is revealed either by the presence of a maximum or minimum in the plot of pK against a or by the construction of a Henderson-Hasselbalch plot.It was found empiricallyll that a plot of pH against log [( 1 - a)/a] yielded a straight line for most polyelectrolytes. Thus (8) pH = pK,, - n log [( 1 - a)/a] where K,, is the average ionization constant and n is a constant depending on the nature and concentration of the polyacid and the ionic strength of the solution. A conformational transition is indicated by a break in the slope.12 The degree of neutralization, a, was determined from the charge/pH isotherms by making the initial assumption that only one acid group was involved in producing the change between the P.Z.C. and the horizontal region of the curve. This assumption is shown to be a good one in the case of cotton linters and bleached sulphate pulp. The horizontal plateau region was taken as the neutralization point and therefore the degree of dissociation at a pH Y was given by (9) a t Y ) = ( Y ) 0 0 /OF where ahY) is the charge at pH Y and 0: is the charge at the point of neutralization. The charge at pH 7 and 0.1 mol dmP3 sodium chloride was taken as an estimate for1532 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Fig.6. Variation of pK with a for cotton linters. [NaCl]/mol dm-3 as follows: 0, 0, 10-2; *, 10-1; 0, loo. 0 7 6 E 5 4 3 Fig. 7. Henderson-Hasselbalch plot for cotton linters. {NaCl]/mol dm-3 as follows 0, 10-3; 0, 10-2; 0, 10-1; 0, 100.T. M. HERRINGTON AND B. R. MIDMORE 1533 Fig. 8. Variation of pK with a for bleached sulphate pulp.[NaCl]/mol dm-3 as follows: 0, 10-3; 0, 10-2; 0, 10-1; a, 100. 9 8 7 PH 6 5 a Fig. 9. Henderson-Hasselbalch plot for bleached sulphate pulp. [NaCl]/mol dm-3 as foiiows: 0, 10-3; 0, 10-2; 0, 10-1; 0, 100.1534 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Saf. Thus for bleached sulphate pulp at pH 6.03 and 0.01 mol dm-3 NaCl: Sor.03) = - 2-59 C 8-1 Saf = -2.90 C 8-l a(6.03) = 0.893 and pK = pH +log [( 1 - a)/a] = 6.03 - 0.92 = 5.1 1. In this way a plot of pK against a was made and extrapolated back to a = 0 to find an estimate of pKo, the pK of the acid at zero dissociation. Using these calculated values for a Henderson-Hasselbalch plots were constructed. The variation of pK with a for cotton linters is shown in fig. 6 and the Henderson-Hasselbalch plot in fig.7. The corresponding plots for bleached sulphate pulp are shown in fig. 8 and 9. DISCUSSION COTTON LINTERS CHARGE/PH ISOTHERMS The form of the charge/pH isotherm in the pH range 2.5-7.0 is characteristic of a monofunctional polyacid. The charge increases as expected with increasing pH as the degree of dissociation of the acid increases. The charge also increases with increasing electrolyte concentration. This can be understood either in terms of the increasing capacitance of the double layer predicted by the differentiation of the Gouy-Chapman equation13 or by the decreasing apparent pK of the acid group. The precise location of the P.Z.C. is not clear, although a value between pH 1.5 and 2.5 is estimated from the isotherms. The 0.1 and 1 .O mol dmP3 isotherms appear to give a slightly higher P.Z.C.than the lower concentrations of electrolyte, and this is probably explained by specific adsorption of chloride at the higher concentrations of sodium chloride. The increase in charge after pH 7.0 may be explained in two ways. It may be that the swelling of the fibre, which occurs in alkali solutions, opens up new acid groups with which the alkali reacts. Alternatively there may be a second weak acid group on the cellulose surface. The behaviour of the 1.0 mol dm-3 NaCl isotherm is anomalous. Below pH 5.0 it lies in the expected order, above the 0.1 mol dm-3 isotherm. However, above pH 5.0 it drops to where it is almost coincident with the 0.01 mol dm-3 isotherm, and at pH 8.5 it levels to a plateau. This behaviour seems to suggest that both swelling and a second weak acid group may be involved.The swelling of the fibre is suppressed by the high electrolyte concentration, which explains the drop in charge of the 1 mol dm-3 isotherm, the effective number of acid groups having been reduced. This would also explain the plateau region above pH 8.5, if the fibre in the lower concentrations of NaCl continued to swell and therefore react after this group had been neutralized. Thus the 1 mol dm-3 isotherm may be considered to be the system in which any complicating effects of swelling are removed. POLYELECTROLYTE PLOTS The straight lines of the Henderson-Hasselbalch plots are indicative of a poly- electrolyte in which no detectable conformational transitions occur. The slopes of unity for the 0.1 and 1.0 mol dm-3 plots indicate that at these NaCl concentrations the system is behaving like a simple monobasic acid, such as acetic, for which n is always unity.The values of n and pK,, at other electrolyte concentrations are given in table 2.T. M. HERRINGTON AND B. R. MIDMORE 1535 Table 2. Values of n and pK,, obtained from the Henderson-Hasselbalch plots for cotton linters and bleached sulphate pulp [NaCl]/mol dmP3 n 0.00 1 0.01 0.1 1 .o cotton linters 1.97 1.44 1 .o 1 .o 5.15 4.20 3.95 3.65 bleached sulphate pulp 0.001 2.05 5.10 0.01 1.92 4.35 0.1 1.58 3.85 1 .o 1.26 3.75 The increase in n with decreasing electrolyte concentration is as expected and reflects the increasing surface potential with decreasing electrolyte concentration. The increase in pK,, is also a reflection of this phenomenon.The form of the pK against a plots is again indicative of a polyelectrolyte in which there is no detectable conformational transition. Such transitions generally show up as local minima in the pK against a plots and it is evident that no such minima exist. The monotonically increasing pK is a reflection of the increasing surface potential with increasing a. In Part 3 of this series it is shown how surface potential can be calculated from these plots. The 1.0 and 0.1 mol dmP3 NaCl systems both show a constant pK, consistent with the very low surface potential caused by the high ionic strength. The plots are extrapolated back to zero dissociation where pK = pKo. From the plots it is clear that the pKo values cluster around a value of ca.4.0; the values are lower, as expected, with increasing activity of NaCl. This is in good agreement with previous estimates obtained by other workers5$ and seems to be consistent with the hypothesis that the alkali adsorption is due to reaction with a carboxylic acid group which has replaced the primary alcohol group in the cellulose structure. BLEACHED SULPHATE PULP CHARGE/~H ISOTHERMS Like the isotherms for cotton linters, the bleached sulphate isotherms are also characteristic of a monofunctional polyacid. However, in this case it is clear that a much higher charge density is involved. This is confirmed by the fact that the neutralization point is achieved at a much higher pH. The precise location of the P.Z.C. is clearer because of the higher charge density and it is estimated to be at pH 1.75f.0.25.There is also clear evidence of charge reversal below pH 1.5. A mechanism for this reversal may be the protonation of the large numbers of oxygen atoms within the cellulose structure. As with the cotton linters, there is a degree of reaction above pH 8 after neutralization has occurred; however, here it is less marked because of the overall high charge density. Again, it is interesting to observe that the 1 mol dm-3 isotherm drops below the 0.1 mol dmP3 isotherm and becomes coincident with that for 0.01 mol dm-3 NaCl. This again is good evidence of the high electrolyte concentration suppressing swelling and thereby reducing the surface charge. The1536 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE e L 0 0 5 1 0 Q Fig.10. Variation of pK with a for various cations on bleached sulphate pulp. A, 0.05 rnol dm-3 BaCI,; 0, 0.05 mol dm-3 CaCI,; 0, 0.1 mol dm-3 NaCl; A, 0.1 mol dm-3 KCl. higher crossover point of pH 6.1, compared with pH 5.0 for cotton linters, also indicates a higher charge density, electronic effects dominating swelling effects until a higher pH. When fig. 5 is studied it is clear that the position concerning the effect on the isotherm of changing the cation is not very straightforward. Perhaps the most striking feature is the calcium isotherm, which after a value of - 2.66 C g-l at pH 5.46 remains constant to within k0.004 C g-l until pH 9.22. This is also the lowest value for So, above pH 6.5 in spite of the high affinity of calcium for carboxy groups.It is suggested that this behaviour is caused by the calcium ions suppressing swelling. This suppression cannot be due simply to the bivalency of the calcium ion, as no such effect is seen with barium. The effect is specific to calcium and probably indicates substantial adsorption of calcium ions into the Stern layer, thereby considerably reducing interlamellar repulsion by lowering ~ 6 . Here the size variations of the hydrated cation are not the dominant factors in determining So,, unlike the silica system studied by Tadros and L~k1ema.l~ They found that at high pH and high salt content, the OH- adsorption increased beyond the surface density of silanol groups. Most of the charge is neutralized by cations penetrating into the pores of the silica-gel structure.The extent of penetration depends on the size of the cation as OH- adsorption increased in the order (C,H,),N+ < Li+ < Na+ < K+ < Cs+. For the cellulose adsorption isotherms it is seen that the OH- adsorption for 0.1 mol dm-3 (C,H,),NCl rises above the potassium isotherm above pH 6.5. Again for silica the charge, for example, at pH 9 for 0.01 mol dm-3 (C,H,), NCl is about one third that of the equivalent sodium chloride concentration, while for bleached sulphate pulp it is approximately the same. This must reflect the relatively large pores in the cellulose fibre compared with those in the silica sol. POLYELECTROLYTE PLOTS In a similar way to the cotton linters, the straight lines of the Henderson-Hasselbalch plots are indicative that again we have a system in which there are no detectableT. M.HERRINGTON AND B. R. MIDMORE 1537 conformational transitions. The larger values for n which lie in the range 1.26-2.05 also indicate greater surface charge density. The pK against a plots for bleached sulphate show a similarity to those for the cotton linters, the pK, values being close to those of the latter. Overall it is clear that the main contributor to the charge is the same as that for cotton linters, namely a carboxy group at the C, position of the glucose molecule. The pK against a plots for the various cations on bleached sulphate pulp are shown in fig. 10. They offer useful ways for comparing the differing affinity of the various cations for the carboxy group. The effect of swelling is eliminated as a,N ( i e .the amount of carboxy groups available) is determined individually for each cation. From the plots the affinity lies in the order Na+ < K+ < Ba2+ < Ca2+. The barium and calcium plots show similarities to the 1 mol dm-3 NaCl plots as they flatten off at a = 0.9. UNBLEACHED SULPHATE PULP The charge/pH isotherm for unbleached sulphate pulp is shown in fig. 4. The charge per gram at pH 7 is four times greater than for bleached sulphate pulp and some 25 times greater than that of cotton linters. There is no tendency to form a horizontal plateau with increasing pH, but on the other hand the charge does not show the rapid increase without limit of a silica The behaviour suggests that several acid groups are contributing to the surface charge. Unbleached sulphate pulp contains a considerable quantity of lignin, so that phenolic and other acid groups may contribute to the surface charge as well as carboxy groups. M. J. Jaycock and J. L. Pearson, J. Appl. Biotechnol., 1975,25, 827. M. B. Donnan, T. W. Healy and P. F. Nelson, Colloids SurJ, 1981, 2, 133. S.M. Neale and W. A. Stnngfellow, Trans. Faraday SOC., 1937, 33, 881. E. Heymann, Trans. Faruhy Soc., 1942,38, 209. G. F. Davidson, J . Text. Inst., 1948, 39, 87. B. R. Midmore, Thesis (University of Reading, 1983). ti M. R. Edelson and J. Hermans, J. Polym. Sci., Part C, 1963, 2, 145. * R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 2nd edn, 1959), p. 433. G. H. Bolt, J. Phys. Chem., 1957, 61, 1166. lo G. H. Parkes and P. L. deBruyn, J. Phys. Chem., 1962,66, 967. l1 A. Katchalsky and P. Spitnik, J. Polym. Sci., 1947, 2, 432. l2 A. M. Kotliar and H. Morawetz, J . Am. Chem. Soc., 1955, 77, 3692. l3 J. Lyklema and J. Th. G. Overbeek, J. Colloid Sci., 1961, 16, 595. l4 Th. F. Tadros and J. Lyklema, J. Electroanal. Chem., 1968, 17, 267. (PAPER 3/ 1620)

ISSN:0300-9599    

DOI:10.1039/F19848001525

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1984,80, 1539-1552 Adsorption of Ions at the Cellulose/Aqueous Electrolyte Interface Part 2.-Determination of the Surface Area of Cellulose Fibres BY THELMA M. HERRINGTON* AND BRIAN R. MIDMORE Department of Chemistry, University of Reading, Reading RG6 2AD Received 14th September, 1983 The surface areas of the fibres studied in Part 1, obtained by the method of negative adsorption, are reported. Corrections were applied for low surface potential using Gouy- Chapman theory and the charge/pH isotherms. The fibres were oxidised to increase the charge and surface areas obtained for oxidised and non-oxidised fibres were in good agreement. The surface area was also obtained for unbleached sulphite pulp which has a high charge. The results were compared with surface areas determined by B.E.T.nitrogen adsorption. All the pulps were initially in a ‘never-dried’ state and for the B.E.T. work they were specially prepared using solvent-exchange techniques such that all the water was replaced by dry pentane. The negative adsorption results for the unbleached sulphate pulp clearly showed the effect of overlapping double layers, and this was treated theoretically. The effect of beating the fibres, as in paper-making, on the surface area was also investigated by both methods as was the effect of drying the fibres once. In order to convert the surface charge obtained for cellulose fibres at different pH values in Part 1 from C 8-l to C m-2, the specific surface area is required. The accepted surface area for cellulose fibre obtained by B.E.T.nitrogen adsorption is < 10m2 g-l.l These authors obtained high values for the B.E.T. surface area of different types of pulp (100-300 m2 g-l) using a method of solvent exchange such that all the water in the pores of the fibre was replaced by dry pentane before drying the fibre for B.E.T. Schofield2 determined the surface area of jute fibres by the negative adsorption of chloride ions. He obtained values of between 130 and 200 m2 8-l for different sources of jute. It was decided to determine the surface area of different types of cellulose fibre by the two methods. One of the main advantages of the negative adsorption technique is that it eliminates all estimations of molecular sizes which are needed in other techniques like gas or dye adsorption. Since the double-layer theory assumes the ions to be point charges their size is immaterial.The degree of negative adsorption also becomes independent of the surface charge, once this has become sufficiently high. The fact that the measurements are carried out in aqueous electrolyte solution makes this technique very suitable for the evaluation of surface area for charge/pH isotherms as these are also determined in aqueous electrolyte solution. Hence ‘wet’ surface area is determined in contrast to a ‘dry’ surface area determined by gas adsorption. The main disadvantage of the technique lies in the generally small change in the concentration of co-ions. This was usually in the range of 1-2% for cellulose fibres. Also for these natural fibres, internal electrical double-layer overlap presents an added complication.Thus negative adsorption does not measure the smallest parts of the fibre and does not ‘see’ any roughness of the surface. Negative adsorption of course can only be used where there is a charged interface and an electrical double layer. 15391540 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE THEORY At the interface between a charged particle and the surrounding solution, the co-ions, bearing the same charge as the interface, are expelled from the surface, i.e. they are negatively adsorbed from the interface. Negative adsorption results in an increase in the concentration of that ion in the bulk solution. As this increase reflects the quantity of ions expelled from the interface, it can be related to the expelling surface area and can be used to determine this area.The interpretation of data in negative adsorption experiments depends upon a suitable application of the Gouy-Chapman theory for the diffuse component of the electrical double layer. This application was first made by S~hofield;~ he related the negative adsorption to electrolyte concentration and surface charge. Lyklema and van den Hu14 gave a more lucid derivation in which potential rather than charge was the primary variable. They derived the following equation for a z : z electrolyte : Ani St = - & 2/ni B[ 1 - exp (- zi elys/2kT)]-l ni where B = (zi2e2/2ekT$. (All symbols are defined at the end of the paper.) The specific surface is given by S = St/o, where o is the mass of solid used.If the potential, ly6, at the start of the diffuse double layer is high, then the surface area is given by For and The two An. St = 2 F Bdn,. (2) ni work with natural fibres it is convenient to use the excluded volume then experimentally observed excluded volume, AGhs, in natural fibres is made up of A V = Ani F/(ni o) (3) AV = S/B2/ni. (4) components, one due to the negative adsorption and the other due to what is essentially a steric effect. This is caused by the inability of ions to penetrate into the smallest pores of the fibre, which can, however, be penetrated by water. Thus AV = A&,, - A & ( 5 ) where A is a constant independent of electrolyte concentration. After this correction a plot of A V against 1 / B d n i has slope S if lys is effectively infinite.In a natural fibre where part of the surface area measured is internal, the effect of overlap between nearby double layers must be considered. De Haan5 successfully solved the mathematical problems involved in an analysis similar in conception to that performed by Schofield3 for the single double layer. This analysis has been clarified by using potential as the primary variable.s The ratio between the negative adsorption for a single double layer and that for an overlapping double layer is given by Also (7) Values for u,, k and dare first calculated and then eqn (7) is solved for ud. These values for ud and u, enable An;/Ani to be calculated from eqn (6).T. M. HERRINGTON AND B. R. MIDMORE 1541 CORRECTION FOR LOW POTENTIAL The factor [l -exp ( -zf ey8/2kT)] only becomes unity, as assumed for eqn (2), when tys is infinite.This of course is never achieved in a real physical situation, so in practice the magnitude of y8 must be taken into consideration. The factor for a 1 : 1 electrolyte comes to within 2% of its limiting value at ca. 200 mV. However, the degree of negative adsorption is not effectively altered once y8 has become sufficiently high. The correction factor can be estimated in the following ways. (1) The charge/pH isotherm may be used. ys is the potential of the double layer at the outer Helmholtz plane, which is related to the diffuse double-layer charge os by the Gouy-Chapman equation. From the titration experiments of Part 1 we know the charge density per gram, So,. One approximation is that 08 = o,, which assumes the absence of cation absorption in the Stern layer.However, in order to determine o,, S must also be known. This problem was solved by two different methods. (a) According to the Gouy-Chapman equation o,, = as = A d n i sinh (zieys/2kT) (8) assuming no cation adsorption in the Stern layer. But from eqn (1) and (3) for negative adsorption of chloride ion (9) S = A VBdn, [ 1 - exp (eys/2kT)]-l and so eliminating ys where q = - So,/(ABni A V ) and the correction factor is [l - 1/(2q- l)]-l. Each individual negative adsorption point can be corrected in this way, the correction factor depending on AV, nt and a value of So, from the charge/pH isotherm. So AV?, the excluded volume if the surface potential were infinity, is given / 1 \-1 by and S = AVF B d n i from eqn (10).A plot of AV," against l/Bdn, then gives S. This correction method has the disadvantage of exaggerating any experimental errors, as values of AV which are initially too high owing to experimental error will lead to correction factors which are also too high, and vice versa for values of AV which are too low. Thus the standard deviation of a series of AVF will be relatively high. To overcome this effect another correction technique was employed. (b) A set of estimated surface areas, S,, from the set of negative adsorption points was used to calculate 0, and hence tys. For each S , the generated surface area S was then calculated in the following way. From the Gouy-Chapman equation1542 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE where y = o,/(Az/n,).From the charge/pH isotherm, So, is known so y is replaced by p , where where Sn is calculated from eqn (4). Hence P = Soo/(ASnz/ni) (15) The intersect at the abscissa of the plot of ( S , / S ) - 1 against S , is an estimate for the best value of S . Another estimate is the minimum in the plot of the standard deviation of the generated surface areas against S,. AVF can then be calculated for each chloride concentration using this best value of S and eqn (16). (2) ws is increased to an effectively infinite value. In the case of the silver iodide and metal oxide sols this is easily accomplished by merely increasing the concentration of the potential-determining ions as there are a very large number of adsorption sites. In the case of cellulose fibres this can only be achieved by increasing the number of adsorption sites, either by chemical oxidation or by adsorption of charged species onto the surface.( 3 ) The assumption is made that tys is equal to the zeta potential. The zeta potential was determined using both microelectrophoresis and streaming potential techniques : this is discussed in Part 3. (4) ws may be estimated from an analysis of the So,/pH isotherm in terms of polyelectrolyte theory. This is also discussed in Part 3 . EXPERIMENTAL PREPARATION OF FIBRE The cotton linters, bleached sulphate pulp and unbleached sulphate pulp A were those used in Part 1. The unbleached sulphite and unbleached sulphate pulp B were obtained from Canada;' they were never-dried Black Spruce. All pulps were stored at 4-6 "C.The fibre was prepared in the same way as in Part 1 except that sodium chloride (0.1 mol dm-+) replaced hydrochloric acid as the electrolyte in which the fibres were soaked. The fibres were washed free of chloride as before. OXIDATION EXPERIMENTS Oxycellulose was prepared in two ways. (1) Oxidation at the C , position using chromate : 20 g of cellulose fibre were oxidised at room temperature in 0.2 mol dm-3 oxalic acid + 0.1 mol dmP3 potassium chromate solution (600 cm3) for 4 h. The fibre was washed thoroughly, soaked in 0.2 rnol dm-3 oxalic acid for 20 h and then reoxidised in 0.1 mol dm-3 sodium chlorate solution containing 3% acetic acid for a further 20 h. Finally the fibre was washed thoroughly. (2) Oxidative cleavage of the glucose ring using periodate: 20 g of cellulose fibre were soaked in 1 dm3 of 0.05 mol dm-3 potassium periodate solution in the dark for 48 h.The fibre was thoroughly washed and then soaked in 0.2 mol dmP3 sodium hypochlorite solution acidified to pH 3 with acetic acid, in the dark. Finally the fibre was well washed. DETERMINATION OF NEGATIVE ADSORPTION OF CHLORIDE ION The principle and method used to determine the negative adsorption were the same as that used in the single-point titration method for the determination of acid concentration. However, fibre in its sodium form was used, which, when dispersed in sodium chloride solution (2 x 10-4-0. 1 mol dm-3), gave a pH of ca. 6.5. The slurries were left overnight at 25 "C to ensure attainment of equilibrium, and the supernatant was analysed for chloride potentiometrically using a silver, silver chloride electrode and silver nitrate as the titrant.A glass electrode, keptT. M. HERRINGTON AND B. R. MIDMORE 1543 only for these titrations, was used as the reference electrode. The titration was carried out at pH 3, using dilute nitric acid, to prevent potential drift. All adsorption data were determined with reference to a blank, the moisture in the fibres being taken into account. A&,, was calculated from the increase in chloride concentration, Ani, using A &,s = Ani &/(ni 0). (17) The value of A%, the steric effect, or the negative adsorption not due to double-layer effects, was determined by carrying out the experiment at pH 1.3 (0.05 mol dmP3 HCl), where the surface potential is zero and there is no double layer.Then A 5 = A &,,(pH 1.3) (18) and A V is calculated from eqn (5). DETERMINATION OF SURFACE AREA BY NITROGEN ADSORPTION The water in the fibre was replaced by organic solvents so that on drying the fibre the pores did not collapse. Dry AnalaR methanol and pentane were prepared by refluxing over iodine + magnesium and sodium, respectively, followed by distillation. For the solvent exchange, 1 g of fibre was dispersed in distilled water and sucked dry on a Buchner funnel. A pad of fibre was formed in a vessel, from which damp air was excluded, and 500 cm3 of dry methanol passed through the pad over a period of 3 h followed by 500 cm3 of dry pentane. The solvent-exchanged fibre was then bottled in dry pentane.The B.E.T. surface area by nitrogen adsorption was determined by Dr A. McLeod of Brunel University using a Carlo Erba Strumentazione instrument. RESULTS Surface areas were determined for the cotton linters, bleached sulphate pulp and unbleached sulphate pulp, for which charge/pH isotherms are given in Part 1, both by negative and nitrogen adsorption. The surface areas of oxidised cotton linters and oxidised bleached sulphate pulp were also determined by the method of negative adsorption. The effect of drying on the surface area of the sample of unbleached sulphite pulp was determined by the negative-adsorption technique. The effect of drying on the surface area of the samples of unbleached sulphate pulp was investigated both by negative and nitrogen adsorption. The effect of beating on the sample of unbleached sulphate pulp B was investigated using negative adsorption.BLEACHED SULPHATE PULP The negative-adsorption results are shown in table 1. A value of 0.138 f 0.010 cm3 g-’ was determined for A V , by doing experiments in 0.05 mol dm-3 hydrochloric acid at sodium chloride concentrations of 0.05 mol dm-3 and zero. An error of f 0.01 cm3 g-l in A Gbs gives an error of f 0.014 cm3 g-l in A V. From fig. 1 the plot of A V against l/Bdni is not linear, indicating that t,us is too small. The low charge was corrected for in three ways, using the two procedures 1 ( a ) and 1 (b) and by oxidation with chromate and periodate as in method 2. Consider method 1 (a) : at a value of ni of 2.08 mol dm-3 and at pH 6.5 it is found from the charge/pH isotherm (fig.3 of Part 1) that So, is -2.80 C g-l. From table 1, AVis 0.347 cm3 g-l and using B = 5.19 x C mol-d m-1 then from eqn (1 1) q = 2.018. Hence from eqn (12), AVF, the excluded volume if the surface potential were infinite, is 0.517 cm3 g-l. Using this procedure for each point gave an average surface area of 128 24 m2 g-l (table 3). In method 1 (6) an approximate value of the surface area, S,, is calculated from A V using eqn (4). From this value and the So, value above, a value of S is calculated from eqn (1 5) and (16). A plot of ( S / S , ) - 1 against S , gave, from the intercept at the abscissa, a best value of S of 126 m2 g-l. From the minimum in the plot of the standard deviation of S against S,, the best mi mol-2 and A = 3.71 x1544 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Table 1.Negative-adsorption results for bleached sulphate pulp A Kbs A V (l/Bdni) A Vr A Vg ni/mol dmP3 / cm3 g-' / cm3 g-1 / cm3 m-2 / cm3 g-l / 1 0-2 cm2 g-' 8.81 x 24.0+ 1.0 10.2f 1.4 20.8 18 22.3 6.56 x low2 29.6 15.8 24.1 35 31.9 4.47 x 36.5 22.7 28.8 48 41.4 2.08 x low2 48.5 34.7 42.8 52 52.8 8.75 x 78.0 64.2 66.0 86 84.6 6.94 x 83.4 69.6 74.1 89 89.8 2.18 x 63.4 49.6 41.3 6 . 5 7 ~ lop3 103.3 89.5 75.2 iodate oxidised pulp 2.23 x 83.0 69.2 40.8 9.11 x 139.8 126.0 63.9 chromate oxidised pulp (l/B,/ni)/104 cm3 m-2 Fig. 1. Negative-adsorption results for bleached sulphate pulp : A, experimentally observed excluded volume, AV; 0, AVcorrected by method (b) (AVg; S = 127 m2 g-l); m, chromate oxidised pulp (S = 120 m2 g-l); -.---, Acheor [eqn (16) using S = S, = 127 m2 g-11.value of S was 127 m2 g-l. A value of AVF was then calculated for this value of S and each chloride concentration from eqn (4). These are plotted as open squares in fig. 1. A value of A V that would be experimentally observed taking into account the low potential is given by putting S= S , = 127 m2 g-l in eqn (16); this is called and is plotted as a dotted line in fig. 1.T. M. HERRINGTON AND B. R. MIDMORE 1545 Table 2. Negative-adsorption results for cotton linters ~~ A &bs A V (1/Bz/ni) A VF A VT n,/mol dm-3 /10-2 cm3 g-1 /lop2 cm3 g-l cm3 rnP2 cm3 g-' cm3 g-l 2.02 x 10-2 14.2 _+ 1.0 9.9 _+ 1.4 43.4 33.9 23.1 4.33 x 10-3 28.1 23.8 98.4 33.2 37.5 3.22 x 39.2 34.9 108.8 53.9 55.2 9.80 x 76.9 72.6 197.2 91.9 92.1 8.52 x 83.7 79.4 21 1.5 99.4 99.6 8.80 x 62.0 57.7 65.0 2.29 x 43.8 39.5 40.3 oxidised by IO;/ClO- The values of AV for the chromate-oxidised pulp (table 1) are plotted in fig.1 as filled squares; they seem to fit eqn (4) and give a very similar surface area (see table 3 below) to the corrected values for the non-oxidised pulp, implying that the surface potential is high. In support of this, the surface charge was determined by titration in 0.01 mol dm-3 NaCl to be 13.0 C g-l at neutralisation, which provides a high surface potential of 150 mV. The periodate oxidation caused a degree of degradation to the fibre, which almost certainly accounts for the high surface area (table 3). The charge density was determined to be 17.0 C g-l at neutralisation.The surface area of the unoxidised bleached sulphate pulp was determined by nitrogen adsorption to be 190 m2 g-l (see table 5 later). COTTON LINTERS The negative-adsorption results for cotton linters are given in table 2. (The value of A& was determined to be 0.043 +O.OlO cm3 g-l.) From fig. 2 it can be seen that they again imply a low surface potential. The surface area was calculated using methods 1 (a) and 1 (b) to give values of 51 and 48 m2 g-l, respectively, (table 3). Oxidation with IOy/ClO- gave a much higher value of 94 5 m2 g-l, possibly because of degradation of the fibre. However, nitrogen adsorption gave a yet higher value of 130 m2 g-l (table 5). UNBLEACHED SULPHITE PULP The effect of drying the sample of never-dried unbleached sulphite pulp was investigated by the negative-adsorption technique.The wet pulp was dried in an air oven at 104 "C for 24 h and then redispersed in water. A& was determined to be 0.15 & 0.01 cm3 g-l, which is the same as unbleached sulphate pulp. The plots of A V against l / B z / n , are shown in fig. 3 for the never-dried and the once-dried pulps. The surface areas were corrected for 'low charge' by method 1 (a), but the corrections were small. The charge density at neutralization was 14.9 C g-l. The never-dried pulp gave a value of 210 m2 g-l, which reduced to 122 m2 g-l for the once-dried pulp (table 3). UNBLEACHED SULPHATE PULP From the negative-adsorption experiments, AK = O.lS+O.Ol cm3 g-l. From the plot of A V against 1 /Bz/n, (fig. 4) it can be seen that this pulp clearly shows the effect of overlapping double layers for chloride concentrations < 1 x lop2 mol dm-3.The pore volume, 5, was estimated by extrapolating the plot of A V against l/Bdni to l/Bz/n, = 0 0 . ~ The value of S was calculated from the negative adsorption at the 51 FAR 11546 1.0- 0.8 - 0.6- I M m E . 2 Q 0.4 - 0.2- THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE / I / /’ 80 160 240 0 (1 /B&.)/ 1 o+ cm3 Fig. 2. Negative-adsorption results for cotton linters : A, experimentally observed excluded volume, A V ; 0, A V corrected by method (b) (A V c = 48 m2 g-l). (l/B,/ni)/lO+ cm3 m-2 Fig. 3. Negative-adsorption results for unbleached sulphite pulp : 0, never-dried pulp; 0, once-dried pulp.T. M. HERRINGTON AND B. R. MIDMORE 1547 Table 3.Comparison of the surface areas of different fibres obtained by negative adsorption method 1 method 2 Pulp (4 (6) CrOt- IO,/ClO- bleached sulphate 128 & 24 126+ 10 120+ 1 184+ 14 unbleached sulphite cotton linters 51 f 14 48+5 - 94i- 5 never dried 210 f 10 once dried 122+5 never dried 210+9 once dried 110+4 unbleached sulphate; sample A 16C 12c - I 04 m E PI 0 80 6 - 2 40 1 0 80 160 (l/B,/ni)/104 cm3 m-2 Fig. 4. Negative-adsorption results for unbleached sulphate pulp : 0, Never-dried pulp; a, Once-dried pulp; ----, Acheor (table 4) for box-shaped pores. highest chloride concentration, where the overlap should be least, using eqn (4). From these two values the average separation between the two double layers, 2d, was estimated assuming the pores to be box-shaped by d = VJS. (19) Values for V, of 1.48 cm3 g-l and for S of 210 m2 g-l [obtained by correction method l(a)] give d = 70.1 A.From this value of d the theoretical negative adsorption could 51-21548 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Table 4. Effect of double-layer overlap on the negative adsorption of unbleached sulphate pulp 2.06 x 71.2 42.5 21.2 8.80 x 10-3 102.1 65.0 32.4 6.94 x 10-3 112.0 73.2 36.5 2.82 x 10-3 132.8 114.8 57.3 9.50 x 10-4 141.3 197.8 98.7 0.207 0.1350 0.1204 0.0774 0.045 1 'd 2.06 x 0.840 98.6 100 70.5 8.80 x 10-3 0.473 84.8 86.5 100.1 6.94 x 10-3 0.385 79.1 82.8 107.0 2.82 x 10-3 0.158 54.7 59.7 121.7 9.50 x 10-4 0.0527 32.5 35.6 129.1 be calculated from eqn (6) in the following way. For ni = 6.94 x mol dm-3 (table 4) and pH 4.5, So, is given by fig.4 of Part 1 as 8.20 C g-l. A value for d u o is calculated from the Gouy-Chapman eqn (14). Rearranging eqn (7) gives rcd+22/uddu, = 22/UdF(Ud,Z/2). (20) This eqn was solved for ud. For this value of ud, the elliptic integrals in eqn (6) were found from tables and hence (An;/Ani)theor was calculated (table 4). These values are compared in table 4 with values of (Ani/Ani)obs obtained by assuming no double-layer overlap at the highest chloride concentration. Using this value of S, eqn (4), Ani was calculated at the other chloride concentrations; An: is the experimentally measured negative adsorption. As can be seen from table 4, the agreement between the two sets of values for An;/Ani is close, lending credance to the theoretical model.In further support of the pore model, A&heor was calculated from (An;/Ani)theor; this is plotted as a dotted line in fig. 4 and shows the same trend as the experimental AV. The effect of drying upon the surface area of this 'never-dried' pulp (sample A) was investigated as for the unbleached sulphite. The surface areas were corrected for low charge by method 1 (a). As can be seen from fig. 3 and 4, their behaviour is similar, dropping from ca. 210 m2 g-l to ca. 110 m2 g-l after drying (table 3). The effect of beating and drying on the second sample of unbleached sulphate pulp (sample B) was investigated both by the negative-adsorption technique and by B.E.T. nitrogen adsorption. Sample B had a slightly lower charge at neutralisation than sample A (6.1 C g-l compared with 8.2 C g-l).It was beaten in a standard PFI mill to three degrees of beating. The extent of beating was defined by a Schopper number (determined by the Schopper-Reigler method). These samples were then dried as before. The surface area obtained by negative adsorption was determined at two chloride concentrations (1 x and 2.5 mol dm-3) and corrected for low charge by method 1 (a); the average value was taken. The surface areas by the two methods are compared in table 5. The two values of the nitrogen surface area for the never-beaten, never-dried pulp B are for the same sample but two different runs of solvent exchange.T. M. HERRINGTON AND B. R. MIDMORE 1549 Table 5. Comparison of the surface areas of different pulps obtained by negative and nitrogen adsorption by by - scr, negative nitrogen /c g-' adsorption adsorption ~ bleached sulphate 2.80 cotton linters 0.46 unbleached sulphate; sample B (1) zero beating, 15 O SR 6.12 (2) beaten to 38.5 O SR (3) beaten to 78 O SR- (1) zero beating, 15 O SR (2) beaten to 38.5 O SR (3) beaten to 78 O SR - - unbleached sulphate; sample B; once-dried 5.02 126+ 10 190 48f5 130 207 & 9 265a, 240a 216f 10 - 226& 10 - 112+6 120 122f7 - 137+7 - a Same sample but for two different runs.The pore size distribution from B.E.T. nitrogen desorption gave a broad band of porosity 22-60 A in radius, which is comparable with our d value of 70 A for the wet pulp. This increase is probably explained by swelling of the pulp in electrolyte solution. DISCUSSION The linear form of the negative-adsorption plots seems to be good quantitative evidence for the applicability of electrical double-layer theory to the cellulose/water interface.It suggests that a double layer does exist which is reasonably well described by the Gouy-Chapman equation. The linearity of the plots increases with increasing charge density, which is as expected since lya is closer to being effectively infinite over a greater range of chloride concentration. The oxidised bleached sulphate (fig. 1) and the unbleached sulphite (fig. 3) pulps show this linearity well. In the case of the bleached sulphate pulps and cotton linters, because of their relatively low charge densities, lya cannot be described as effectively infinite. The negative adsorption therefore declines considerably as the increasing electrolyte concentration reduces lya as given by eqn (8).This reduction seems to agree well with the theoretical reduction predicted from eqn (9), as shown by the dotted line of A Kheor in fig. 1 ; the slope decreases as the concentration of chloride ion increases. This again is not only good evidence for the applicability of the negative adsorption aspects of double-layer theory but also more generally for the Gouy-Chapman equation, as this is involved in the estimation of lya. Both of the correction procedures applied to account for a low surface potential seem to work quite well. The estimated surface area for the oxidised pulp of 126 m2 g-l agrees within experimental error with that obtained for the oxidised pulp (120 m2 g-l) where the surface potential was raised effectively to infinity by chemical means.The effect of double-layer overlap is shown clearly by the negative-adsorption plot for the unbleached sulphate pulp (sample A) in fig. 4. The theoretical negative adsorption AKheor of table 4 is plotted as a dotted line in the figure and agrees1550 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE reasonably well with the experimental values. Exact agreement was unlikely, as the calculations are based on a single double-layer separation value of d which is a simple numerical average obtained by assuming box-shaped pores. This situation is most unlikely within the complex porous structure of the cellulose fibre. However, the results do show that both the calculated surface area S and the pore volume are at least of the right order of magnitude, and hence give a reasonable estimate for d.Indeed for an unbleached sulphate pulp has been independently determined by Stone and Scallan.8 They determined the negative adsorption of dextran molecules with sufficiently high molecular weight that they were excluded from the interior of the fibre. A value of 1.4 cm3 g-l was obtained, which compares very well with the value of 1.48 cm3 g-l obtained in this work. They found a higher pore volume for sulphite pulp (ca. 1.7 cm3 g-l); from fig. 3 it can be seen that the pore volume for unbleached sulphite pulp must be > 1.8 cm3 g-l. It is clear that in every case drying the fibre at 104 "C for 24 h and redispersion in water considerably reduced the surface area.In general there is a reduction by about a half (tables 3 and 5). This reduction must be caused by permanent closure of some of the pores in the fibres. They are permanently closed because they are not reopened on redispersal in water. There is evidence that these pores are small pores. The surface area of the unbleached sulphate pulp (B) obtained by nitrogen adsorption is 30% larger than that obtained by negative adsorption (265 m2 g-l compared with 207 m2 g-l; table 5). This is presumably because the negative-adsorption technique does not de- tect the smallest pores. However, after one drying the difference reduces to only 7% (120 m2 g-l compared with 112 m2 g-l; table 5). This may be explicable by the closing up of the smallest pores which are now no longer accessible to either technique.From the results in table 5 on the unbleached sulphate pulp (B), beating increases the surface area of the fibres, but not dramatically so. This is considered to be achieved by the increased external fibrillation of the fibre^.^ The fact that the reduction of surface area on drying for all three degrees of beating is the same indicates that this loss in surface area is due to reduction of internal and not external fibrillation. The negative-adsorption surface-area values compare reasonably well with those obtained by nitrogen adsorption. They are all, however, lower than the latter, which is as expected for the internal surface of a porous material. This effect was found by Tadros and LyklemalO for the surface area of silica gel, where the surface area obtained by negative adsorption was 35 m2 g-l and that by nitrogen adsorption was 50 m2 g-l.This represents a reduction of 30%. So 30% of the surface (that due to the smallest pores) was not available for negative adsorption. This compares well with the results obtained for cellulose, where for the unbleached sulphate pulp 22% of the surface is undetected by negative adsorption and 34% in the case of bleached sulphate pulp. The 'wet' value for the cotton linters (48 m2 g-l) is, however, considerably lower than the nitrogen-adsorption value (130 m2 g-l). This is accounted for by the very low charge density of the cotton linters. At these low charge densities (4 x C m-2, calculated from the charge/pH isotherm with S = 130 m2 g-l) the double-layer theory, which assumes an indiscreet and smeared surface charge density, can no longer apply. The charges have a separation of 63 A at this char e density, compared with pulp (B) (2.3 x C m-2).The degree of effective shielding the discreet charges will experience will be greater than that of a ' smeared ' charge, and so the effective surface potential will be considerably lower. Thus our method of correction using Gouy- Chapman theory cannot apply, as it relies on an estimation of the surface potential that is too high; from this method a surface area which is too low is obtained. The only other negative-adsorption study of cellulose fibres is that of Schofield2 on jute. 32 A for bleached sulphate pulp (1.6 x C m-2) and 26 x for unbleached sulphateT.M. HERRINGTON AND B. R. MIDMORE 1551 He obtained a surface area for jute from different sources of 150-200 m2 g-l, which is the same range as the surface areas for cellulose fibres in this work. The nitrogen- adsorption surface areas obtained by Stone and Scallans also compare well with this work; four of the pulps they studied had surface areas in the range 182-230 m2 g-l. They also found a loss of surface area on bleaching; the surface area of unbleached sulphite pulp dropped from 182 to 93 m2 g-l after bleaching. This is paralleled by the lower surface area of bleached sulphate (1 26 m2 g-l negative adsorption, 190 m2 g-l nitrogen adsorption) compared with unbleached sulphate (2 10 m2 g-l negative adsorp- tion) in this work. Probably the digestion of finely divided areas in the fibre occurs in the bleaching process.It is considered that the negative-adsorption technique offers the best approximation for the surface area in the context of the adsorption of ions from solution on to the cellulose surface as in determining chargelpH isotherms. A B d E e F ni Ani P 4 S Sn st K A Vobs AKi Ud UO AV A VF A V g Y zi u) & K *b 0 0 Wb W O ryd GLOSSARY OF SYMBOLS = ( 8 ~ k T ) ; = (z2e2/2ckT); half the distance apart of two overlapping double layers elliptic integral of the first kind charge of electron elliptic integral of the second kind concentration of co-ion i in bulk solution increase in concentration of co-ion i in bulk solution = - Soo/(ABni A V ) surface area per gram estimated surface area total surface area = exp ( -zi etyd/kT) = exp ( - zi ey/,/kT) total volume of solution excluded volume, A V = Ani VJ(ni co) experimentally observed excluded volume excluded volume accessible to water but not ions excluded volume corrected for low potential by method 1 (a) excluded volume corrected for low potential by method 1 (b) mass of solid = *o/(Adni) valency of ion i, sign included electrical permittivity of the medium Debye reciprocal length, K = ( 2 4 zf e2/&kr): charge per unit area at Stern plane charge per unit area at surface potential at Stern plane potential at surface potential at midpoint of two overlapping double layers = S~oI(ASndni)1552 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE M. B. Donnan, T. W. Healy and P. F. Nelson, Colloids Surf., 1981, 2, 133. R. K. Schofield and 0. Talibuddin, Discuss. Faraday SOC., 1948,3, 51. R. K. Schofield, Nature (London), 1947, 160,408. H. J. van den Hul and J. Lyklema, J. Colloid Interface. Sci., 1967, 23, 500. F. A. M. de Haan, J. Phys. Chem., 1964, 68, 2970. B. R. Midmore, Thesis (Reading University, 1983) communication. J. E. Stone and A. M. Scallan, Pulp Paper Mag. Can., 1965, 66 (8), 407. ' A. M. Scallan (Pulp and Paper Research Institute of Canada, Point Claire, P.Q., Canada), personal * J. E. Stone and A. M. Scallan, Tappi, 1967,50 (lo), 496. lo Th. F. Tadros and J. Lyklema, J. Electroanal. Chem., 1968, 17, 267. (PAPER 3/1621)

ISSN:0300-9599    

DOI:10.1039/F19848001539

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1984, 80, 1553-1566 Adsorption of Ions at the Cellulose/Aqueous Electrolyte Interface Part 3.-Calculation of the Potential at the Surface of Cellulose Fibres BY THELMA M. HERRINGTON* AND BRIAN R. MIDMORE Department of Chemistry, University of Reading, Reading RG6 2AD Received 14th September, 1983 The zeta potential, (, for cotton linters and bleached sulphate pulp has been calculated from measurement of the streaming potential. The measurements were taken at the same electrolyte and pH conditions as the charge/pH isotherms of Part 1. A comparison is made of the surface potential calculated under the same conditions from the charge/pH isotherms, negative adsorption experiments, polyelectrolyte theory and streaming potential measurements. It is shown that the [ potential is not a good relative measure of the surface charge and cannot be used for qualitative comparison between such similar materials as bleached sulphate pulp and cotton linters.Since the discovery of electrokinetic phenomena in the early nineteenth century, extensive data have been accumulated in this field. This has been partly for historical reasons and partly because of the relative ease with which electrokinetic phenomena may be studied. Electrokinetic phenomena involve the tangential displacement of a liquid along what is known as the slipping plane. From theoretical considerations it is possible to calculate the potential at this slipping plane from measurement of electrokinetic phenomena. This potential is known as the zeta potential. The zeta potential, [, has a different character from v0, the surface potential, generally being considerably smaller in magnitude.It has often been approximated to vS, the potential of the diffuse double layer, implying that the slipping plane coincides more or less with the first layer of adsorbed i0ns.l However, c only gives exact quantitative information about the nature of the double layer if the position of the slipping plane is known, which is in turn only determinable from knowledge of both [ and vS. It might appear from these considerations that the extensive data accumulated in the field of electrokinetics were unjustified, taking into account that the absolute value of [ only has relevance within the context of electrokinetic phenomena. However, it has always been thought that c has a semi-quantitative significance, potentials within a given system (or different types of similar material) being comparable with each other.The ‘isoelectric point’, where [ = 0, is also determinable and allows a comparison with the point of zero charge from the charge/pH isotherms. Moreover, the measurement of the [ potential gives information about the nature of the diffuse layer which is totally independent of adsorption phenomena and therefore provides a very useful and independent comparison with values of I,U determined from charge/pH isotherms or negative-adsorption data. THEORY When a liquid is forced through a compacted plug of particles a convection current is produced because of the removal of excess counter-ions from the double layer by the 15531554 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE flowing fluid.In order to maintain electrical neutrality this convection current is balanced by a conduction current in the opposite direction, which induces the ‘streaming potential’ between either side of the plug. The interpretation of these streaming potentials depends upon the following assumptions : (1) only external fibre surfaces are involved, (2) each portion of fibre is electrokinetically identical, (3) the flow is lamellar, (4) the velocity gradient or shear rate at the plane of shear is constant throughout the pad and ( 5 ) the internal pore network is statistically uniform throughout the pad. The Helmholtz-Smoluchowski equation for streaming potential is given by where is the pore factor. (Other terms are defined at the end of the paper.) It is related to the conductivity of the bulk liquid and the conductivity of the surface by r = 1 /(&+2As/r).Eqn (1) applies to single capillaries, and its extension to porous plugs presents problems when the surface conduction becomes important, i.e. at low concentrations of electrolyte. The surface-conduction problem can be overcome by measuring streaming current since surface conduction is not involved : I, = - ~ a P c / q L . (3) This formula would apply to plugs composed of granular particles of fairly uniform size and shape. Chang and Robertson2 worked with compressible plugs of beaten and unbeaten cellulose fibres, nylon, dacron, etc., and found that a plot of In (I, q/&P) against c was linear, i.e. (4) where c is the concentration of solid material in the plug.In terms of the streaming potential ( 5 ) so that the pore factor c is given by c = exp (-Jc)/(A,+ 2Ls/r), and the slope of a plot of E, against P would depend on the density of the plug. However, we found that it was only possible to alter the packing density of the plug within narrow limits, but within these limits the slope was independent of the plug density at constant electrolye concentration. Thus it was decided to use eqn (2) and to eliminate < by the method of Brigg~.~ At high concentrations of electrolyte (0.1 mol dm-3 KCl) < = l/Ao = R,/L. At low electrolyte concentrations Is = - EaPc exp (- Jc)/qL ESP = E c exp ( - Jc)/h(A, + 2AsIr)l = l/(Ao+2As/r) = R/L.Then Es/P = ERc/vA,, R, = ERc/qK (4) where K is the cell constant. EXPERIMENTAL The potassium chloride and sodium chloride used were A.R. grade. The hydrochloric acid and sodium hydroxide solutions used for adjusting the pH were prepared as in Part 1, as were the fibre slurries. The streaming-potential apparatus is described in detail el~ewhere.~ It consisted of two reservoirs connected by two-way taps to the streaming-potential cell. The streaming solution was forced through the base of one reservoir, through the cell and into the other reservoir via a vertical glass tube, which ensured that the hydrostatic head opposing theT. M. HERRINGTON AND B. R. MIDMORE 1555 flow of solution remained constant. Two smaller vessels were connected to the side of each reservoir and enabled the pH, conductivity and temperature of the solution to be determined.The solution was forced from one reservoir to the other using C0,-free nitrogen. All joints were either Teflon-sleeved cone-and-socket or else Sovril. The fibre was formed between two perforated platinum discs covered with nylon mesh. Platinum wire attached to the discs was used to measure the resistance, R, of the pad. Ag,AgCl electrodes were positioned either side of the platinum discs to measure the streaming potential. These were the same as used in the titration work. The streaming potential was determined using a Vibron electrometer model 62 A (Electronic Instruments, Richmond, Surrey). The conductivity of the solution and pad was determined by means of a Wayne-Kerr bridge, model B605 (Wilmot-Breeden Electronics, Bognor Regis, Sussex) and the pH of the solution was determined using a pH meter, model pHM 64 (Radiometer, Copenhagen), with glass and calomel electrodes. The fibre pad was formed from a slurry in demineralized water.The settling of the fibre was accelerated by suction; care was taken to keep the pad wet and free of air bubbles. Fibre concentrations of ca. 0.15 g The cell constant of the fibre pad was determined by repeatedly passing fresh KC1 solution (0.1 mol dm-3) through the pad until the resistance, R,, of the pad remained constant. The temperature of the solution was recorded. From tables of the conductivity of KCl solutions at that temperat~re,~ the cell constant, K, of the pad was calculated. To measure the streaming potential, the 0.1 mol dm-3 KCl solution was replaced with the streaming solution mol dm-3 NaCl).Again the plug was washed with solution until a constant resistance of the plug was recorded. Some of the solution was then forced into a side vessel and its pH and temperature were determined. The conductivity was also determined in the side vessel as a check to ensure adequate washing of the plug. The pH was varied using sodium hydroxide and hydrochloric acid. All the plots of E, against P were linear. The electrophoresis measurements were carried out using the Rank mark I1 electrophoresis apparatus (Rank Bros, Bottisham, Cambridge). The silica flat cell, fitted with a double- platinum-electrode system, was used. The cellulose fines were obtained by beating a 1 % slurry in a single-rotor blender; after filtering, the fines were collected by centrifugation and kept in wet storage at 4-6 "C. were used.mol dm-3 or RESULTS Streaming-potential measurements were made for cotton linters and bleached sulphate pulp over the whole pH range (2-10.5) in mol dm-3 NaCl. The zeta potential was determined from the linear plots of E, against P. Data for the whole pH range in 10-2moldm-3 NaCl were also determined for cotton linters. The variation of c with pH is shown in fig. 1 for cotton linters and in fig. 2 for bleached sulphate pulp. Electrophoresis data were also determined for cotton linters, and the zeta potential so obtained is shown by way of comparison in fig. 1. The data are in agreement within experimental error.the variation of potential with distance from the surface is given by According to the analysis of Gouy and exp ( - K X ) = tanh [zery(x)/4kTl/tanh (zeryd/4kT) (7) where ~ ( x ) is the potential of the double layer at a distance x from the surface, which is equal to c at the plane of shear. Now as is low for cellulose fibres (< 25 mV) then tanh (ze[/4kT) ze[/4kT. (8) Thus at pH a exp ( - Icx,) = sit, 4k T (9)1556 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE / 1 I4 12- 10. 0 . -S/mV - 6. 4 - 2 - 0.. PH Fig. 1. Variation of the zeta potential with pH for cotton linters from measurements of streaming potential. 0 , l 0-3 mol dm-3 NaCl ; 8 , l 0-2 mol dm-3 NaCl ; , 1 0-2 mol dm-3 NaCl (micro-electrophoresis data) ; ----, [, values (see text). - 0- _ _ - - - 4 - - - / / //020Y /’ .6’ /4.r P ” / / / / , A I . 2 3 4 5 6 7 8 9 10 I I where t , = tanh (zey/f/4kT); x,, [, and y/$ refer to the distance of the plane of shear, the zeta potential and the surface potential at pH a. Similarly for a second pH at the same effective ionic strength, pH b:T. M. HERRINGTON AND B. R. MIDMORE 1557 Table 1. Electrokinetic data for cotton linters in mol dm-3 NaCl from streaming- potential measurements 10.64 12.0 9.52 11.0 7.98 9.3 6.42 8.9 4.83 7.5 4.34 6.4 3.09 2.0 2.24 0.3 - - 24.0 1.563 13.9 - - 18.8 1.188 10.6 - - 17.2 1.083 9.6 - - 1 5.9a 1.000 8.9 0.843 4.10 11.7 0.729 6.5 0.719 3.93 8.9 0.552 4.9 0.225 3.63 3.4 0.208 1.9 0.034 3.70 a vs(NA) = - 10.5 mV at pH 6.5 estimated from negative-adsorption experiments assuming S = 130 m2 g-l. If the position of the plane of shear does not change on altering the pH, then x, = X b and This equation enables a quantitative comparison between zeta potential and surface potential without an accurate estimate of c necessarily being known, as this is eliminated in the ratio.Also, if tya is low (a condition that is fulfilled in the case of cotton linters in lov2 mol dm-3 NaCl where rys < 25 mV) then [eqn (14) below]. Thus if the acid group is totally dissociated at pH b, then l a / T b = a (13) where a is the degree of dissociation at pH a. This allows values of pK to be calculated at different values of the pH. A plot of pK against a gives pKo from electrokinetic data. The values of pK are given for cotton linters in mol dm-3 NaCl in table 1.A plot of pK against a extrapolates to a pKo of 3.65. [This compares with a pKo of 4.05 for mol dm-3 NaCl (fig. 6 of Part 1) using values for a calculated from the charge/pH isotherm.] Alternatively values of aa were obtained from the chargelpH isotherm for cotton linters in mol dm-3 NaCl (fig. 2 of Part 1). Taking cb as - 8.9 mV at the point of neutralisation (pH 6.42), then la can be calculated from eqn (13). These values are also given in table 1 and they are plotted as a dashed line in fig. 1. To show that the approximation for eqn (12) is valid, values of tys were calculated from the Gouy-Chapman equation : od = ( 2ni k Tc)42 sinh (zi e tya/2k T ) (14) assuming that 06 = oo and thus tya x ry,(GC). The nitrogen-adsorption surface area of 130 m2 g-l for cotton linters was used to estimate od from the Sao/pH isotherm (fig.2 of Part 1). As is shown in table 1, the values of rya are c 25 mV. mol dm-3 NaCl, tyo(GC) calculated in this way is much higher, and eqn (1 1) must be used to calculate la. These values are given in table 2 and are plotted in fig. 1. In the case of bleached sulphate pulp in mol dm-3 NaC1, For cotton linters in1558 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Table 2. Electrokinetic data for cotton linters in mol dm-3 NaCl using streaming- potential measurements 10.03 8.46 7.31 6.34 6.18 6.08 4.81 4.50 3.78 3.10 2.71 2.15 24.1 22.8 22.8 21.5 20.6 20.2 16.8 15.7 8.9 4.3 1.2 0.2 55.6 47.4 42.0 37.2b 36.3 35.5 21.7 15.7 6.9 4.3 1.147 1 .ooo 0.900 0.805 0.788 0.771 0.484 0.353 0.144 0.096 26.1 22.8 20.5 18.4 18.0 17.6 11.0 8.0 3.3 2.2 a vANA) = - 17.1 mV at pH 6.5 if estimated from negative-adsorption measurements assuming S = 130 m2 g-l.* tb = 0.4307. Table 3. Electrokinetic data for bleached sulphate pulp in lop3 mol dmP3 NaCl using streaming-potential measurements 10.31 8.55 6.53 5.82 4.78 4.08 3.44 2.51 13.9 13.5 12.4 12.2 10.4 7.6 4.9 1.6 110.8 107.5 99.9 89.2 69.7 42.4 15.8 2.2 1.015 1 .ooo 0.962 0.897 0.756 0.501 0.196 0.028 13.7 13.5 13.0 12.1 10.2 6.8 2.6 0.4 the nitrogen-adsorption surface area of 190 m2 g-' was used to calculate yS. Values of [, from eqn (1 1) are given in table 3 and plotted in fig. 2. Values of tyS (NA) were calculated from the negative-adsorption results at pH 6.5 for cotton linters in low2 and mol dmP3 NaC1. The values are given in tables 1 and 2.DISCUSSION COTTON LINTERS The value of the zeta potential in mol dm-3 NaCl and pH 7.0 agrees reasonably well with results obtained by other workers. Zeta potentials of between 25 and 30 mV have been previously determined for cotton.8 The suppression of the double layer onT. M. HERRINGTON AND B. R. MIDMORE 1559 increasing the electrolyte concentration is clearly seen by comparing the zeta potentials at lop3 and The low surface charge density on the cotton linters will almost certainly lead to an over-estimation of the effective surface potential, yd, using the Gouy-Chapman eqn (14). This problem is not encountered if yd is calculated from the negative- adsorption experiments as then only the ‘effective’ surface potential is registered.It is therefore interesting to note the better agreement between the zeta and the surface potential calculated from the negative-adsorption data at both sodium chloride concentrations. This would indicate that for cotton linters the zeta potential calculated from electrokinetic phenomena is a reasonable estimate of the effective surface potential. The zeta potential [, predicted from the charge/pH isotherms is shown as a broken line in fig. 1. The general form of both isotherms is the same. This is clearly shown in the case of 10-2mol dm-3 NaC1. Again this is additional evidence that the adsorption of hydroxide ions is responsible for the double layer. The ‘isoelectric’ point is also close to the point of zero charge (P.z.c.) (pH 2.0 compared with pH 1.5-2.5 for the adsorption isotherm), which is good evidence for the mutual source of the two phenomena.mol dm-3 NaCl. BLEACHED SULPHATE PULP The general form of the predicted zeta potential, CS, is close to that of the measured values (fig. 2), and together with a good comparison between the ‘isoelectric’ point and P.Z.C. (pH 2.0 and 1.7) this provides good evidence that the adsorption of hydroxide ions at the cellulose surface is the source of the surface charge and hence double layer. However, the zeta potential is very much smaller than the calculated surface potential. Indeed it is lower than in the case of cotton linters in spite of an increased surface charge density. This phenomenon, of the zeta potential being lowered on increasing the charge, has previously been observedg and has been used as an argument against adsorption of hydroxide ions being the source of the surface charge.1° The weight of evidence for the ionogenic source of the surface charge shown in this work discounts such an argument.Our work shows quite dramatically that the [ potential is not a measure of the surface charge and cannot be used for a comparison of the surface charge of even very similar materials. Note that the zeta potential is reduced as the swelling in the fibre increases. Thus the swollen volume of bleached sulphate pulp is 2.50 cm3 g-l compared with 1.10 cm3 g-l for cotton.1° It seems probable therefore that, because of the swelling of amorphous cellulose or perhaps carboxylated cellulose chains at the surface, the position of the shear plane is shifted further out into the double layer, thereby reducing the zeta potential.This model was first postulated by Goring and Mason.ll Indeed such micro-fibrils at the surface of cellulose have been detected in electron-microscope images of the surface of unswollen fibres. A value of 8 nm for the position of the shear plane was calculated from eqn (7) for bleached sulphate pulp in lop3 mol dmp3 NaCl at pH 6.5 and the data of table 3. SWELLING IN CELLULOSE FIBRES Grignon and Scallan12 interpreted the swelling in cellulose fibres in terms of gel theory. They considered the swelling to be caused by an osmotic pressure differential and used the Donnan theory to describe the distribution of the ions. The osmotic pressure is caused by a difference in concentration of mobile ions between the interior of the fibre and the external solution.An alternative model to explain the swelling is advanced here using double-layer theory. The repulsive pressure, nR, at the midpoint1560 1500- 1000- ? -u 500- THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE 0’ 2 3 4 5 6 j 8 PH 0- 60 50 40 ? “a 20 10 Fig. 3. Comparison of observed and theoretical swelling in pulps. The upper figure shows the experimentally observed swelling in carboxymethylated cotton (a super absorbent pulp). The lower figure shows the calculated swelling assuming the repulsive double-layer overlap model. [NaCl]/mol dm-3 as follows: 0, 2.5 x loF2; m, 2.5 x 10-I; 0, 0, 10-l; +, 1. of two flat overlapping double layers is equal to the osmotic pres~ure.~ If it is assumed that the attractive pressure, nA, between adjacent lamellae follows Hook’s law then (1 5 ) nA = 2dY where Y is Young’s modulus and 2d is the lamella separation. At equilibrium nA = nR, so that exp (ey/,/2kT)- 1 exp (ey/,/2kT) + 1 2dY = 64nikTexp(-22icd) Now yo can be estimated from eqn (14) and so for a given K and Y, d may be calculated.The effect of ionic strength and pH on d was calculated for the unbleached sulphate pulp (sample A of Part 2). Y was calculated by taking the d value of 70.1 A (obtained in Part 2) in mol dmP3 NaCl at pH 6.0. The results are shown in fig. 3.T. M. HERRINGTON AND B. R. MIDMORE 1561 The experimentally observed swelling for a super-absorbent pulp (carboxymethyl- ated cotton) is shown by way of comparison.12 The water-retention values given as grams of water per gram of fibre are converted to an interlamellar spacing by assuming that the volume of water is equal to dS and S is 200 m2 g-l.This model is at least satisfactory on a semi-quantitative level. The plots are both of the same general form, which is similar to the chargelpH isotherms. The swelling of the super-absorbent pulp drops by 50% between 0.025 and 0.25 mol dm-3 NaCl, whereas that of the unbleached sulphate pulp drops by 54% between This pore closure offers a plausible explanation for the lowering of charge found in 1 mol dm-3 NaCl for both the cotton linters and bleached sulphate pulp (fig. 2 and 3 of Part 1). As the pores close up, through suppression of the double layers with increasing electrolyte concentration, then the smallest pores become less accessible to hydroxide ions, thereby reducing the number of carboxy groups available for reaction.and 10-1 mol dm-3 NaCl. CONCLUSIONS POTENTIAL AT THE CELLULOSE/ WATER INTERFACE There are four ways of estimating the surface potential using the data obtained in this series of papers: (1) the Gouy-Chapman equation using the surface charge density So,, (2) negative-adsorption measurements using the surface area determined from negative or nitrogen adsorption, (3) potentiometric titrations and polyelectrolyte theory and (4) use of the zeta potential determined from electrokinetic data. (1) THE GOUY-CHAPMAN EQUATION In eqn (14) So, is known from the chargelpH isotherms, so if it is assumed that oh = oo then all that is needed is a value for S .Two are available from negative and nitrogen adsorption. The negative-adsorption value might be regarded as the most appropriate as it is determined in aqueous solution. It does not detect the small pores which are inaccessible to ions but accessible to nitrogen. Now o,, may only be approximated to od if there is little or no cation adsorption in the Stern layer. The concept of the Stern layer was invoked, as at a high potential (w,, = 200 mV) and high electrolyte concentration (0.1 mol dm-3 NaCl) at 25 "C the concentration of counter- ions at the surface would be greater than that of solid sodium chloride. By assuming a Stern layer or molecular condenser, the potential can drop to reasonable values. In the case of cellulose near the neutralisation point, the surface charge is approximately constant at all electrolyte concentrations.Now if the potential is high eqn (14) becomes o6 = (2ni EkT)iexp (eyd/2kT) but n, = n,exp(ewd/kT) (18) n, = $/2~kT (19) < ot/2~kT. (20) For the bleached sulphite pulp o, = -2.50 x C m-2 at pH 6.5 (fig. 3 of Part 1 and S = 120 m2 g-l) and hence n+ is 0.18 mol dm-3, which is well below the solu- bility of sodium chloride at 25 "C. Thus it seems reasonable to assume that it is not necessary to define a Stern layer and oo can be assumed to be a good approximation to os, the charge at the start of the diffuse double layer. Eqn (14) then gives a value of w,, which will be designated vo(GC). therefore1562 THE CELLULOSE/ AQUEOUS ELECTROLYTE INTERFACE Table 4. Comparison between the surface potentials calculated from negative-adsorption and surface-charge measurements for bleached sulphate pulp in different concentrations of sodium chloride, n, ni/10-2 mol dm-3 -y/j20(NA)/mV -y/i2*(GC)/mV -y/jgO(NA)/mV - y/$gO(GC)/mV 8.81 27.0 32.8 15.3 21.9 6.56 40.6 36.7 21.7 24.3 4.46 53.6 42.5 27.5 28.8 2.08 57.9 57.1 28.6 39.8 0.875 85.5 75.0 36.9 54.5 0.694 78.4 79.0 35.0 59.3 a &)(NA) = potential determined by negative adsorption assuming surface area S.y/iS)(GC) = potential determined from charge/pH isotherms assuming surface area S. (2) CALCULATION OF SURFACE POTENTIAL FROM NEGATIVE ADSORPTION From eqn (9) of Part 2 S = A VB2/n,[ 1 - exp (etys/2k T)]-l (21) so that if S and A V are known, then lys can be calculated. Values of SW for the oxidised pulp (when tps is effectively infinite) and SW from nitrogen adsorption are used to calculate AVW and then [l -exp(etys/2kT)] = A V / A P .(22) These values of tys will be designated y/ANA). (3) CALCULATION OF SURFACE POTENTIAL USING POLYELECTROLYTE THEORY From eqn ( 5 ) of Part 1 pK = pKo + 0.434AG(a)/kT (23) y/ = kT(pK-pKo)/0.434e. (24) where AG = ety. Thus The potential determined in this way will be designated tyel. (4) CALCULATION OF SURFACE POTENTIAL FROM ELECTROKINETICS It has been shown in the discussion section of this paper that the zeta potential is not a good estimate of the surface potential at least for bleached sulphate pulp. COMPARISON BETWEEN THE SURFACE POTENTIALS CALCULATED BY METHODS (I), (2) AND (3) BLEACHED SULPHATE PULP The negative-adsorption measurements were carried out over a range of chloride concentrations, but only at one pH, 6.5.The surface-charge/pH isotherms were determined over a range of pH but only four chloride concentrations (fig. 3 of Part 1). Interpolation of the isotherms enabled So, to be estimated at the same chloride concentrations as were used in the negative-adsorption experiments. To calculate both tyo(GC) and tys(NA), there is a choice of two values for S. One obtained from negative-adsorption measurements on the oxidised pulp, assuming no change in poreT. M. HERRINGTON AND B. R. MIDMORE 1563 Table 5. Comparison between the surface potentials calculated from polyelectrolyte theory, yel, and surface-charge measurements, y0(GC) (S = 190 m2 g-l), for bleached sulphate pulp in sodium chloride solutions 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 Wac11 = loF3 mol dm-3 0 0 10.1 20.5 23.7 38.5 33.7 53.3 43.2 65.4 53.3 75.5 65.1 84.1 77.0 91.5 94.7 98.0 124.3 103.8 [NaCl] = 10-l mol dm-3 0 0 0.6 2.1 2.1 4.2 3.3 6.3 4.7 8.4 6.5 10.5 8.3 12.6 11.2 14.6 17.8 16.6 33.7 18.6 [NaCl] = mol dm-3 0 0 1.5 6.7 3.0 13.2 4.5 19.6 10.1 25.7 18.9 31.4 30.8 36.9 43.2 42.0 58.6 46.7 81.7 51.3 [NaCl] = 1 mol dm-3 0 0 0.6 0.7 1.2 1.3 1.5 2.0 2.4 2.7 2.9 3.3 4.7 4.0 8.8 4.7 14.8 5.3 21.3 6.0 structure (120 m2 g-l), and the other a dry (solvent-exchanged) surface area obtained from nitrogen adsorption (190 m2 g-l).Values of y,(GC) and yB(NA) were calculated for both values of the surface area and are given in table 4.The larger surface area gives lower potentials. The agreement is excellent except for the two lowest chloride concentrations. From the pH/charge isotherms both yo(GC) and yel are calculated. They are compared at the four different sodium chloride concentrations at a range of values of a (i.e. pH) in table 5. For yo(GC) S is taken as 190 m2 g-l. The agreement is reasonable except that y,(GC) tends to be higher than yel at low values of a. This may be explained by the fact that here we do not have a situation where the charge is smeared over the whole surface but consists of discrete point charges so that the Gouy-Chapman theory is not applicable. The reasonable agreement between y,(GC), ya(NA) and yel has a number of consequences. (i) The Gouy-Chapman treatment of the double layer is applicable to the cellulose/wafer interface. The three methods of calculating y rely on the determination of five independent variables and therefore provide each other with mutually supportive evidence. method variables1564 THE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE Table 6.Comparison between the surface potentials cal- culated from negative-adsorption, y/,(NA), and surface- charge measurements, y/,(GC) ( S = 130 m2 g-l), for cotton linters in different concentrations of sodium chloride, ni nil mol dmb3 - ly,(NA)/mV - v,,(GC)/mV 20.2 9.9 1 1 . 1 4.33 11.0 21.9 3.22 14.6 24.6 0.980 17.1 38.3 0.852 17.5 39.8 (ii) There is little or no adsorption of ions in the Stern layer. As lyb(NA) and probably lyel are estimates of the potential at the start of the diffuse layer, their similarity with ly,(GC) suggests that this is situated more or less at the cellulose/water interface, i.e.the Stern layer is empty. (iii) If lyel, ly8(NA) or ly,(GC) provides a reasonable estimate for the surface potential of bleached sulphate pulp then the zeta potential is a very poor estimate of this potential. COTTON LINTERS Calculated values of ly8(NA) and ly,(GC) are given in table 6 for S = 130 m2 g-l for the range of chloride concentrations used in the negative-adsorption experiments and at pH 6.5. ly,,(GC) and lyel from the charge/pH isotherms are given in table 7 again using S = 130 m2 g-l. The agreement between the different methods of calculating ly is not so good. If the anomalously high values of tyel at mol dm-3 NaCl are ignored, then lye, is in reasonable agreement with ly,(GC).The lower values of lyB(NA) compared with ly,(GC) might be explained by the low charge density, but this is not reflected in a lower lye,, which might be expected to have similar character The discrepancy may lie in an over estimation of the surface area. If 06 is calculated from eqn (14) using the values of lye, at a = 0.8 and electrolyte concentrations of mol dm-3 and low2 mol dm-3, values of -0.0121 C m-2 and -0.0068 C m-2 are obtained, giving an average of -0.0095 C m-2. From the charge/pH isotherms (fig. 2 of Part 1) So, is -0.39 C g-l, so that this suggests an effective surface area of 40 m2 g-l. This indicates that a wet surface area of 48 m2 g-l, obtained from negative adsorption, may be a better estimate of the surface area rather than 130 m2 g-l.This value for the surface area would increase both lys(NA) and lyo(GC) and then even the zeta potentials obtained for the relatively unswollen cotton linters become a poor estimate of tyS. to lyB(NA)- THE ORIGIN OF THE ELECTRICAL CHARGE OF CELLULOSE It has been tacitly assumed throughout Parts 1, 2 and 3 that the origin of the electrical charge on cellulose is the adsorption of hydroxide ions, and that the mechanism for this adsorption is the reaction of these hydroxide ions with occasional carboxy groups in the cellulose structures. It can be asked what justification there is for this. (1) The surface potential calculated from negative adsorption compares very wellT. M. HERRINGTON AND B.R. MIDMORE 1565 Table 7. Comparison between the surface potentials calculated from polyelectrolyte theory, ye,, and surfacecharge measurements, y,(GC) (S = 130 m2 g-l), for cotton linters in sodium chloride solutions 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 [NaCl] = mol dm-3 0.0 0.0 6.0 5.0 12.4 10.0 20.6 14.8 32.5 19.6 44.3 24.1 59.0 28.6 76.8 32.8 97.6 36.8 120.0 40.7 [NaCl] = 10-l mol dme3 0.0 0.0 0.2 0.5 0.4 1 .o 0.5 1.5 0.7 2.0 0.9 2.5 . 1.1 3.0 1.2 3.5 1.4 4.0 1.6 4.5 [NaCl] = mol dm-3 0.0 0.0 1.5 1.6 3.0 3.2 4.1 4.7 6.8 6.3 10.1 7.9 14.8 9.4 20.7 11.0 28.4 12.5 37.6 14.1 [NaCl] = 1 mol dm-3 0.0 0.0 0.0 0.2 0.0 0.3 0.0 0.5 0.0 0.6 0.0 0.8 0.0 1 .o 0.0 1 . 1 0.0 1.3 0.0 1.4 with that calculated from the Gouy-Chapman equation using the charge/pH isotherms.The former method does not rely on any assumption about the origin of the charge on the surface, while the latter relies on the assumption that the adsorption of hydroxide ions is responsible for the surface charge. This is also true for the potential calculated from polyelectrolyte theory. (2) The P.Z.C. and ‘isoelectric point’ (from zeta potential) occur at similar values of the pH for both cotton linters and bleached sulphate pulp, indicating that adsorption of hydroxide ions is responsible for the diffuse double layer. (3) In the sodium adsorption experiment of Part 1 it was found that the adsorption of hydroxide ions is balanced by the adsorption of sodium ions. This would not be true if other negative ions were responsible for the surface charge of cellulose. We are indebted to S.E.R.C.and to Wiggins-Teape Ltd for joint sponsorship of this research. We also thank Dr Th. Tadros for helpful discussions. We would also like to thank the late Dr Heinz Corte, without whose industry, enthusiasm and, above all, inspiration this work would not have taken place. GLOSSARY OF SYMBOLS a area of cell Es streaming potential e charge on electron 4 streaming currentTHE CELLULOSE/AQUEOUS ELECTROLYTE INTERFACE proportionality constant cell constant length of cell concentration of co-ion i in bulk solution concentration of positive ions pressure across cell negative logarithm of ionization constant negative logarithm of ionization constant at a = 0 resistance of pad at low electrolyte concentrations resistance of pad at 0.1 mol dm-3 KCl radius of pore surface area per gram = tanh (zeva/4kT) = tanh (zey8 / 4 k T ) excluded volume Young's modulus valency of ion i, sign included degree of dissociation from charge/pH isotherm surface excess bulk conductivity surface conductivity electrical permittivity of the medium zeta potential viscosity of medium Debye reciprocal length, K = (2ni zf e2/&kT)b attractive disjoining pressure charge in density at surface charge in density at plane of shear pore factor potential at surface potential calculated from the Gouy-Chapman equation potential at the plane of shear potential calculated from negative adsorption potential at distance x from interface i H. R. Kruyt, Colloid Science (Elsevier, Amsterdam, 1952), vol. 1. D. K. Briggs, J . Phys. Chem., 1928, 32, 641. B. R. Midmore, Thesis (Reading University, 1983). C. W. C. Kaye and T. H. Laby, Tables of Physical and Chemical Constants (Longmans, London, 1973). * M. Y. Chang and A. A. Robertson, Pulp and Paper Res. Inst. (Canada), Tech. Rep. No. 485, 1965. 6 G. Gouy, J. Phys. (Paris), 1910, 9, 457. 7 D. L. Chapman, Philos. Mag., 1913, 25, 475. 8 S. G. Mason, Tappi, 1950, 33, 8. lo S. G. Mason, Tappi, 1950,33,413. 11 D. A. Goring and S. G. Mason, Can. J . Res., 1950, 6, 307. S. M. Neale and R. H. Peters, Trans. Faraday Soc., 1946, 41,478. J. Grignon and A. M. Scallan, J. Appl. Polym. Sci., 1980, 25, 2929. (PAPER 3/ 1622)

ISSN:0300-9599    

DOI:10.1039/F19848001553

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1984,80, 1567-1578 Hydrogenation of Nitric Oxide on (0 0 1) and (1 1 10) Surfaces of Ruthenium BY TETSUYA NISHIDA, CHIKASHI EGAWA, SHUICHI NAITO* AND KE~NZI TAMARU Department of Chemistry, Faculty of Science, University of Tokyo, Hongo, Bunkyo-ku, Tokyo 113, Japan Received 15th September, 1983 Hydrogenation of nitric oxide on flat Ru(0 0 1) and stepped Ru(1 1 10) surfaces has been studied by Auger electron spectroscopy, X-ray photoelectron spectroscopy, ultraviolet photo- electron spectroscopy and thermal-desorption studies as well as by investigation of reaction kinetics. The products of the reaction were N, and H,O under the investigated conditions ( 10-s-10-4 Pa of NO pressure and between 300 and 1100 K). A high reaction probability (0.5) was obtained on an oxygen-free clean surface, while the reaction did not take place on an oxygen-covered surface, which depends not only on the reaction temperature but also on the ratio of the partial pressures of NO to H,.Although the rate of formation of N, on an oxygen-free clean surface is almost proportional to the partial pressure of NO above 520 K, the rate of desorption of N(ads) from the surface determines the overall rate of the reaction; N(ads) exhibited heterogeneous behaviour and its activation energy was shown to be 32 kcal mol-l. By comparison with the activity of N(ads) during ammonia decomposition, it was found that the rate of desorption of N(ads) in the reaction between NO and H, was accelerated by the presence of a small amount of oxygen (perhaps < 1 % of a monolayer).The activity of N(ads) above 520 K was the same for the flat and stepped surfaces, although the rate of N, formation on the Ru(1 1 10) stepped surface was 1.5 times as fast as that on the Ru(0 0 1) flat surface at lower temperatures (450 K). It is generally accepted that ruthenium, like iron, is a good catalyst for ammonia synthesis from nitrogen and hydrogen. At room temperature, however, the sticking probability of N, on an Ru surface is considered to be very small, while that on Fe is reported to be ca. lo-’ above room temperature.l Therefore, few investigations of chemisorbed nitrogen on Ru have been carried out under ultrahigh-vacuum conditions. Danielson et aZ.2 studied the adsorption and desorption of ammonia on an Ru(0 0 1) surface by thermal desorption studies and LEED.Ammonia, which was adsorbed at 100 K, was desorbed molecularly at 133 and 186 K, whereas the adsorption of ammonia at temperatures between 300 and 500 K took place via an activated process and gave rise to a (2 x 2) LEED pattern through its dissociation. Klein and Shih3 stud- ied the chemisorption of NO on Ru(l0 T 0) by using FEM and thermal-desorption methods. Nitric oxide showed three and two binding states when it was adsorbed at 120 and 295 K, respectively, while it was dissociatively adsorbed above 250 K. The changes in the state of NO adsorbed on an Ru(0 0 1) surface at low temperature upon subsequent heating have been studied by multi-technique approaches, using methods such as X.P.S., u.P.s., thermal desorption studies and work-function measurements.* Two molecular states of NO below 150 K were distinguished and the dissociation of NO took place at 370 K on an NO-saturated surface.The same behaviour was also observed in EELS s t ~ d i e s . ~ The behaviour of nitrogen adsorbed on an Ru surface, however, has never been 15671568 HYDROGENATION OF NO ON RU SURFACES investigated during the reaction comprising NO hydrogenation and ammonia decomposition.6 In this study it was found that atomic nitrogen accumulates on the Ru surface during NO hydrogenation, as revealed on Pd7 and Rh8 foils, where the formation of N, and NH, proceeds through adsorbed atomic nitrogen. Therefore, the reaction mechanism of NO hydrogenation can be discussed through the behaviour of adsorbed nitrogen and may be compared with that for the decomposition of ammonia.Recently, much attention has been paid to the step sites in a catalyst surface; these are considered to be the active sites for some catalytic reactions. For instance, Blakely and Somorjaig investigated the dehydrogenation and hydrogenolysis ofcyclohexane and cyclohexene on a stepped Pt surface and concluded that the steps play an important role as the active sites for C-H and H-H bond-breaking processes. Since the dissociation of NO takes place at temperatures as low as 200 K on an Ru(1 1 10) stepped surface,l0 the effect of steps in the Ru surface in NO hydrogenation has been investigated in this study by comparing the activity on flat Ru(O0 1) and stepped Ru(1 1 10) surfaces. EXPERIMENTAL The experiments were carried out using two different stainless-steel ultrahigh-vacuum systems, one equipped with 4-grid LEED-A.e.s.optics and mass filter as described pre- viously,ll and the other an ESCALAB5 (Vacuum Generators) apparatus used in the X.P.S. and U.P.S. measurements. The sample used in the experiment was an Ru single crystal (99.999% purity) obtained from Metal Research Ltd. The (0 0 1) and (1 1 10) planes of the crystal were prepared by spark erosion and standard polishing after the normal to the desired plane had been determined by an X-ray diffraction method. The sample was 8.0 mm in diameter and 0.9 mm thick for the (0 0 1) plane and 7.7 mm in diameter and 0.5 mm thick for the (1 1 10) plane. The sample was mounted on a sample holder made of spot-welded Ta wire (0.3 mm in diameter), through which the sample was resistively heated by a d.c.source. This facilitated the use of the thermal desorption technique. The sample temperature was measured using a Pt-Pt/Rh (13 %) thermocouple spotwelded at the centre edge of the sample. Gas pressures were measured by an ion gauge and a mass filter. At the beginning the sample surface was contaminated by carbon and sulphur which were detected by A.e.s. These impurities were burned off under 2 x Pa of oxygen at 900 K and then the sample was reduced by 1 x Pa of H, at 800 K. After the repetition of this procedure for 2 h, a clean surface was obtained and no impurities were detected by A.e.s. during the reaction. Edge effects were eliminated by the sulphidation of the surface by the adsorption of 5 langmuir of H,S, followed by Ar-ion bombardment (1.3 kV and 4 pA for 30 min) on the defined surface only.Following the cleaning procedure a hexagonal LEED pattern of (1 x 1) structure was observed in the case of the (0 0 1) plane. The (1 1 10) plane of the stepped surface was identified by the splitting of the LEED spots. The direction and width of this splitting gave information about the terraces and steps on the surface. Further, the occurrence of splitting in (0 0) and (1 0) spots for various beam voltages makes it possible to determine the height of the step. From an analysis of these LEED patterns a surface structure of five terrace widths and one step height was determined, as is depicted in fig. 1 ; this means that the step density is 20% on the surface.The surface density of the (1 1 10) plane was calculated to be 1.51 x 1015 atom ern-,, which is a little smaller than that of the (0 0 1) plane (1.58 x lOI5 atom ern-,). The rates of the reaction in the steady state were determined using the following equation for a flow reactor: AP. S . r . = - BAP, a RT,An,T. NISHIDA, C. EGAWA, S. NAITO AND K. TAMARU 1569 17.6' Fig. 1. Surface atomic structure of the stepped Ru(1 1 10) surface. where ri is the specific rate of formation of product i in the unit of turnover frequency, Api is the change in partial pressure of species i during the reaction, S, is its pumping speed, R is the gas constant, q is the gas temperature, A is the sample area and no is the surface density of Ru atoms per cm2.The amount of species adsorbed on the surface during the reaction was measured by the thermal desorption technique, calibrated by A.e.s. This amount, nads, was calculated by the following equation : where S,, Tp, R, A and no are defined above, a is the number of atoms of a specific type per desorbed molecule and the integral gives the intensity of the thermal desorption peak measured by a mass filter, where nads is expressed as a coverage. Ta was used as the support in this study because its nitride and oxide are very stable and had no effect on the measurements. A preliminary study showed that nitrogen adsorbed on Ta never desorbed to the gas phase and diffused into the bulk when heated to higher temperatures (ca. 2000 K). The 15N0 isotope was employed in the experiments in order to distinguish N, produced from CO, a residual gas, during the reaction.One of the gas inlet systems was improved so that gas could be leaked through a stainless-steel capillary (0.3 mm in diameter) on top of the sample. As a result, the gas pressure near the sample was ca. one order of magnitude greater than the background pressure. By using such an improved gas inlet system the reaction-gas pressure decreased (or increased) to about one-tenth (or ten-fold) at the instant of stopping (or introducing) the gas stream. The binding energy in X-ray and ultraviolet photoelectron spectra is referred to the Fermi level of the Ru sample.1570 rr a 0 4-l E 0 WDROGENATION OF NO ON RU SURFACE? i \ I &+ I t Fig. 2. Rate of formation of N, and nitrogen coverage plotted as a function of temperature during the reaction of NO and H, on Ru(0 0 1).pNO = 2.8 x low5 Pa, pHz = 1.2 x Pa. 0, rate of formation of N,; 0, nitrogen coverage. RESULTS AND DISCUSSION STEADY-STATE REACTION RATE AND ITS SURFACE CONDITIONS Fig. 2 shows a typical result of the reduction of nitric oxide with hydrogen on an Ru(0 0 1) surface under pNO = 2.8 x Pa. The reduction took place above 500 K, and N, and H,O were formed. Neither N,O nor NH, was produced under the conditions investigated, which is characteristic behaviour over Ru catalysts.12 The rate of the reaction increased with temperature rapidly up to 600 K and then remained almost constant, at which point the reaction probability (the ratio of N, formation to NO collision frequency) attained a high value, 0.5.During the reaction only nitrogen was adsorbed on the Ru surface, and its coverage decreased with increasing temperature. Other adsorbed species were not detected by A.e.s. and thermal-desorption methods. On the other hand, the reaction did not occur at temperatures below 500K, for which adsorbed oxygen was observed on the Ru surface by A.e.s. Once oxygen covered the Ru surface, the NO+H, reaction did not occur unless adsorbed oxygen was first removed from the surface; this required heating the sample to above 1000 K. Surface species formed during the reaction were also identified by X.P.S. and U.P.S. measurements as shown in fig. 3-5. As shown by the A.e.s. technique, atomic nitrogen was only observed at 490 K under pNO = 5 x loa6 Pa and pHz = 5 x Pa, which was identified by N 1s core-level emission at 397.3 eV [fig.3(a)] and a 5.6 eV peak in He I1 valence region [fig. 4(b)]. This assignment is supported by the fact that both peaks are also observed during NH, decomposition above 500 K.l0 By lowering the temperature from 500 to 350 K [fig. 4(c) and (41 the 5.6 eV peak shifted to the higher-binding-energy side of 6.5 eV and increased in intensity, and a small new peak appeared at 9.5 eV at 350 K. This spectral change indicates that at lower temperatures the surface was mainly covered by the dissociated oxygen and a small amount of NO was adsorbed in molecular form, because, as shown in fig. 6, molecularly adsorbed NO gives an intense peak at 9.5 eV (5a+ In) together with the peak at 13.9 eV (k), Pa and pH, = 1.2 xT.NISHIDA, C. EGAWA, S. NAITO AND K. TAMARU 1571 0 1s N 1s A at (c) 500K (b ) at 350K - (a at 490K 535 530 525 400 3 95 P C " " " ' . . L binding energylev Fig. 3. N 1s and 0 1s X-ray photoelectron spectra during the reaction of NO and H, on the Ru(1 1 10) surface.pNo = 5 x Pa,pHz = 5 x lop5 Pa. (a) At 490 K; (b) temperature lowered to 350 K; (c) at 500 K (deactivated surface). 15 10 5 0 15 10 5 O = E , binding energy/eV Fig. 4. He I1 ultraviolet photoelectron spectra and difference spectra during the reaction of NO and H, on the Ru(1 1 10) surface. pNO = 5 x Pa. (a) Clean surface; (6) at 490 K; (c) temperature lowered to 420 K; (d) temperature lowered to 350 K. Pa, pHe = 5 x and the dissociated oxygen formed by heating following the adsorption of NO displays the peak at 6.5 eV.This is also confirmed by 0 1s core-level emission at 530.3 eV in the X-ray photoelectron spectra [fig. 3(6)] under the same conditions, where the shoulder on the higher-binding-energy side (53 1.8 eV) is due to molecular NO.l0 Fig. 5 gives the ultraviolet photoelectron spectra for surface species on the oxygen-dissociated surface, where the reaction did not proceed at higher temperatures. No atomic nitrogen was detected at 505 K on this surface [fig. 3(c)], although molecularly adsorbed NO1572 HYDROGENATION OF NO ON RU SURFACES 15 10 5 O = E , binding energy/eV Fig. 5. He I1 ultraviolet photoelectron spectra for a mixture of NO+H, on deactivated Ru(1 1 1O).pN0 = 5 x Pa. (a) At 380 K; (b) heated to 420 K; (c) heated to 505 K.Pa,p,, = 5 x could be observed at 420 K [fig. 5(b)]. Accordingly, the removal of oxygen formed by NO dissociation from the surface is essential for the catalytic cycle of NO reduction with high activities. EFFECT OF PARTIAL PRESSURES ON THE STEADY-STATE REACTION RATE The temperature at which dissociated oxygen starts to cover the surface as demonstrated above depends greatly on the ratio of pressures of NO to H,, as reported for Rha and Ir(1 1 O).13 When a 1 : 1 mixture of NO and H, was introduced, the dissociated oxygen began to spread over the surface above 600 K and the rate of N, formation became negligible. On the other hand, oxygen was not adsorbed near 400 K under the conditions of pNO/pH, = 0.03, where the reaction proceeded readily even at lower temperatures. The partial pressure of hydrogen which is sufficient to keep the surface free of oxygen is therefore determined by the partial pressure of NO as well as the reaction temperature.Accordingly the reaction mechanism of NO hydrogenation on an Ru surface not covered by dissociated oxygen was investigated in more detail as follows. The dependence of the rate of N, formation upon NO pressure under conditions of an excess of H, is given in fig. 7, where the rate of N, formation is nearly pro- portional to NO pressure in the temperature range 520-750 K on both the Ru(O0 1) and Ru( 1 1 10) planes. However, even when the hydrogen pressure was far higher than that defined above, both the rate of N, formation and the product distribution were independent of the partial pressure of H,, i.e.no ammonia formation was detected; this is unlike the reaction on Rh,8 Pd,' Ir13 and Pt.14 The rate of N, formation was independent of H, pressure and a high reaction probability was always attained.T. NISHIDA, C. EGAWA, S. NAITO AND K. TAMARU 1573 , ‘ a a . - ‘ * . . . ” . . I 15 10 5 O=E, binding energy /eV Fig. 6. He I1 ultraviolet photoelectron spectra for the adsorption of NO on Ru(1 1 10) and subsequent heating. (a) 2.3 langmuir exposure at 210 K; (b) 4.3 langmuir exposure at 300 K; (c) heated to 370 K; (d) heated to 450 K. Fig. 7. NO pressure dependence of the rate of formation of N, during the reaction of NO and H, on Ru(O0 1) and Ru(1 1 10) surfaces. Open symbols Ru(1 1 10); filled symbols Ru(O0 1).0, 520; 0, 562; A, 598; 0, 633 K.1574 HYDROGENATION OF NO ON R U SURFACES Fig. 8. Desorption of N(ads) on Ru(0 0 1). 0 , 5 2 0 ; 0 , 5 6 2 ; V, 578; A, 598; 0 , 6 3 3 ; c), 663 K. Table 1. Activation energy for the desorption of N(ads) on Ru(0 0 1). peak heating initial activation Tm/K P/K s-l no(&) E,/kcal mold' temperature, rate, coverage, energy, 665 12 9.2 x 31.8 700 1 1 5.1 31.5 715 10 4.2 32.1 735 9.5 2.8 31.8 DESORPTION OF N(ads) ON Ru SURFACES As the nitrogen-containing product is only molecular nitrogen, the process of desorption of adsorbed nitrogen that had accumulated during NO hydrogenation was investigated according to the following procedure. A given amount of N(ads) was prepared on the Ru surface at various temperatures in the range 520-633 K at steady state in the NO+H, reaction, i.e. pNO = 5 x lo-' Pa andpH2 = 4 x lod5 Pa.At t = 0 both the NO and H, gases were evacuated and after a given period for desorption the amount of N(ads) remaining on the surface was measured by thermal desorption. The change in amount of N(ads) on the Ru surface is plotted as a function of time in fig. 8, and it is seen that nitrogen coverage decreases linearly with the logarithm of time for nitrogen coverages above 0.0 1. Accordingly, it is seen that the process of desorption of N(ads) obeys neither first-order kinetics, i.e. ud = ke,, nor second-order desorption kinetics, i.e. vd = kek, but instead may be expressed by the Zeldovich- Roginsky equation, i.e. ud = k exp (he,); this shows the heterogeneous behaviour of N(ads) on the Ru surface.From Arrhenius plots of the rate constant, k, of N(ads) desorption the activationT. NISHIDA, C. EGAWA, S. NAITO AND& TAMARU 1575 energy for N(ads) desorption on an Ru surface at near-zero coverage can be estimated as 32.1 & 1 kcal mol-l, with the following parameters h = 1.2-1.5 x 10, k = 7.3 x lo5 exp (- 32.1 kcal mol-l/RT) (s-l) where vd = k exp (he,) and h is a heterogeneous parameter. Since the heterogeneity of N(ads) on the surface was revealed by the isothermic desorption method, the activation energy for the desorption of N(ads) was also estimated from the temperature of the thermal-desorption peak as follows. The rate of desorption of N(ads) can be described by the following equation: v(t) = -- dn(t) = v exp [hn(t)] exp (- E,/RT) dt where n is the concentration of adsorbed species. When the sample temperature was raised linearly at a rate of B K s-l the following equation is valid at the thermal- desorption peak temperature, T, : dv(t) dt - = 0.If the nitrogen coverage at T, is assumed to be in,, from the thermal-desorption curve, where no is the initial coverage, the activation energy can be derived from the following equation : The activation energy obtained from this equation for several different initial coverages and B is summarized in table 1, and it is found to be in good agreement with that obtained from isothermic desorption. Note that there exists a general trend in activation-energy change for N, desorption on Ru, Rha and Pd7 surfaces, i.e. 32, 25 and 16 kcal mol-l, respectively.The homogeneous-desorption treatment was also applied to this system, employing the following equation : where rn is the kinetic order of the desorption. After various amounts of N(ads) were accumulated on the surface, the sample temperature was lowered to 435 K and then thermal-desorption spectra were measured at a low heating rate (ca. 5 K s-l). Plots of reciprocal temperature against log (desorption rate) at a certain coverage or of log(nitrogen coverage) against log(N, desorption rate) did not have a linear relationship. It is therefore confirmed that the homogeneous treatment cannot be employed for the desorption of N(ads) in this case. By using these parameters it is possible to calculate the rate of desorption of N(ads) at any coverage and temperature, this is plotted as the filled points in fig.9. The open points in fig. 9 give the rate of N, formation and nitrogen coverage at a steady state of NO hydrogenation, indicating that the rate of the NO+H, reaction in the steady state is also proportional to the exponential of nitrogen coverage, i.e. rNg cc exp(&), and equal to the rate of desorption of N(ads). Accordingly it is concluded that the process of desorption of N(ads) from the surface is the rate-limiting step in NO hydrogenation on the Ru surface.1576 HYDROGENATION OF NO ON RU SURFACES 0 0.5 1 nitrogen coverage (X 10-1) Fig. 9. Rate of N, formation plotted against nitrogen coverage during the reaction of NO and H, on Ru(0 0 1). 0, 520; 0, 562; A, 598; 0, 633 K. St 640 750 K Fig.10. Thermal desorption spectra of N, on Ru(0 0 1): (a) during NH, decomposition (pNHa = 1 x Pa); (b) during NH, decomposition after 0.32 langmuir exposure of 0,; (c) dunng NH, decomposition after 0.97 langmuir exposure of 0,; ( d ) during the reaction of NO and H,. Thermal-desorption spectra from 520 K.T. NISHIDA, C. EGAWA, S. NAITO AND K. TAMARU 1577 COMPARISON OF THE BEHAVIOUR OF N(ads) ON THE RU SURFACE BETWEEN NO HYDROGENATION AND NH, DECOMPOSITION The behaviour of N(ads) formed during NO hydrogenation was compared with that formed during NH, decomposition. For the decomposition of ammonia6 relationships similar to those of fig. 9 could be obtained between the rate of N, formation and nitrogen coverage in the steady state under various ammonia pressures at each temperature. However, the rate of formation of N, in the hydrogenation of NO is about ten times faster than that of ammonia decomposition at the same nitrogen coverage between 520 and 630K, which corresponds well to the change of peak temperatures in the thermal desorption of N,, i.e.the peak temperature of the desorption of N(ads) during NO hydrogenation appeared at ca. 650 K, while that formed during ammonia decomposition was observed to be ca. 750 K. It is therefore considered that the desorption of N(ads) in NO hydrogenation is affected by a small amount of oxygen on the surface. To make this point clearer, the effect of a small amount of adsorbed oxygen on ammonia decomposition was examined. When a small exposure of 0, (below 1 langmuir) was introduced, temperature peaks corresponding to N, desorption appeared at both 640 and 750 K, in accordance with the increase in the rate of formation of N,, as shown in fig.10. Consequently it may be inferred that the desorption of N(ads) during NO hydrogenation is accelerated by the presence of a small amount of adsorbed oxygen which cannot be detected by electron spectroscopic methods and is therefore estimated to be below 1 % of a monolayer. DIFFERENCES BETWEEN THE FLAT Ru(0 0 1) AND STEPPED Ru( 1 1 10) SURFACES IN NO HYDROGENATION As reported by Umbach et aZ.,4 the NO molecule is dissociatively adsorbed on the Ru(0 0 1) surface at temperatures above 370 K; this is observed on both terrace and stepped surface.l0 It is therefore likely that the dissociation of NO readily takes place on both surfaces under reaction conditions.In addition, taking the high reaction probability (0.5) into consideration, a difference in the rate of N, formation between the two surfaces would not be expected. In fact, the rate of N, formation above 520 K in the NO + H, reaction was the same for both the flat and stepped surfaces. However, when the reaction temperture was lowered to 450 K under conditions of an excess of H,, the rate of N, formation on the stepped Ru( 1 1 10) surface was observed to be 1.5 times as fast as that on the Ru(0 0 1) surface. At similar temperatures (below 500 K) the rate of ammonia decomposition on the stepped Ru(1 1 10) surface is initially one order of magnitude faster than that on the flat Ru(0 0 1) surface, where a second peak, N*(ads), was observed at 570 K in thermal-desorption spectra.This was considered to be an activated nitrogen atom formed selectively on the stepped sites and has been confirmed to have an important role in N, formation.6 We therefore consider that the difference in the rate of N, formation at 450 K during the NO+H, reaction is related to the formation of N*(ads) on the stepped surface. CONCLUSIONS NO reduction with H, on flat Ru(0 0 1) and stepped Ru(1 1 10) surfaces has been investigated in this study. The products of the reaction were N, and H,O; neither NH, nor N,O was formed under the conditions investigated. The rate of N, formation exhibited a high reaction probability (0.5) almost irrespective of the reaction temperature on an oxygen-free clean surface, while the reaction did not take place when oxygen atoms covered the surface.The surface conditions during the reaction 52 FAR 11578 HYDROGENATION OF NO ON RU SURFACES described above are dependent on the reaction temperature as well as the ratio of partial pressures of NO and H,. The rate of N, formation on an oxygen-free surface is proportional to NO pressure and independent of H, pressure. The overall rate of NO reduction is determined by that of N(ads) desorption on the surface, which is heterogeneous in nature, and the activation energy is found to be 32 kcal mol-I. By a comparison with the decomposition of ammonia, where surface nitrogen also exhibits heterogeneous behaviour, it is found that the desorption of N(ads) is accelerated by the presence of a small amount of oxygen on the surface (perhaps below 1 % of a monolayer).The rate of N, formation in NO hydrogenation is an order of magnitude faster than that during ammonia decomposition at the same nitrogen coverage. Although the rate of the reaction of NO and H, above 520 K was the same for both the flat and stepped surfaces, the rate of N, formation on the Ru(1 1 10) surface was observed to be 1.5 times as fast as that on the Ru(0 0 1) surface at temperatures as low as 450 K. This work was supported by a Grant-in-Aid for Scientific Research from the Japanese Ministry of Education, Science and Culture. F. Bozso, G. Ertl, M. Grunze and M. Weiss, J. Catal., 1977, 49, 18. L. R. Danielson, M. J. Dresser, E. E. Donaldson and J. T. Dickinson, Surf. Sci., 1978, 71, 599. R. Klein and A. Shih, Surf. Sci., 1977, 69, 403. E. Umbach, S. Kulkami, P. Feulner and D. Menzel, Surf. Sci., 1979,88,65; P. Feulner, S. Kulkami, E. Umbach and D. Menzel, Surf. Sci., 1980,99,489. G. E. Thomas and W. H. Weinberg, Phys. Rev. Lett., 1978, 41, 1181; P. A. Thiel, W. H. Weinberg and J. T. Yates Jr, Chem. Phys. Lett., 1979, 67, 403. C. Egawa, T. Nishida, S. Naito and K. Tamaru, to be published. ' A. Obuchi, S. Naito, T. Onishi and K. Tamaru, Surf. Sci., 1982, 122, 235. * A. Obuchi, S. Naito, T. Onishi and K. Tamaru, Sur- Sci., 1983, 130, 29. D. W. Blakely and G. A. Somorjai, J. Catal., 1976, 42, 181. lo C. Egawa, S. Naito and K. Tamaru, Surf. Sci., submitted for publication. l1 C. Egawa, S. Naito and K. Tamaru, Surf. Sci., 1983, 125, 605. l2 K. C. Taylor and R. L. Klimisch, J. Catal., 1973, 30, 478. l3 D. E. Ibbotson, T. S. Wittrig and W. H. Weinberg, Surf. Sci., 1981, 111, 149. l4 G. Pirug and H. P. Bonzel, J. Catal., 1977, 50, 64. (PAPER 3/ 1625)

ISSN:0300-9599    

DOI:10.1039/F19848001567

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1984, 80, 1579-1594 Metal Dispersions on Zirconium Phosphates Part 2.-Hydrogen Reduction of Silver(1)-exchanged a-Zirconium Phosphate BY SOOFIN CHENG-~ AND ABRAHAM CLEARFIELD* Department of Chemistry, Texas A&M University, College Station, Texas 77843, U.S.A. Received 2 1st September, 1983 The kinetics of the reduction by hydrogen of Ag+ ions in a-zirconium phosphate and the formation of silver particles have been studied. Reduction proceeded under relatively mild conditions. The time dependence of the percentage of Ag+ ions reduced in this reaction is always expressed by an S-shaped curve with a maximum reaction rate at ca. 15 % conversion. An initial induction period was observed and is suggested to result from the extremely slow rate of initiating silver nuclei.The remainder of the reaction rate was found to conform to an Elovich-typed equation. The initial rate was found to be directly proportional to the Ag+ loading and to the surface area of the support but proportional to the half-power of the hydrogen pressure. It is also suggested that the reduction reaction occurs on the surface with the reaction itself determining the rate after the induction period. Supported metals are widely used as catalysts. Metal is usually introduced to the support as cations, either from aqueous solutions or from suspensions, by one of several processes such as impregnation, ion exchange, deposition or coprecipitation, followed by drying and hydrogen reduction. Such catalysts, which contain very small metal crystallites, have certain definite advantages over the bulk metal, which can be in the form of films, wires or powders, because the high dispersion of the metal leads to a high surface area and to increased resistance to sintering.However, the dispersion of metal particles and the activity of the catalyst strongly depend on the conditions under which the catalysts have been prepared and on the particular support ad0pted.l Many efforts have been made to develop finely dispersed metal particles by reduction of transition-metal ions in zeolites.2+ The possibility of these catalysts being used as highly active and/or bifunctional catalysts is the incentive for such work. Kinetic and physicochemical studies have been reported on the formation of metal particles in Nevertheless, incomplete characterization is an inextricable problem owing to the diversity of sites existing in zeolites and to the limitation of analytical methods employed in the studies.There is theoretical and practical interest in studying the redox behaviour of transition-metal ions in zirconium phosphates. They form synthetic hydrogen- containing analogues of clays and possess cages akin to those of zeolites. As a comparison, zirconium phosphates provide a simpler system for studying metal-particle formation. a-Zirconium phosphate, Zr(HPO,), H,O, is a crystalline ion-exchanger. It has a layered structure with an interlayer distance of 7.56 A. Ion exchange occurs by replacement of the orthophosphate hydrogen by cations which then occupy positions between the layers.' The flexible interlayer distance in zirconium phosphate makes it possible to detect the formation of small metal clusters in between the layers by X-ray powder diffraction methods.In contrast, structural differences in zeolites t The work reported in this paper formed a part of the Ph.D. Thesis of S. Cheng. 1579 52-21580 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES are not easily detected using X-ray diffraction techniques because of their rigid structures. From a practical point of view, studying the redox behaviour of cations in zirconium phosphate (often referred to as ZrP) will lead to a better understanding of the durability and applicability of this material in various catalytic reactions. In previous work8 the reduction of CulI by hydrogen in copper-exchanged a-zirconium phosphate, ZrCu(PO,),, was found to proceed in two stages.Below ca. 150 Torr* the product was ZrCuH(PO,),, as CuI1 was reduced to Cul. Then, further reaction with hydrogen at pressures above 150 Torr yielded copper metal and A-zirconium The rates of both reactions were reported to conform to an Elovich-type equation. In the present work, silver reduction is investigated. Since silver reduction is only a one-stage process, the study of this simpler system is expected to provide more information on the metal-formation step. The influence of hydrogen pressure, temperature, AgI loadings and the surface area of a-zirconium phosphate on the process is examined. EXPERIMENTAL SAMPLE PREPARATION A series of zirconium phosphate samples with different crystallinities was prepared by refluxing the zirconium phosphate gel in H3P04 solution for periods of time as described by Clearfield et aZ.l0 These samples are labelled ZrP (4.5: 48), (6:48), (9: 48), (12: 50) and (12: 336), where the first number represents the molar concentration of H3P04 acid and the second number the duration of reflux in hours.The crystallinity increases with the reflux time and concentration of acid. Silver-exchanged zirconium phosphate samples were prepared as described elsewhere. l1 Samples for reduction were obtained by dehydrating the air-dried, silver-exchanged zirconium phosphates at CQ. 100 "C in a vacuum overnight and were kept in a vacuum desiccator. CHEMICALS Reagent-grade silver acetate (Fisher Scientific Co.) was used as the source of Ag+.Ultrahigh-purity (99.999 %) hydrogen was obtained from Matheson Gas Products. Zirconium phosphate gel was obtained from Magnesium Elektron Inc. Reagent grade phosphoric acid (Fisher Scientific Co.) was used without further purification. Distilled deionized water was used throughout. APPARATUS As mentioned in a previous paper,* the apparatus in which the reductions were carried out consists of a U-tube Pyrex reactor connected to a recirculatory loop which forms part of a vacuum system. The temperature of the reaction zone inside the reactor was read via a thermocouple held inside the sample and connected to a Doric digital trendicator model no. 400A. A side-arm of Pyrex tubing was attached for collecting samples in situ, thus avoiding any exposure to the atmosphere.In order to minimize the dead space, most of the assembly (loop) was made of capillary tubing. The total volume of the system was ca. 100 cm3 and that of the reactor was ca. 27 cm3. The reactor was heated by means of a well-insulated tubular furnace; the heating rate and the temperature were regulated using a Weathermeasure temperature controller. A Pyrex tube for blowing cold air was inserted inside the heating zone of the furnace to conduct away heat evolved from the reaction. The temperature control of the reaction zone in the oven with this air circulation was ca. 1 "C. The pressure changes occurring during the process of reaction were monitored by a Validyne differential-type digital manometer connected to a linear single-pen recorder.The precision of the manometer was f 0.1 Torr. REDUCTION PROCEDURES A weighed amount of dehydrated sample (ca. 0.2 g) was placed inside the reactor and spread over supported glass wool to achieve maximum surface exposure. The reactor was mounted onto the vacuum system and heated at the reduction temperature for > 5 h under vacuum to * 1 Torr = 101325/760 Pa.S. CHENG AND A. CLEARFIELD 1581 remove any absorbed moisture which may have been picked up during the transfer process. The reduction was started by bleeding-in a known amount of hydrogen which was kept in the outer chamber at a known pressure. The hydrogen uptake was followed by recording the decrease in pressure with time. The end of the reaction was determined when no further change in pressure could be detected for ca.30 min. Part of the reduced sample was examined by X-ray powder diffraction methods immediately after the reaction. The rest of the sample was transferred to the side-arm tube and sealed off under vacuum pressure for further analysis. The amount of Agl reduced was then calculated from the amount of hydrogen consumed. INSTRUMENTAL The powder diffraction patterns were obtained using a Seifert-Scintag Pad I1 X-ray diffractometer with Cu Ka radiation. The average silver particle size was determined from line-broadening measurements on the (1 1 1) reflection of silver metal with the aid of Scherrer's equation.12 These values were not corrected for errors resulting from instrument or strain. The data were used only as a basis of comparison.The silver-particle dispersion was determined by means of scanning electron microscope with a Jeol model JSM-35 oscilloscope. RESULTS The reduction of anhydrous samples of Ag+-substituted a-zirconium phosphate, i.e. ZrAg,H,-,(PO,), where 0 < x d 2, was found to be an exothermic reaction. As hydrogen was consumed, the colour of the samples changed from white to sepia or olive, indicating the reduction of Ag+ ions to Ag atom aggregates. X-ray powder diffractions patterns of the reduced samples contained reflections at 28 = 38.2 and 44.4", which correspond to the Ag (1 1 1) and (200) planes, respectively. The reduction rate was found to be a function of the initial hydrogen pressure, the temperature of the reaction zone, the Ag+ ion loading, and the surface area of the zirconium phosphates.Kinetic studies were mainly carried out on highly crystalline zirconium phosphate (12:336) in order to avoid difficulties in temperature control caused by the abrupt release of enormous amounts of heat during the reduction of the high-surface-area samples. FORMATION OF METALLIC SILVER PARTICLES X-ray powder diffraction patterns of the completely reduced samples (12: 336) showed reflections due to silver metal and c-ZrP [composition Zr(HP0,),]7i rather than the A-ZrP (same composition) observed on reduction of CuII-ZrP. In addition, the high-surface-area samples gave different results. Instead of the &ZrP phase, reflections characteristic of an Ag-I1 phasell with composition ZrAg,~,,H,~,,(PO,), were observed along with the reflections due to silver metal.However, if the samples were only partially reduced, diffraction patterns characteristic of metallic silver and the anhydrous phases of Ag+-containing zirconium phosphate phases, which could be Ag-I [ZrAg,,,,H,.,,(PO,),], Ag-I1 or Zr(AgPO,),,* depending on the extent of reduction, were obtained. The Ag-I phase had an assumed composition of The sizes of the silver particles formed by the reduction were zvaluated by X-ray line-broadening measurements based on the width at half height of the Ag (1 1 1) peak. Table 1 shows the average silver particle size of the samples obtained by reducing Zr(AgPO,), (12: 336) at various temperatures and at different initial hydrogen pressures. The particle size does not seem to be a function of these two reaction conditions.It will be shown later that all the silver comes to the surface. Apparently ZrAg0.22Hl.78(Po4)2.'1 * The X-ray pattern of the fully exchanged phase exists over a range of compositions, ZrAg, H2JPO&, with x = 1.9-2."1582 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES 400 Table 1. Average silver particle size of the samples obtained by reducing Zr(AgPO,), (12: 336) at different temperatures (pH, = 580 Torr) or hydrogen pressures (T = 150 "C) (A) - - av. particle av. particle T/OC size/A p/Torr size/A 100 0 62 392 124 525 100 393 191 424 145 426 354 54 1 190 389 583 446 - - - - I I 1 1 1 I l l , , surface area of ZrP support/mZ g-' Fig. 1. Average Ag particle size of the samples obtained by (A) reducing Ag+-exchanged a-ZrP (1 2 : 336) with various Ag+ loadings ( T x 100 "C, p H 1 x 585 Torr) and (B) reducing Zr(AgPO,), as a function of surface area of the supports (T x 100 "C pH2 x 280 Torr).J.Chem. SOC., Faraday Trans. 1, Vol. 80, part 6 S. CHENG AND A. CLEARFIELD Plate 1 I: %- - n a m m (u .. 3 w r( I M n en (Facing p . 1582)J. Chem. SOC., Faraday Trans. I , Vol. 80, part 6 Y Plate 2 S . CHENG AND A. CLEARFIELDS. CHENG AND A. CLEARFIELD I583 the very small surface area of the support [2.2 m2 g-l for (12:336)] relative to the amount of silver deposited leads to the formation of relatively large particles with a wide distribution range. As the loading of Ag+ in the exchanger was reduced there was an almost parallel reduction in the average particle size as a function of loading [fig.1 (A)]. The effect of increasing the surface area of the support is shown in fig 1 (B). Again it is seen that the particle size decreases as the space available on the surface increases. The silver-particle distribution was examined by scanning electron microscopy. Plates 1 (A) and (B) show the silver particles on ZrP( 12 : 336) samples which contained 6.64 and 1.18 m equiv g-l of Ag+, respectively, before reduction. Reduction was carried to completion. The ZrP crystals have a shape resembling hexagonal platelets13 but the composite with silver is more ellipsoidal. The highly loaded sample gives a wide range of silver particle-size distributions [plate 1 (A)]; the particle diameter varies from ca. 0.6pm to a few hundred Angstroms. However, the size distribution on the low silver-loaded sample is relatively uniform, as shown in plate 1 (B).In order to understand further the formation and distribution of silver particles on the surface of a-ZrP, single crystals of a-ZrP with a diameter of ca. 0.1 cm were exchanged with silver acetate, reduced by H, and examined by the scanning electron microscope. For comparison, a single crystal of a-ZrP without Ag content was exampled by the same technique. Without a silver coating, both crystal faces are clean and smooth. For the single crystal containing Ag, the micrographs [plates 2(A) and (B)] show two kinds of silver-particle distribution, one on the side and the other on the base of the platelet. Plate 2(A) shows one of the cracks on the side of the platelet and the very thick layers of large silver particles which accumulate on its surface.However, a very uniform distribution of fine silver particles can be observed on the basal surface of the hexagonal platelet, as shown in plate 2(B). After the surface of the crystals had been examined, they were sliced through the side so that the interior layers could be revealed. Although no differences were discernible between the surface and interior layer of the pure a-ZrP crystal, as expected, marked differences were noticeable for the silver-coated crystal. Large silver particles were found to aggregate on the side surface of the platelet. The particle sizes become smaller and smaller as the particles approach the edge of the crystal. Silver particles also accumulate along cracks existing in the interior layer.This reveals that the silver particles can be located anywhere, providing space is available. KINETICS OF REDUCTION The reduction of silver-exchanged a-ZrP can be expressed as Zr(AgPO,), + H, -+ Zr(HPO,), + 2Ag. The kinetic curves for the reduction of dehydrated Zr(AgPO,), (1 2 : 336) as a function of temperature are shown in fig. 2 (A). These curves are typically S-shaped. The value of the percentage conversion, a, increases rapidly after the initial induction period and reaches a maximum increasing rate at a point between 5 and 20% conversion, depending upon temperature. After this the rate of increase of a slows down gradually. This variation is shown in fig. 2(B), where da/dt is plotted a a function of time. This type of kinetic curve has also been observed in other reduction s y s t e m ~ .~ ~ - ~ ~ However, attempts to represent the curves by a simple algebraic equation have not, on the whole, been successful. The best fit was reported with the Elovich equation, as it applies to most of the data in the middle part of the curves.8v l4 In the present study the Elovich equation was found to fit the curves in the range a = 0.1-0.8, where a is the extent of conversion. This portion of the reaction will be treated separately from the initial induction period or accelerating rate period (ARP).1584 100 80 h 60 G .- E 2 40 > 20 04 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES tlmin Fig. 2. (A) Kinetic curves for the reduction of (12:336) Zr(AgPO,),; (B) time dependence of the reduction rate at various temperatures.A, 86; a, 104; 0, 121 and 0, 132 "C. INDUCTION PERIOD The reduction reactions are characterized by an initial induction period which extends to a x 0.1 and is not fitted by the Elovich equation. The range of the induction period was found to depend upon the reduction temperature, hydrogen pressure, the Ag+ loading and the surface area of the support. Similar induction periods haveS. CHENG AND A. CLEARFIELD 1585 commonly been observed in solid reaction ~ y s t e m s . l ~ - ~ ~ Garner et aZ.18 suggested that the induction periods were mainly due to an abnormally slow rate of nuclei growth. RATE OF REDUCTION REACTION The Elovich equation, which fits the middle part of the kinetic curves, was originally applied to the rate of adsorption of gases on solid surface.21q22 The rate decreases exponentially with an increase in the amount (or fraction) a of gas adsorbed: daldt = Re-aa. (2) The exponent a is constant for a given sample and for fixed experimental conditions.The other constant, R, represents the initial rate when a = 0. The integrated form of eqn (2) can be written as with a = (2.3/a) log (t+ to) - (2.3/a) log to to = l/(aR). (3) (4) With a correct choice of to, the plot of a as a function of log(t+t,) should give a straight line with a slope of 2.3a. Then, the initial rate, R, is obtained from eqn (4). The disposable parameter, to, is found by trial and error; if to is too small the curve of a against log (t+ to) is convex, and if to is too large the curve is concave.23 The resultant data for the temperature-dependent kinetic curves (shown in fig.3) are collected in table 2. A general expression for any rate law is r = -dC,/dt = kC$ Cb,... CE lo3 KIT Fig. 3. Arrhenius plot of In R against 1/T for the reduction of (12: 336) Zr(AgPO,),.1586 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES Table 2. Values of zo, a, R and R' for the reduction of Zr(AgPO,), (12: 336) at various temperatures; pH2 M 580 Torr Elmo1 (Ag+) g-* T/OC 2, a R (% min-I) (catalyst) min-' 85.8 7.0 0.0150 9.50 3.82 x 10-4 104.0 2.0 0.0193 25.90 1.04 x 10-3 121.0 1.0 0.0203 49.30 1.98 x 10-3 132.0 0.5 0.0165 121.20 4.88 x 10-3 tlmin t/min Fig. 4. (for legend see opposite).S. CHENG AND A. CLEARFIELD 1587 t/min Fig. 4. Kinetic curves for (A) the reduction of Ag+-exchanged (1 2 : 336) a-ZrP with various Ag+ loadings [ T x lOO"C, pH, x 585 Torr; a, 0.324; A, 1.18; 0, 3.09 and 0, 6.64 mequiv g-l (a-ZrP)], (B) the reduction of (12: 336) Zr(AgPO,), at various hydrogen pressures [T x 100 "C; pH2 = 0, 280; 0, 376; A, 478 and 0, 587 Torr] and (C) the reduction of Zr(AgPO,), with various surface areas of the support [T x 100 "C, pH2 E 280 Torr; 0, ZrP (9: 48); A, ZrP (12:50) and 0, ZrP (12:336)].where C, is the concentration of the reactant A and CB,. . ., CN are the concentrations of either the reactants or products. For the reduction of Ag+-exchanged ZrP the rate law can be written in the following form: R' = - d[Ag+]/dt = k[Ag+]'pE2 Sn.. . (6) where p is the partial pressure of hydrogen, S is the surface area of the support and R' is the initial rate obtained from the product of R values from the Elovich plots and the initial concentrations of Ag+ in ZrP.According to the Arrhenius equation, the rate constant, k , is a function of temperature in the form k = Aexp(-E,/RT) (7) where A is the frequency factor, E, is the activation energy of the reaction and R is the ideal gas constant. Therefore the initial rate of the reduction is temperature-dependent : R' = Aexp(-Ea/RT)[Ag+l1pg2S n.... (8) The temperature effect was studied by keeping all other variables constant. Thus eqn (9) (8) reduces to with the constant B = A[Ag+Izpg2 Sn. A plot of In R' against 1/T should be linear, and the slope gives the value of -E,/R (fig. 3). The activation energy obtained has a value of 15.2kO.2 kcal mol-l.? The values of I, rn and n in eqn (6) were then determined by examining the kinetic curves obtained by varying the concentration of R' = Bexp (- E,/RT) t 1 cal = 4.184 J.1588 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES 7 9- .z 7 - E 8- Table 3.Values of to, a, R and R for the reduction of Ag+-exchanged ZrP at various Agf loadings, initial hydrogen pressures and surface areas of the support expt. no. Ag+/ T pH2 mequiv g-l S ZrP /"C /Tom (a-ZrP) /m2 g-l to (12:336) 101 280 6.64 2.2 5.0 (12:336) 102 376 6.64 2.2 4.0 (12:336) 103 478 6.64 2.2 3.8 (12:336) 99 587 6.64 2.2 3.0 (12:336) 101 593 3.09 2.2 17.0 (12:336) 99 586 1.18 2.2 50.0 (12:336) 102 584 0.034 2.2 43.0 (12:50) 100 282 6.64 5.0 2.0 (9:48) 101 281 6.64 7.8 2.0 U 0.0224 0.020 6 0.020 5 0.023 7 0.008 04 0.00420 0.005 16 0.0148 0.0105 R (% min-l) 8.93 12.10 12.80 14.10 7.32 4.76 4.51 33.80 47.60 R'lmol (Ag+) g-l (catalyst) min-l 3.59 x 10-4 4.87 x 10-4 5.15 x 10-4 5.67 x 10-4 5.26 x 10-5 1.53 x 10-5 1.36 x 10-3 1.92 x 10-3 1.78 x iO-* 11 I I I I I I I I 1 2 3 4 5 6 7 8 9 ' p",/lo* tom 3 [ Ag+l /mequiv g-' (a-ZrP) S/mZ g-' Fig.5. Logarithmic plot for variation of initial reduction rate with (a) the initial hydrogen pressure, (b) Ag+ loading and (c) the surface area for the support.S. CHENG AND A. CLEARFIELD 1589 one component and keeping those of the other components constant. Fig. 4(A), (B) and (C) show the kinetic curves as a function of Ag+ loading, initial hydrogen pressure and the surface area of the support, respectively. These curves have been fitted to Elovich plots and the resultant values of to, a, R and R' are listed in table 3.In experiments no. 1-4 the initial hydrogen pressure varied from 280 to 587 Torr. The value of rn in eqn (6) is equal to the slope of the straight line when log R' is plotted against logpH*, as shown in fig. 5(a). The experimentally determined value is 0.59. In experiments no. 4-7, the loading of Ag+ ions has been varied from 6.64 to 0.324 mequiv (Ag+) g-l (a-ZrP). Fig. 5 (b) shows that by plotting log R' against log [Ag+] a straight line results with a slope of 1.2, which is the value of I in eqn (6). The effect of the surface area of the support is obtained from experiments no. 1, 8 and 9. The surface area has been varied from 2.2 to 7.8 m2 g-l (a-ZrP) and initial hydrogen pressure has been kept at 280 Torr.The plot of log R' against log S is shown in fig. 5(c). The slope has a value 1.4, which is the value of n in eqn (6). Eqn (6) can thus be written in the form R' = -d[Ag+]/dt = k'[Ag+]1*2&fBS1.4 where k' is the rate constant independent of the Ag+ loading, the hydrogen pressure and the surface area of the support. DISCUSSION The reduction of Ag+ ions in a solid system is generally found to proceed easily under mild conditions. Several studies on the reduction of silver(1) oxide with hydrogen have been r e p ~ r t e d . ~ ~ - ~ ' The reduction reaction has been found to proceed at a measurable rate at temperatures above 40 "C with an activation energy of 8.0-1 5.0 kcal mol-l. More recently, several articles dealing with the reduction- reoxidation behaviour of silver(1) ions in zeolite systems have been p ~ b l i s h e d .~ ~ ~ ~ 28-30 None of the kinetic curves reported in those studies showed the initial induction period characteristic of the present system. In the research performed by Iwamoto et aL5 the activation energy was determined by plotting the logarithm of the initial rates of reduction, calculated from the amount of hydrogen .consumption for 10 s, against the reciprocal absolute temperatures between 273 and 741 K. The value obtained was 23.4f1.6 kJmol-l or 5.6k0.4 kcal mol-l, which is much smaller than that found in the present study (15.2 cal mol-l). This large difference is attributed not only to the different methods used in determining the initial reduction rates but also to the possibility that the reaction mechanisms and rate-determining steps are different in these two systems.On the other hand, the reduction of the Ag+ ions in a-zirconium phosphate was found to proceed much more rapidly than in X-type zeolites. For example, Iwamoto et aL5 reported that the zeolite-X sample consumed an amount of hydrogen corresponding to the complete reduction of Ag+ ions within 10 min at 450 "C. In contrast, complete reduction of Zr(AgPO,), (12: 336) could be achieved in a few seconds at much lower temperatures. Beyer et aL6 found a marked temperature dependence for the reduction of Ag+ in zeolites. At temperatures lower than 430 K, they could only achieve a maximum reduction of ca. 60% of the silver ions. They reported that two different mechanisms were operative at high and low temperatures.Above 430 K the reaction is first-order in Ag+ concentration and independent of the hydrogen pressure and has an activation energy of 97.6 kJ mol-1 (24.2 kcal mol-l). The migration of cations from SI sites of Y zeolites is thought to control the rate. At low temperatures the reaction is dependent1590 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES on hydrogen pressure with an activation energy of 40 kJ mol-l (9.5 kcal mol-l). The reaction was suggested to be a catalysed reaction with the regeneration of the surface-active sites as the rate-limiting step. Some of the impurities in the Y zeolites were thought to be active sites, although this could not be proved. The reduction of silver(1) in zirconium phosphates was found to proceed at relatively low temperatures, even at room temperature ! Moreover, the reaction rate was strongly dependent on the surface area of the supports.For a high-surface-area sample, such as Zr(AgPO,), (4.5:48) with a surface area of ca. 30 m2 g-l, complete reduction was achieved in 2 h at 21 "C with B hydrogen pressure of 440 Torr. However, for a low-surface-area sample, Zr(AgPO,), (1 2 : 336) (surface area ca. 2.2 m2 g-l), reduction of 30% of the Ag+ ions was achieved after 7 h at 43 "C andpH2 = 580 Torr. The kinetics of the reoxidation of the silver particles will be reported in a subsequent paper. However, our preliminary studies show high-surface-area samples reoxidize much faster than the low-surface-area samples. The metallic silver supported on ZrP is converted to Ag+ ions which diffuse back into the lattice of the support even at room temperature.This reoxidation phenomenon probably accounts for the difference observed between the X-ray powder diffraction patterns of the completely reduced samples of high- and low-surface area a-ZrP. Although the patterns were taken immediately after complete reduction, air oxidation of silver must occur and perhaps is accelerated by the heat produced from the X-ray beam. As a result the high- surface-area samples, where reoxidation of silver is rapid, give patterns of Ag+ partially exchanged phases instead of the c-ZrP phase. SILVER-PARTICLE SIZE AND DISTRIBUTION Silver-particle size can be changed by varying the amount of Ag+ loading or the surface area of the supports. The dispersion of the silver particles was not uniform; neither was the particle-size distribution.This can cause considerable deviation when the average particle size is determined by X-ray diffraction line-broadening measurements. Thus the data should be used only as a basis of comparison and as information required to understand the reaction mechanism, rather than as absolute values. Several factors can explain the discontinuity observed in the linear relationship between particle size and Ag+ loading when the loading is very small [fig. 1 (A)] and also the discontinuity between particle size and surface area of the support [fig. 1 (B)]. The electron micrographs show that the particle-size distribution is relatively more uniform when the Ag+ loading is low (plate 1) or the surface area of the support is higher.Different size distribution curves can lead to different average sizes. Further- more, any reoxidation which occurs on the high-surface-area samples tends to reduce the silver-particle size REACTION MECHANISM Gas-solid reactions involve a series of complicated steps which can be summarized as follows: (1) gas diffusion to a surface, (2) adsorption of gas molecules on a surface which can be a solid reactant or solid product, (3) movement of adsorbates on the surface of the solid phase, (4) adsorption on active sites, (5) reaction itself and (6) diffusion of products from the active sites. The unraveling of these complex reaction mechanisms presupposes an understanding of the various physical stages reached during the course of the total reaction.It is also clear that the relative contribution of the different steps to the overall reaction may change during the course of the reaction. It is generally accepted that hydrogen gas diffuses into the internal porous cavities of zeolites and that the reduction process proceeds both on the surface and in the internal pores.30331 However, the degree of hydrogen-gas diffusion into theS. CHENG AND A. CLEARFIELD 1591 interior of zirconium phosphate crystallites is considered to be negligible. If hydrogen gas did diffuse into the interior so that reduction proceeded inside the cavities, then nucleation of silver atoms in the cavities should result in the formation of silver clusters uniformly distributed in the interlamellar regions.As a result, the interlayer distance should become larger. However, this is not what is actually observed, i.e. the observed c-ZrP phase, Zr(HPO,),, has an interlayer spacing of only 7.41 A,9 and the silver collects on the surface of the crystals and along cracks and fissures. Therefore hydrogen diffusion into the interior of ZrP crystallites is thought to be minimal. The reason may be that the windows of the cavities of a-ZrP are too small for hydrogen-gas diffusion. The maximum free diameter of the windows in the a-ZrP crystallites is 2.64 A. The interlayer distance for Zr(AgPO,), is 7.76 A compared with 7.56 A for a-ZrP. Thus the windows in the former phase are only slightly larger. However, there are two Ag+ (radius 1.26 A)32 in each cavity of dimensions 5.3 x 7 A.Thus the interlamellar region is quite crowded and hydrogen-gas diffusion will be greatly restricted. The most likely mechanism, therefore, is that hydrogen is adsorbed on the surface of Zr(AgPO,), crystallites and the reduction reaction proceeds on the surface.8 The proposed reaction sequence is as follows: (1) H, molecules are adsorbed on the surface and diffuse to the active sites; (2) the reduction reaction proceeds on the surface: Ag atoms and protons are formed; (3) Ag atoms leave the active sites and form clusters; (4) protons diffuse into the interior and Ag+ ions diffuse out, enabling the reaction to proceed continuously on the surface. It is proposed that the initial induction period in this reaction is mainly due to an abnormally slow initial rate of silver-nuclei growth.Several investigationslap 3 3 9 34 have examined nuclei formation in solutions and on crystals and indicated that the formation of nuclei of a new phase is kinetically difficult. Nevertheless, after the induction period the number of nuclei increases at a rate proportional to a power of the time which is usually larger than 2. In other words, if the rate-determining step in our system is the nucleation of silver atoms, then the observed reaction rate must increase after the induction period and its dependence on time must be expressed by a curve having relatively high curvature. We can imagine that somewhere near the end of the induction period processes other than nucleation become rate-deter- mining; a concomitant rate deceleration occurs.Because the initial rates derived from Elovich plots are nearly proportional to the square root of the hydrogen pressure, the reaction mechanism must involve the dissociative adsorption of hydrogen molecules. The active sites thus have to be species which are able to bond with hydrogen atoms. Because it is well accepted that silver metal cannot chemisorb hydrogen,35 the autocatalytic reaction involving the spillover of H, on silver is excluded. The only other possible active species are the ion-exchange sites on the surface where the Ag+ ions are situated. The mechanism thus proposed is as follows:1592 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES k3 ,”:\\ - 0----Ag The protons formed in reaction (14) must diffuse into the lattice while unreacted Ag+ ions diffuse out to the surface in order to continue the reduction processes.The diffusion of H+ and Ag+ ions in zirconium phosphate is expected to occur more rapidly than the reduction reaction on the surface. Jerus and Clear~ield~~ have studied the kinetics of the reaction between gaseous HC1 and Zr(NaPO,), : Zr(NaPO,), + HCl, --+ Zr(HPO,), + NaCl. (15) This reaction should occur on the surface since all of the NaCl is present on the surface and there is no mechanism for chloride-anion diffusion between the layers of ZrP. Thus the HCl at the surface must ionize in order to form protons which then diffuse into the interior. Sodium ions are displaced and diffuse out to the surface; this counter-diffusion of ions was found to be rate-determining.36 A comparable reaction rate should be found in the present study if the counter-diffusion of cations is the rate-determining step.However, the half-life of reaction (1 5 ) in the same temperature and pressure range is one half of that observed in the present study. In addition, the activation energy of the migration of Na+ ions in Zr(NaPO,), is 10.3 kcal m ~ l - ’ , ~ ~ which is significantly smaller than that of the Ag+ reduction reaction (1 5.2 kcal mol-l). Therefore the rate-determining step of the reduction reaction should be some process other than the migration of the cations. Reaction (14) is actually the nucleation of silver atoms, which is slow during the induction period but very fast thereafter. Hence the rate-determining step, after the induction period, is most likely to be the process depicted by reaction (1 3). The rate equation is thus expressed as (see Appendix) rate = k, [k, k,/(k-, k-,)]ip~2[(0-Ag+),] (16) where (0-Ag+), represents the Ag+ ions at the surface ion-exchange sites.Since the surface concentration of Ag+ ions is proportional to the loading of Ag+ ions and to the surface area of the support where c is a constant. Upon substituting into eqn (16) the following equation results: with rate = k’p&,[Ag+], k’ = ck3[k, k,/(k-, k-,)];. This means that the initial reduction rate is directly proportional to the concentration of Ag+ loading and the surface area of the support, but only to the square root of the initial hydrogen pressure. Eqn (18) is similar to the empirical rate equation, eqn (10). The only discrepancy between eqn (10) and (18) is in the order of the surface-area term; 1.4 in eqn (10) and 1 .O in eqn (1 8).One reason that an order greater than oneS. CHENG AND A. CLEARFIELD 1593 is observed may result from the different activities on the sides and on the bases of the ZrP platelets. Electron micrographs show that large silver particles accumulate on the side faces, while fine silver particles cover the surface of the bases. This indicates that the side surface is more active than the basal surface towards the reduction processes. The probable reason is that the windows of the six-sided cavities of ZrP crystallites have a large free distance on the sides perpendicular to the layers; the diffusion of the cations, therefore, is easier in the direction parallel to the layers.As a result, the reduction reaction proceeds more readily at the edges of the layers which compose the side surface of the platelets and is therefore not a linear function of surface area. Another reason for the difference is that the concentration of surface-active sites varies as a function of the surface area. For crystalline a-ZrP, the surface area is increased mainly by decreasing the crystal size. However, the length of the edge on the base and the thickness of the platelets do not vary proportionally as the crystal size changes. Electron micrographs show that the ratio of the surface area of the sides to that of the bases increases as the size of the crystals decreases. Consequently, the dependence of the reduction rate on the surface area of the support should have a power greater than one.Two possible reasons can explain the decrease in the reduction rate as the reaction proceeds. One is a decrease in the concentration of the surface cations, which are the active sites for dissociative adsorption of hydrogen. As the reaction proceeds, the concentration of Ag+ ions in the interior decreases and the proton concentration increases. Since the rate of the diffusion process is proportional to the concentration gradient, the rate of counter-diffusion of Ag+ ions and the protons in the ZrP lattice should decrease as the reduction reaction proceeds. Actually, Jerus and Clea~Iield~~ have shown that the activation energy of cation migration increases as the loading of cations decreases.When the reduction reaction is nearly complete, the cation diffusion is probably so slow that it determines the rate. That explains why the Elovich equation can never be fitted to the decelerating rate in the kinetic curves. The other reason may be the increased diffusional resistance to H, of the silver particles formed on the surface as the particle size increases. As the Ag particles grow and cover the surface, hydrogen gas must travel through the pores or spaces between silver particles to get to the surface of the ZrP crystallites and then diffuse to the active sites buried under the silver particles. That a thick layer of silver particles covers the surface has been shown by the electron micrographs of the reduced samples. We acknowledge financial support of this study by the National Science Foundation under grants CHE79-16160 and CHE81-14613.We thank the Texas A&M EM Center for help in obtaining the electron micrographs. APPENDIX If reaction (1 3) is the rate-determining step, the reaction rate can be expressed as rate = k, [I2]. (20) Application of the steady-state treatment to the intermediates I, and I, gives eqn (21) and (22), respectively : (21) (22) d[I,lldt = k,[I11 [(O-Ag+),l- (k-,[IzI+ k, [I211 = 0 4Illldt = kl([O-Ag+),lP,* - k-,[I,I- k,[I11 [(O-Ag+),l+ k-,[I2I2 = 0. An assumption is made to simplify eqn (21); i.e. k-,[I,] %- k,.1594 METAL DISPERSIONS ON ZIRCONIUM PHOSPHATES Because reaction (13) is the rate-determining step, it seems to be reasonable to assume that its rate constant, k,, is negligible compared with the value k-&].Therefore eqn (21) simplifies (23) to or (24) Eqn (23) can be written in the form k-2[I,I2 = k,[I,1"0-Ag+),l. (25) Upon substituting expression (25) into eqn (22), the following equation results : k111(O-Ag+)slPH2 - kl [Ill = 0. (26) Thus [Ill = (kl/k-l) "O-Ag+)slP,, (27) (28) and [I21 = (kl k,lk-l~-2)~PL2[(o-Ag+~,l. The rate expression is therefore rate = k,(k, k,/k-, k-,):p~,[(O-Ag+),]. (16) K2 [Ill [(O-Ag+)sl- k-2[I,I2 = 0 [I21 = (kZ/k-ZY [41: W-Ag+),l~. J. R. Anderson, Structure of Metallic Catalysts (Academic Press, New York, 1975), p. 163. Kh. M. Minachev and Ya. I. Isakov, ACS Monogr., 1976, 171, 552. P. A. Jacobs, Carboniogenic Activity of Zeolite (Elsevier, Amsterdam, 1977). R. G. Herman, J. H. Lunsford, H. Beyer, P.A. Jacobs and J. B. Uytterhoeven, J. Phys. Chem., 1975, 79, 2388. M. Iwamoto, T. Hashimoto, T. Hamano and S . Kagawa, Bull. Chem. Soc. Jpn, 1981, 54, 1332. H. Beyer, P. A. Jacobs and J. B. Uytterhoeven, J. Chem. Soc., Faraday Trans. I , 1976, 72, 674. A. Clearfield, in Inorganic Zon Exchange Materials, ed. A. Clearfield (C.R.C. Press, Boca Raton, Fla, 1982), chap. 1. A. Clearfield, D. S. Thakur and H. Cheung, J. Phys. Chem., 1982,86, 500. A. Clearfield and S . P. Pack, J. Znorg. Nucl. Chem., 1975, 37, 1283. lo A. Clearfield and J. A. Stynes, J. Znorg. Nucl. Chem., 1964, 26, 117. l1 A. Clearfield and S. Cheng, J. Znorg. Nucl. Chem., 1980, 42, 1341. l2 S. Cheng, Ph.D. Thesis (Department of Chemistry, Texas A&M University, 1982), p. 36. l3 A. Clearfield and G. D. Smith, Znorg. Chem., 1969, 8, 431. l4 T. Inui, T. Ueda, M. Sushin and H. Shingu, J. Chem. Soc., Faraday Trans. I , 1978, 74, 2490. l5 A. Y. Rozovskii, V. D. Stytsenko and V. F. Tret'yakov, Kinet. Catal., 1973, 14, 953. l6 H. H. Voge and L. T. Atkins, J. Catal., 1962, 1, 171. l 7 N. F. H. Bright and W. E. Garner, Nature (London), 1934, 133, 570. l8 W. E. Gamer, W. R. Southon, J. Chem. Soc., 1935, 2, 1705. 2o W. E. Gamer, A. S. Gomm and H. R. Hailes, J. Chem. SOC. 1933, 2, 1393. 2* (a) Ya. Zeidovich, Acta Physicochim. USSR, 1934, 1, 449; (b) S. Roginskii and Ya. Zeldovich, Acta 22 S. Yu Elovich and G. M. Zhabrova, Zh. Fiz. Khim., 1939, 13, 1761. 23 H. A. Taylor and N. Thon, J. Am. Chem. Soc., 1952,74,4169. 24 J. M. Dunoyer, C . R. Acad. Sci., 1942, 214, 556. 25 B. D. Averbukh and G. I. Chaparov, Zh. Fiz. Khim., 1949, 23, 37. 26 J. A. Allen, Austr. J. Chem., 1963, 16, 193. 27 A. van Tiggelen, L. Vanreusel and P. Neven, Bull. SOC. Chim. Belg., 1952, 61, 651. 28 K. Tsutsumi and H. Takahashi, Bull. Chem. SOC. Jpn, 1972,45, 2332. 29 G. R. Eulenberger, J. G. Keil and D. P. Shoemaker, J. Phys. Chem., 1967, 71, 1812. 30 Kh. M. Minachev, G. V. Antoshin, E. S. Shipro and T. A. Navruzov, Izu. Akad. Nauk SSSR, Ser. Khim., 1944, 9, 2081. 31 M. Kermarec, M. F. Guilleux, D. Delafosse, M. Briend-Fauer and J. F. Tempere, Proc. 8th. Int. Symp. Reactions in Solids, ed. J. Wood, 0. Lindquist and C. Helgesson (Plenum Press, New York, 1977), p. 689. 32 B. E. Douglas and D. H. McDaniel, Concepts and Methods of Inorganic Chemistry, (Blaisdell Publishing Co. New York, 1965). 33 A. Ya. Rozovski, V. D. Stytsenks and V. F. Tret'yakov, Kinet. Katal., 1978, 18, 991. 34 J. Haber, A. Kostowska and J. Stoczynski, Proc. 8th Znt. Symp. Reactions in Solids, ed. J. Wood, 35 G. C. Bond, Heterogeneous Catalysis (Oxford University Press, Oxford, 1974). 36 P. Jerus and A. Clearfield, J. Znorg. Nucl. Chem., 1981, 43, 21 17. Groen, Nature (London), 1954, 174, 1836. Physicochim. USSR, 1934, 1, 554. 0. Lindquist and C. Helgesson (Plenum Press, New York, 1977), p. 331. (PAPER 3/ 1665)

ISSN:0300-9599    

DOI:10.1039/F19848001579

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1984, 80, 1595-1604 Ammonia Decomposition on (1 1 10) and (0 0 1) Surfaces of Ruthenium BY CHIKASHI EGAWA, TETSUYA NISHIDA, SHUICHI NAITO* AND KENZI TAMARU Department of Chemistry, Faculty of Science, University of Tokyo, Hongo, Bunkyo-ku, Tokyo 1 13, Japan Received 21st September, 1983 The decomposition of ammonia on stepped Ru(1 1 10) and flat Ru(0 0 1) surfaces has been investigated by Auger electron spectroscopy, low-energy electron diffraction, thermal-desorption studies and kinetic studies. The reaction takes place at ca. 400 K, and N, and H, are formed stoichiometrically. At lower temperatures (< 500 K) the reaction proceeds through the com- bination of N*(ads), which is equilibrated with H, and NH, in the gas phase, and the rate of N, formation is expressed as follows : rNO a exp Cpe,,).The reaction takes place via a typical Temkin-Pyzhev mechanism in which an isotope effect for the hydrogen atoms in the ammonia molecule is observed, i.e. the rate of NH, decomposition is 1.5 times as fast as that of ND,. The formation of N*(ads), which is observed as a peak at 570 K in the thermal-desorption spectra for N,, takes place preferentially on the stepped sites. Therefore, in the transient state the rate of N, formation on the stepped Ru( 1 1 10) surface is an order of magnitude faster than that on the flat Ru(O0 1) surface, although the steady-state reaction rate is twice that on the flat surface. This means that stepped sites on the Ru surface play an important role in breaking the N-H bond.On the other hand, at higher temperatures (> 600 K) the rate of reaction is linearly dependent on ammonia pressure only and independent of hydrogen and nitrogen pressures; the amount of adsorbed hydrogen on the surface was negligible. Since no isotope effect was observed, the reaction is thought to proceed through the recombination of N(ads) on the surface. Amano et aZ.l have reported that Ru metal, like Fe, has a high activity for both ammonia synthesis and decomposition reactions, with surface nitrogen being formed according to the Temkin-Pyzhev mechanism. On the other hand, Danielson et aL2 have reported that molecularly adsorbed ammonia at 100 K on an Ru(0 0 1) surface was desorbed at 133 and 186 K from the surface as molecular ammonia, while atomic nitrogen was formed through an activation process with a very small sticking probability at temperatures between 300 and 500 K, giving a (2 x 2) LEED pattern on dissociation. Ammonia decomposition on flat Pt(1 1 1) and stepped Pt(5 5 7) surfaces was recently investigated by the molecular-beam technique ;3 the decomposi- tion rate on the stepped surface is 16 times faster than that on the flat surface. The step sites on the Pt surface are therefore considered to be effective for N-H bond scission.As reported el~ewhere,~ it has been observed in X-ray and ultraviolet photoelectron spectra that NH(ads) species, as well as atomic nitrogen, was preferen- tially formed on the Ru(1 1 10) surface during ammonia decomposition. Many studies of ammonia decomposition on various transition metals have been rep~rted.~-l~ In this study the reaction mechanism of ammonia decomposition on flat Ru(0 0 1) and stepped Ru( 1 1 10) surfaces, the latter one step per five atomic distances, has been elucidated by studying the behaviour of surface nitrogen species, in a similar I5951596 DECOMPOSITION OF NH, ON RU SURFACES way to the reaction of NO and H, reported in ref.(15), where atomic nitrogen exhibited heterogeneous behaviour and its desorption obeyed the Zeldovich-Rodinsky equation. EXPERIMENTAL The experiments were performed in a stainless-steel ultrahigh-vacuum system equipped with 4-grid LEED-A.e.s. optics and a mass filter, as described previously.16 The Ru(0 0 1) and Ru(1 1 10) planes were cut from a single-crystal rod (99.999% purity, Metal Research Ltd) and prepared by standard methods.The Ru(0 0 1) surface has a hexagonal close-packed structure, while the Ru(1 1 10) surface has a stepped structure, identified as described e1~ewhere.l~ Edge effects were determined in the following manner. After the sample surface was treated with 5 langmuir of adsorbed H,S, both planes were cleaned by Ar-ion bombardment (1.3 kV and 4 pA for 30 min). The turnover frequency of the reaction in the steady state was obtained from the change in partial pressure of the product species from reaction temperature to room temperature for flow-reactor conditions, taking the effective pumping speed and the number of Ru atoms per cm2 into consideration. This completely eliminated any possible effects on the filament of a mass-filter or the wall of the stainless-steel chamber.The amount of adsorbed species on the surface during the reaction was estimated from the area under the desorption peak in thermal desorption spectra in the same manner as the reaction rates. A change of heating rate or a sudden decrease of ammonia pressure did not have an influence on the amount of adsorbed species on the surface. The sample was mounted on a sample holder constructed of oxygen-free copper by spot- welded tantalum wires 0.3 mm in diameter, through which the sample was resistively heated using direct current. A preliminary study employing this sample holder without a sample re- vealed that the Ta wire was inactive in the reaction and had no influence on the measurements of the reaction rate or the amount of adsorbed species.The sample temperature was measured by a Pt-Pt/Rh (13%) thermocouple spot-welded at the centre edge of the sample. The isotope 15NH3 was employed in this study to avoid the overlapping of N, and CO mass peaks. RESULTS AND DISCUSSION The decomposition of ammonia on the Ru(0 0 1) surface took place at ca. 400 K and nitrogen and hydrogen were produced stoichiometrically in the steady state as shown in fig. 1, where the temperature dependence of the rates of N, and H, formation, together with surface concentrations of nitrogen and hydrogen, are given. The rate of ammonia decomposition increased up to ca. 560 K, where the reaction probability for the collision of ammonia molecules was ca. 0.05, and thereafter decreased at higher temperatures under 1.9 x Pa of ammonia.Both nitrogen and hydrogen coverage, which were estimated from the amount of desorbed N, and H, by thermal desorption studies, also decreased with temperature, approaching zero at 770 and 560 K, respectively. Hydrogen is considered to be adsorbed as NH,(ads) species because its desorption temperature is higher than that from the dissociated adsorption of H,. The reaction temperature giving the maximum reaction rate shifted to higher temperatures with increasing ammonia pressure, as shown in fig. 2(A). The temperature at which ammonia decomposition started did not change appreciably with ammonia pressure, but the rate of N, formation on the low-temperature side of the peak in fig. 2(A) depended slightly on ammonia pressure, while that on the higher-temperature side increased with ammonia pressure.The dependence of the rate of N, formation at 520 to 633 K upon ammonia pressure is plotted in the range of ammonia pressure between lop5 and Pa ammonia pressure. However, it became less dependent at higher ammonia pressures and lower reaction temperatures. A similar dependence was obtained on the Pa in fig. 2(B) and shows a linear dependence at ca. 633 K andC. EGAWA, T. NISHIDA, S. NAITO AND K. TAMARU 1597 Fig. 1. Rate of N, and H, formation and surface coverages as a function of temperature during ammonia decomposition on Ru(0 0 1). pNHa = 1.9 x loF5 Pa. 0, N, formation rate; A, H, formation rate; , nitrogen coverage; A, hydrogen coverage. - I d 1 0 E C 0 .- c p 10' Fig. 2. (A) Relation between rate of N, formation and reaction temperature for various ammonia pressures on Ru(0 0 1).pNH,/Pa: (a) 2 x ( d ) 4 x lo-* and (e) 1.2 x Pa. (B) Dependence on ammonia pressure of rate of N, formation during ammonia decomposition on Ru(0 0 1) (open symbols) and Ru (1 1 10) (closed symbols). 0, 0, 520; 0, ., 562; A, A, 598 and 0, +, 633 K. (b) 5 x ( c ) 1.5 x1598 DECOMPOSITION OF NH, ON Ru SURFACES 1 2 nitrogen coverage (X lo-' ) Fig. 3. Rate of N, formation plotted against nitrogen coverage during ammonia decomposition on Ru(0 0 1). 0, 520; 0, 562; A, 598 and 0 : 633 K. Ru( 1 1 10) surface, as shown in fig. 2(B). These results suggest that a different reaction mechanism for ammonia decomposition may be operating at different temperatures. The rate of N, formation as well as nitrogen coverage in the steady state were independent of N, pressure up to 1 O-, Pa at 520-600 K under 3 x Pa of ammonia, although its coverage was much less than saturation.In fact, the dissociative adsorption of molecular nitrogen above room temperature could not be observed, at least up to lo5 langmuir exposure. Consequently, it is concluded that during NH, decomposition surface nitrogen is not equilibrated with molecular nitrogen in the gas phase. As to the effect of hydrogen pressure, both the rate of N, formation and nitrogen coverage were independent of H, pressure at 598 K. Since a negligible amount of hydrogen is adsorbed at higher temperatures in the steady state, N, formation is thought to proceed through the recombination of N(ads) on the surface, as in the cases of W16 and M017 surfaces.In contrast, at lower temperatures the rate of N, formation was reduced, together with nitrogen coverage, with increasing hydrogen pressure. This dependence is more predominant on the stepped surface, as will be discussed later in detail. The rate of N, formation at 520-633 K was plotted as a function of nitrogen coverage in the steady state under various ammonia pressures. As is demonstrated in fig. 3, the rate of N, formation is exponentially proportional to nitrogen coverage, which is similar to NO hydrogenation on Ru, whose rate is one order of magnitude faster than that during ammonia decomposition at the same nitrogen coverage.15 The rate of the reaction and nitrogen coverage in the period before reaching steady state were also investigated.At higher temperatures the rate of the reaction increased with time in accordance with nitrogen coverage on both the flat and stepped surfaces, as shown in fig. 4. On the other hand, as shown in the lower half of fig. 4, at lowerC. EGAWA, T. NISHIDA, S. NAITO AND K. TAMARU 1599 t1103 s Fig. 4. Changes in the rate of N, formation and nitrogen coverage for the transient process up to steady state for ammonia decomposition on Ru(0 0 1). pNHB = 2.7 x lop5 Pa. Upper part, T = 578 K; lower part, T = 520 K. temperatures the rate of N, formation showed a maximum at first and then decreased to its steady-state value, although the total nitrogen coverage increased monotonically toward steady state. This behaviour was more prominent on the stepped surface, as demonstrated in fig.5. The change in the rate in the initial stages, which were not observed on the flat surface for which edge effects had been eliminated, is therefore a characteristic feature of the stepped surface with step sites. To attain steady state even at higher temperatures on Ru(0 0 1) required longer times than on Ru(1 1 lo), which is consistent with activated adsorption as reported by Danielson et al., The initial rate on Ru(1 1 lo), in particular, was one order of magnitude faster than that on Ru(0 0 l), although that on Ru(1 1 10) at steady state was about twice that on Ru(0 0 l), as shown in fig. 2 (B). The changes in the rate of ammonia decomposition on Ru(1 1 10) at lower temperatures were studied in more detail by the thermal desorption of N(ads) during the reaction.As is shown later in fig. 10, a shoulder at 570 K "*(ads)], in addition to the main peak at 750 K [Ns(ads)], was observed during the reaction at 450 K. In the transient state at the beginning of the reaction the amount of N*(ads) showed a maximum in a similar manner to that of the rate of N, formation, although the amount of Ns(ads) grew with time until the steady state was reached. Accordingly, it is suggested that N*(ads) plays an important role in N, formation at lower temperatures and that the formation of N*(ads) may be inhibited by the presence of Ns(ads) on the surface at a later stage of the reaction.1600 DECOMPOSITION OF NH, ON R U SURFACES 2 0 1 t/103 2 Fig. 5. Changes in the rate of N, formation for the transient process up to steady state of ammonia decomposition on Ru(0 0 1) and Ru(1 1 10).pNHI = 2.7 x Pa, T = 520 K. (a) Ru(0 0 1) (edge effects eliminated); (b) Ru(1 1 10). --. Y Q) 2 E .4 Y E 2 z J l -0 -0 - I m .75 3 2 a z 0 5 10 t/min Fig. 6. Changes in the rates of HD and N, formation for the transient process up to steady state of ammonia decomposition on Ru(1 1 10). pNHs = 3.5 x Pa, pHz = pDz = 2.4 x lo-’ Pa, T = 450 K. H,-D, exchange was also investigated on both surfaces during ammonia decom- position. As is shown in fig. 6, the rate of HD formation, which was larger than that on a clean surface at the initial stage in the transient state, was reduced rapidly with the increase of N, formation and then reached a steady-state value that was a quarter of the rate of H,-D, exchange on a clean surface (dashed line).This suggests that the exchange between NH,(ads) formed on step sites and D, proceeds effectively at the beginning of the reaction, because the rate of HD formation on Ru(0 0 1) did not exceed the dashed line. In addition, the decrease in the rate of formation of HD in the later stage indicates that the formation of NH,(ads) is hindered on the surface, which is correlated with the change in the amount of N*(ads) and Ns(ads) mentioned before. The different behaviour of surface nitrogen species formed during ammonia decomposition was obtained on both surfaces in X-ray and ultraviolet photoelectronC. EGAWA, T. NISHIDA, S. NAITO AND K. TAMARU 1601 0 Fig. 7. Effect of H, partial pressure on ammonia decomposition at lower temperatures on Ru(1 1 10).pNH3 = 4.1 x lops Pa, T = 447 K. 0, N, formation rate; A, ON, and 0, nitrogen coverage (X (6N*) Fig. 8. Rate of N, formation plotted against nitrogen coverage (ON*) during ammonia decomposition on Ru( 1 1 10). T = 447 K. spectroscopic measurements as reported elsewhere.* NH(ads) species characterized by an 8.8 eV peak and atomic nitrogen at 5.6 eV below the Fermi level were observed in the u.p. spectra during the reaction at 360 K on Ru(1 1 lo), while atomic nitrogen was detected -on Ru(O0 1) in an amount only half that on Ru(1 1 10). It is thus considered that the activation energy for N-H bond breaking is lowered on step sites and accordingly NH,(ads) species, including atomic nitrogen, are effectively formed on step sites.The relationship between the rate of N, formation and the amount of N*(ads) was investigated under various H, pressures at 450 K on Ru(1 1 10) as shown in fig. 7. The amount of Ns(ads) on the surface was unchanged under different H, pressures; however, both the rate of N, formation and the amount of N*(ads) decreased with increasing H, pressure. This indicates that N*(ads) plays a decisive role in N, formation at lower temperatures. As is demonstrated in fig. 8, the rate of N, formation may be1602 DECOMPOSITION OF NH, ON Ru SURFACES Fig. 9. Isotope effect Open symbols NH,; I n I 1 oi4 1 015 collision frequency/molecule s-l cm-2 on rate of N, formation during ammonia decomposition on Ru(O0 1). filled symbols ND,. 0, a, 520; 0, +, 541; 0, a, 562 and A, A, 598 K.expressed as follows; rN2 GC exp(PB,*). Under the conditions of 520 K and Pa of ammonia the reaction order of ammonia decomposition may be expressed as & & J I E ~ ~ . Accordingly, it is suggested that the formation of N, at lower temperatures proceeds through the combination of N(ads), which is equilibrated with H, and NH, in the gas phase. In other words, the reaction takes place via a typical Temkin-Pyzhev mechanism. Isotope effects on ammonia decomposition were examined by employing NH, and ND,. As is shown in fig. 9, at lower temperatures the rate of N, formation during NH, decomposition was 1.5 times as fast as that during ND, decomposition (the abscissa is expressed as ammonia collision frequency). In contrast, N, formation at higher temperatures is similar for both NH, and ND, decomposition, which is consistent with the independence of N, formation on hydrogen pressure.According to the Temkin-Pyzhev mechanism, the kinetic isotope effect on ammonia decompo- sition would be obtained as (KH/KD)O-,, where KH(KD) is the equilibrium constant 2NH,(ND,) + 2N(ads) + 3H,(D,) for the reaction if the intermediate is N(ads), taking the reaction order described above into account. Its value can be calculated to be 1.4 at 520 K, which is in good agreement with the observed isotope effect. However, if the reaction intermediate is NH(ads), its value would be 1.2, which is a little lower than the experimental value, Finally, the effect of oxygen on the reaction is briefly discussed.As mentioned before, the rate of N, formation during the NO+H, reaction is one order of magnitude faster than that during ammonia decomposition at the same nitrogen coverage. This is in agreement with the different peak temperatures in the N, thermal- desorption spectra; i.e. the desorption temperature of N(ads) during the NO + H, reaction is ca. 650 K, which is 100 K lower than that during ammonia decomposition. Addition of 0, below 1 langmuir during ammonia decomposition on Ru(0 0 1) caused a second peak at 650K, together with the enhancement of N, desorption. It is therefore concluded that the rate of N, formation during the NO+H, reaction is enhanced by the presence of oxygen on the surface (below 1 % of a monolayer). On the other hand, as is demonstrated in fig.10, the presence of oxygen on Ru(1 1 lo), which is preadsorbed by introducing a small amount of 0, (below 0.3 langmuir) at high temperatures (ca. 1200 K), affects the rate of nitrogen desorption in a similar wayC. EGAWA, T. NISHIDA, S. NAITO AND K. TAMARU 1603 (0 ) 450 570 770 TIK Fig. 10. Effect of preadsorbed oxygen on N, thermal-desorption spectra during ammonia decomposition on Ru(1 1 10). 0, exposure: (a) 0, (b) 0.02, (c) 0.07, (d) 0.13 and (e) 0.27 langmuir. Thermal desorption from 450 K, pNHB = 4 x Pa. to the formation of N*(ads). The peak at 570 K in the thermal-desorption spectra was also observed during NO+H, reaction on Ru(1 1 10) at 450 K, where its rate was 1.5 times as fast as that on Ru(0 0 1). However, the nature of the oxygen species on the stepped surface has not been explored in detail in this study.CONCLUSIONS In this study ammonia decomposition has been investigated on stepped Ru( 1 1 10) and flat Ru(0 0 1) surfaces. At lower temperatures (< 500 K) the reaction proceeds through the combination of N*(ads), which is equilibrated with H, and NH, in the gas phase; i.e. the reaction takes place via a typical Temkin-Pyzhev mechanism, where the rate of N, formation may be expressed as rN, cc exp (PON*). The isotope effect of hydrogen in the ammonia molecule is also observed under these conditions. The rate of N, formation during NH, decomposition was 1.5 times as fast as that during ND, decomposition. The formation of N*(ads), which is observed as a peak at 570 K in N, thermal desorption spectra, takes place preferentially on step sites.Especially in the transient state, the rate of N, formation on Ru(1 1 10) is one order of magnitude faster than that on Ru(0 0 1) for which edge effects have been eliminated. Accordingly, it is concluded that step sites on the Ru surface play a decisive role in breaking the N-H bond in ammonia. On the other hand, at higher temperatures (> 600 K) the reaction rate was nearly proportional to ammonia pressure and independent of hydrogen and nitrogen pressures, and no isotope effect was observed. It is thus interpreted that the reaction proceeds through the recombination of N(ads) on the surface, since the amount of hydrogen adsorbed on the surface during the reaction is negligible.1604 DECOMPOSITION OF NH, ON R U SURFACES A. Amano and H. Taylor, J. Am. Chem. Soc., 1954,76,4201. W. L. Guthrie, J. D. Sokol and G. A. Somojai, SurJ Sci., 1981, 109, 390. C . Egawa, S. Naito and K. Tamaru, Surf. Sci. in press. S. R. Logan and C. Kemball, Trans. Faruduy Soc., 1960,56, 144. M. Grunze, F. Bozso, G. Ertl and M. Weiss, Appl. Surf. Sci., 1978, 1, 241. M. Weiss, G. Ertl and F. Nitschke, Appl. Surf. Sci., 1979, 2, 614. 3, 217. K. Kishi and M. W. Roberts, Surf. Sci., 1977, 62, 252. * L. R. Danielson, M. J. Dresser, E. E. Donaldson and J. T. Dickinson, Sug. Sci., 1978, 71, 599. * M. Drechsler, H. Hoinkes, H. Kaarmann, H. Wilsch, G. Ertl and M. Weiss, Appl. Surf Sci., 1979, lo R. P. H. Gasser and D. P. Green, Surf. Sci., 1979, 82, L582. l1 A. Vavere and R. S. Hansen, J. Catal., 1981, 69, 158. l2 C . W. Seabury, T. N. Rhodin, R. J. Purtell and R. P. Merrill, Surf. Sci., 1980, 93, 117. l3 K. Jacobi, E. S. Jensen, T. N. Rhodin and R. P. Merrill, Surf Sci., 1981, 108, 397. l4 J. L. Gland and E. B. Kollin, Surf. Sci., 1981, 104, 478. l5 T. Nishida, C. Egawa, S. Naito and K. Tamaru, J. Chem. Soc., Faruday Trans. I , 1984, 80, 1567. l6 H. Shindo, C. Egawa, T. Onishi and K. Tamaru, J. Chem. Soc., Faruday Trans. I , 1980, 76, 280. l7 M. Boudart, C. Egawa, S. T. Oyama and K. Tamaru, J. Chim. Phys., 1981,78,987. (PAPER 3/ 1667)

ISSN:0300-9599    

DOI:10.1039/F19848001595

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. I, 1984,80, 1605-1616 Use of Supported Rhodium and Cobalt Carbonyls as Catalysts for the CO+H, Reaction Effect of the Support and the Metal BY ALESSANDRO CERIOTTI, SECONDO MARTINENGO AND LUCIANO ZANDERIGHI Universita di Milano and Centro CNR sui bassi stati di ossidazione, Via G. Venezian 21, 20133 Milano, Italy, AND CLAUDIO TONELLI, ANTONIO I ANNIBELLO AND ALBERTO GIRELLI* Stazione sperimentale per i Combustibili, Viale A. De Gasperi 3, 20097 San Donato Milanese, Italy Received 23rd September, 1983 The use of cobalt and rhodium catalysts derived from carbonyl clusters and supported on A1,0,, Al,O,-ZnO, ZrO, and SiO, for the CO+H, reaction has been investigated. Only Rh-containing catalysts were active in the formation of oxygenated compounds (essentially ethanol).In all the runs the selectivity to methane plus ethanol was > 90%. The ratios of the rates of ethanol and methane formation or the rate of ethanol formation and the CO reaction reach a maximum when equal amounts of Rh and Co are present on the catalyst surface. Cobalt alone does not form oxygenates, but when mixed with rhodium it increases the selectivity to ethanol. In recent years there have appeared a large number of papers on the direct catalytic synthesis of chemicals from carbon monoxide and hydrogen. Of special interest is the selectivity between hydrocarbons and oxygenated compounds. Until now only processes for the synthesis of one-carbon-atom compounds can operate with high selectivity, other processes giving a broad spectrum of product distribution (hydro- carbons and oxygenated compounds).Mixtures of higher hydrocarbons are produced with good yields and the use of zeolites has greatly improved the selectivity for low-molecular-weight hydrocarbons. For oxygenated compounds, on the other hand, the situation is less satisfying and much work has been done on this problem. Oxygenated compounds can be produced from CO+ H, using Fe, Ni, Pd, Os, Ir and Pt carbonyls or salts as catalysts., The synthesis of methanol on simple metal centres has been reported by Rathke,3 who used mononuclear carbonyls of Mn and Co, and by Bradle~,~ who used Cu carbonyls. Methanol is a major product in the synthesis of ethylene glycol using Rh as catalyst and Pruett and Walker5 have tested a variety of metal catalysts, including Co, Ru, Rh, Pd, Ir, Pt, Cu, Mn, Sn and Pb, but only Rh and Co produced polyhydric alcohols.Cobalt carbonyl hydride catalyses the formation of methanol and higher alcohols up to C,, as well as their formates, acetaldehyde etc.g The literature on homogeneous catalysis shows that the Group VIII coordination complexes, when active for the CO + H, reaction, catalyse the formation of oxygenated products, while hydrocarbons are produced when the metal alone is present.’ The only exception is a homogeneous Fischer-Tropsch reaction catalysed by II-~(CO)~,, under particular experimental conditions.8 Among the Group VIII metals, Rh seems to have 16051606 CO+H, REACTION OVER Rh AND Co CARBONYLS the unique ability of being able to produce by heterogeneous catalysis two-carbon-atom compounds with good sele~tivity.~ Supported rhodium catalysts have been extensively studied by Ichikawa,lo who has shown that the activity and selectivity of the catalyst depend on the precursor, the support, the preparation conditions and the heat treatment.Catalysts prepared from cluster carbonyls show remarkable selectivity for alcohol synthesis : supported on MgO, CaO and ZnO they promote methanol formation and on TiO,, ZrO,, La203 and Ce,03 they promote ethanol formation. With SiO, support, the activity and selectivity change with the pretreatment conditions. Rhodium appears to be a promising catalyst for the CO + H, reaction as its activity and selectivity change for different precursors and supports. Other metals such as Ni and Co do not have such properties.The aim of this work is a study of the activity and selectivity of catalysts prepared by supporting Rh and Co carbonyls on different oxides in the reaction between CO and H, for the direct synthesis of ethanol. As shown by Ichikawa,lo this reaction occurs with good selectivity; nevertheless an improvement in the activity and selectivity of the reaction is necessary before a new synthesis process can be developed. Our work is experimental and in some aspects is a development of previous studies of the activation of carbon monoxidell and of alcohol synthesis.12 EXPERIMENTAL PREPARATION OF THE CATALYSTS The supports used were y-Al,O, CK-300 (surface area 200 m2 g-l) and SiO, SIL-3 E (surface area 200 m2 g-') from Akzo Chemie and Al,O,-ZnO (surface area 53 m2 g-l) and ZrO, (surface area 70 m2 g-l) from Strem Chemical Inc.A1,0, doped with BaO was prepared from CK-300. The pellets were crushed and the fraction between 50 and 150 mesh used. The precursors were cobalt, rhodium or mixed Co-Rh dodecacarbonyls [Co,RH,-,(CO),, (n = 4,3,2, O)] prepared according to the 1iterat~re.l~ All the supports were thermally treated (Al,O, at 820 K for 6 h, ZrO, at 770 K for 7.5 h) in an inert atmosphere; after cooling at room temperature a weighed amount of the support was suspended in an anhydrous degassed organic solvent (pentane or toluene) and a solution of a known amount of carbonyl in a suitable anhydrous solvent (pentane or toluene) was added drop-wise to the stirred suspension in an inert atmosphere: the cluster was adsorbed by the support and the solution lost the colour of the carbonyl.When all the carbonyl was added the stirred mixture was left to equilibrate for 2 h and the liquid was then drained off. The low adsorption properties of silica did not allow it to adsorb significant amounts of the carbonyls ; therefore these catalysts were prepared by the pore filling technique with heptane as solvent. The solid was dried by evaporation in uucuo (low2 Torr) at room temperature for 2 h. The catalyst so obtained was kept in a nitrogen atmosphere in a sealed glass tube. Table 1 gives details of the prepared catalysts. CATALYTIC ACTIVITY MEASUREMENTS The catalytic runs were performed in a recirculating reactor system (Temkin type) equipped with a condensation trap (liquid nitrogen or water at 273 K) to collect the liquid products.The stainless-steel tubular reactor was 200 mm long with a diameter of 3 mm; in each run ca. 2-3.5 g of catalyst were used. Repeated runs were performed for each catalyst in the 453-523 K temperature range at a pressure of 1.33 x lo5 N m-2 and with H,: CO feed ratios of 3 : 1 and 3:2. The feed and the gaseous products were analysed by gas chromatography with a normal gas- sampling valve on line. Two columns were used for the analysis of the liquid and gaseous products: Porapak Q S (length 7 m; diameter 1/8 in; programmed temperature 9CL200 "C; 5 "CA. CERIOTTI et al. 1607 Table 1. Summary of the experimental conditions of catalyst preparation metal (wt%) catalyst no. precursor solvent support c o Rh 1 2 9 14 15 16 18 19 20 21 22 25 30 31 32 pentane pentane toluene toluene pentane pentane pentane pentane pentane pentane pentane pentane heptane heptane heptane A1203 A1203 A1203 Al,O,-BaO A1203-ZnO A1203-ZnO Al,O,-ZnO ZrO, ZrO, ZrO, Si02 SiO, Si02 ZQ zro2 1.69 0.77 0.74 0.8 1.87 1.17 0.46 - - 0.48 0.51 1.6 - 1.35 2.75 1.28" 1.39 3.08 0.8 1 1.58 0.2 1 0.28 0.89 1.6 - - - a 5wt% BaO.min-' ; carrier gas H,) for CO, CO, and gaseous hydrocarbons and polyethylene glycol 1500 on Chromosorb W-AW (length 3 m; diameter 1/8 in; programmed temperature 60 "C; carrier gas He) for the oxygenated products. The analytical results of each run were checked by the CO material balance. The catalyst was transferred under the nitrogen atmosphere from the sealed tube to the reactor (using this procedure the supported carbonyl never came into contact with air).Preliminary tests in a nitrogen atmosphere had shown that decomposition of the carbonyl begins at 350 K and is complete at 570 K. Before each catalytic run all the catalysts were treated for 2 h at 495 K in hydrogen; all the supported carbonyls were thus decarbonylated to a stable catalyst with highly dispersed metal. RESULTS A1,0, SUPPORT All the catalysts prepared with this support give only hydrocarbons (Cl-C,) as measurable products with a selectivity to methane > 70%. Plots of initial reaction rates (CO molecule atom-l s-l) of the carbon monoxide (H, : CO = 3 : 1) against T1 for the different catalysts (no. 1, 2, 9 and 14) and the temperature coefficients of the reaction rates are reported in fig.1 and in table 2, respectively. All the catalysts containing Rh have approximately the same temperature coefficient but different activities : the catalyst prepared from Co,Rh,(CO),, (no. 2) is significantly more active than that prepared from Rh,(CO),, (no. 9). When the alumina surface is doped with BaO (no. 14) the activity decreases sharply. The influence of the H,:CO ratio in the feed is shown in fig. 1 for two catalysts (no. 2 and 9): in all cases, increasing the CO: H, ratio in the feed decreases the activity, without any significant change in the temperature coefficient of the rate; this behaviour is common to all the catalysts tested, as discussed later.1608 " I v) 2 1 0 - 3 . E CO+H, REACTION OVER Rh AND Co CARBONYLS I " 1.8 1.9 2 .o 103 KIT Fig.1. Plots of CO reaction rate against T1 for different carbonyls on Al,O,: open symbols, CO: H, = 1 : 3; closed symbols, CO: H, = 2: 3. Type of catalyst: no. 1, 0 ; no. 2, V; no. 9, ; no. 14, A. Table 2. Temperature coefficients for the initial reaction rate on Al,O,-supported carbonyls E/kJ mol-1 catalyst no. H,:CO = 3: 1 H,:CO = 3 : 2 1 2 9 14 34 106 112 119 - 100 114 - Al,03-ZnO SUPPORT The temperature coefficients of the initial reaction rates of carbon monoxide over supported CO,(CO),, (no. 16), Co,Rh,(CO),, (no. 15) and Rh,(CO),, (no. 18) are reported in table 3. The main difference with respect to the y-Al,03 support is the inversion of the activity between Rh and Co, mainly because of the strong variation of the temperature factor on supported cobalt.Moreover, the bimetallic Co,Rh, catalyst (no. 15) has behaviour intermediate between the two monometallic catalysts (no. 16 and 18). Detectable amounts of oxygenated products, mainly ethanol, were obtained with Rh, (no. 18) and Co,Rh, (no. 15) catalysts; the selectivity to ethanol was not greater than a few percent, while the selectivity to methane was at least 70%. By changing the feed composition from H,: CO = 3 : 1 to 3 : 2 the activity decreases without any significant change in the selectivity to oxygenated products.A. CERIOTTI et al. 1609 Table 3. Temperature coefficients for the initial reaction rate on Al,O,-ZnO- supported carbonyls E/kJ mol-' catalyst no. H,:CO = 3: 1 H,:CO = 3:2 15 16 18 I26 155 75 106 85 - 10-2 c( I $ 10-3 E 1 o - ~ \ 1.9 2 .o 2.1 2 .2 Fig. 2. Plots of CO reaction rate against T1 for different carbonyls on ZrO,: open symbols, CO:H, = 1 : 3 ; closed symbols, CO:H, = 2:3. Type of catalyst: no. 19, 0; no. 20, A; no. 21, 0; no. 22, v; no. 25, 0; no. 19+no. 20, 1 : 1 , *; no. 19+no. 21, 1:3, +. 1 0 3 K/T ZrO, SUPPORT Plots of the initial reaction rates of the carbon monoxide against T-l are shown in fig. 2 and the temperature coefficients are given in table 4. The Rh, catalyst (no. 21) is more active than the Co, catalyst (no. 19); the activities of the bimetallic catalysts Co,Rh, (no. 20) and Co,Rh (no. 25) are similar and intermediate between those of Rh and Co. Also with these catalysts a decrease in activity has been observed on changing the feed composition from H, : CO = 3 : 1 to 3 : 2.53 FAR 11610 CO + H, REACTION OVER Rh AND Co CARBONYLS Table 4. Temperature coefficients, E/kJ mol-l, for the initial reaction rates on ZrO,-supported carbonyls co methane ethanol catalyst H,:CO H,:CO H,:CO H,:CO H,:CO H,:CO no. = 3:l = 3:2 = 3:l = 3:2 = 3:l = 3:2 ~~ - - - 77 - 19 83 20 112 131 139 141 96 77 21 145 146 162 157 91 84 22 143 25 110 162 124 162 91 - 162 - - - - 19+21 = 1:l 103 90 117 108 54 78 19+21 = 1:3 101 132 117 141 50 80 In the fig. 2 the activities of the bimetallic catalysts Co,Rh, (no. 20) and Co,Rh (no. 25) are compared with that of a mechanical mixture of the catalysts 19 and 21 in 1 : 1 and 1 : 3 ratios; the Co,Rh (no. 25), Co,Rh, (no. 20) and 1 : 1 and 1 : 3 mechanical mixtures of CO, (no. 19) and Rh, (no.21) have practically the same activity. In all cases the activity decreases with increasing partial pressure of CO. To study the influence of the rhodium surface concentration on the activity and selectivity, a catalyst (no. 22) containing 7.5 times less rhodium than the reference catalyst (no. 21) was tested. The initial reaction rate of CO was four times lower without any change in the temperature coefficient (fig. 3). The decrease of the activity is due only to the decrease of the pre-exponential factor of the kinetic constant. Unlike the previous catalysts, all the ZrO, catalysts containing rhodium give measurable amounts of ethanol and other detectable oxygenated products. The amount of methane and ethanol together constitutes > 90% of the products; acetaldehyde does not exceed 1 %.Neither methanol nor C, oxygenates were found in the reaction products. The lack of oxygenated C, compounds, particularly methanol, was confirmed by gas-chromatography mass-spectroscopic analysis. With the most active catalysts the selectivity to ethanol decreases sharply on raising the temperature from 493 to 523 K, while the total CO conversion increases. For this reason only the runs at lower temperature have been considered in the analysis of the ethanol data. The initial rates of ethanol and methane formation as a function of temperature are reported in fig. 3 for Rh and Co-Rh catalysts and the temperature coefficients are given in table 4. The rate of ethanol formation on a bimetallic catalyst is slightly lower than on a monometallic Rh, catalyst (no.21); by decreasing the surface concentration of Rh, (no. 22) the reaction rate decreases. The methane formation rate is higher on monometallic Rh catalysts than on all the others; on bimetallic catalysts the formation rate decreases on increasing cobalt content, this decrease being greater than the decrease in the ethanol formation rate. The values of the temperature coefficients of methane are greater than those measured on the other catalysts and 43 kJ greater than those observed for the formation of ethanol. These high values suggest that methane is formed by CO dissociation ; in comparison with ethanol formation, this reaction is not energetically favoured Ecause of its high activation energy.A. CERIOTTI et al. 161 1 - I v) -..- + E 1 0-' 1 I I 10 -3 " I v1 .t l 0-l 2 10- 1.9 2 .o 2 . 1 103 KIT 2 . 2 Fig. 3. Plots of reaction rate of methane and of ethanol formation on different ZrO, catalysts against T1 (CO:H, = 1 :3). Type of catalyst: no. 20, A; no. 21, 0; no. 22, [7 ; no. 25, 0 ; no. 19+no. 21, 1:1, *;no. 19+no. 21, 1:3, +. SiO, SUPPORT The activity data are shown in fig. 4. The cobalt catalysts are more active than the rhodium and bimetallic catalysts; the latter have an activity a little higher than rhodium. Small amounts of oxygenated compound are formed, mainly ethanol, acetaldehyde and dimethylether. The selectivity to ethanol was 2-5 %. The presence of dimethylether suggests the formation of methanol or of some of its precursors. The values of the temperature coefficients (table 5 ) are similar to those found with Al,O, and Al,O,-ZnO.53-21612 10-5 CO+H, REACTION OVER Rh AND Co CARBONYLS \ \ 0 '\ 1.9 2 .o 2.1 2 . 2 1 0 3 KIT Fig. 4. Plots of CO reaction rate against T1 for different carbonyls on SiO, (CO: H, = 1 : 3). Type of catalyst: no. 30, A; no. 31, 0; no. 32, 0. Table 5. Temperature coefficients for the initial reaction rate on Si0,-supported carbonyls cat a1 ys t E/kJ mol-l no. H,:CO = 3: 1 30 104 31 108 32 123 DISCUSSION Clearly, cobalt and rhodium catalysts behave differently : rhodium catalysts are in general more active for carbon monoxide conversion than the cobalt catalysts. For instance, the reaction rates at 493 K and H,:CO = 3:l are 1.15 x (molecule CO s-l atom-') on Rh/Al,O,, 1.4 x on Rh/ZrO, and 4 x on Rh/SiO,; under the same experimental conditions the cobalt catalysts give the following values: 2.5 x on Co/Al,O,, 3.5 x 10+ on Co/Al,O,-ZnO, 2.5 x on Co/ZrO, and 4.4 x on Co/SiO,.The low activity of the Co/Al,O,- on Rh/Al,O,-ZnO, 2.2 xA. CERIOTTI et al. 1613 Rh, RhcCo, Rh,CO, R h Coj Co, Fig. 5. Selectivity to ethanol as a function of Rh:Co ratio on ZrO,. Open symbols, T = 473 K; closed symbols, T = 493 K. Circles, H, : CO = 3 : 1 ; triangles, H,: CO = 3 : 2. ZnO catalyst suggests an interaction between cobalt and ZnO, for instance the formation of a weakly active or inactive catalytic phase. The influence of the support on the selectivity of Rh catalysts is emphasized by the formation of oxygenated products with ZrO, support but only hydrocarbons with A1,0,.Very small amounts of oxygenated compounds are obtained on Rh/SiO, and on Rh/Al,O,-ZnO; in the latter case this is probably due to the direct action of the zinc oxide, In all the experiments, except those using SiO,, methanol was not detected; the results of chromatographic analysis were confirmed in some cases by mass-spectro- scopic analysis. The absence of methanol in the reaction products is one of the more surprising results obtained with the rhodium catalysts. Moreover, the only products present in significant amounts in all the runs were methane and ethanol. The support has a strong influence on the selectivity of Rh, while it has no influence on the selectivity of Co. The activity of bimetallic Co-Rh catalysts is intermediate between that of mono- metallic Co or Rh catalysts, with a decrease in methane formation on increasing the amount of Co; there is also a decrease in ethanol formation but to a lesser extent.1614 CO + H, REACTION OVER Rh AND Co CARBONYLS Maximum selectivity to ethanol is obtained for a 1 : 1 Co: Rh ratio (fig.5). It is difficult to explain the different catalytic behaviour of cobalt and rhodium on the basis of metal-support interactions or differences in the colligative properties alone ; the work functions of the two metals are 4.12-4.35 and 4.52 eV, respectively, and the first ionization potentials are 7.8 1 and 7.7, respectively. Therefore it appears more reasonable to ascribe the different activity and selectivity to the reactivity of the two metals towards the reacting molecules. Both metals can easily dissociate dihydrogen with the formation of active hydrogen on the surface.On the other hand, experimental data and theoretical calculations point to dissociative chemisorption of CO on cobalt16 and non-dissociative chemisorption on rh0dium.l’ By simple theoretical calculations Miyazakil* has shown that the energy of desorption of CO adsorbed on cobalt is similar to the energy of dissociation, while the energy of desorption is significantly lower on rhodium. This different reactivity of the two metals towards chemisorbed CO is due mainly to the metal-carbon and metal-oxygen bond energies, which are greater for cobalt than for rhodium. Moreover, should the metal-support interaction be strong with cobalt, the localisation of a negative charge on the metal would favour the dissociation of CO.The invariance of the selectivity to hydrocarbon formation with support should be ascribed to this specific reactivity of cobalt. As far as the activity and selectivity of rhodium are concerned our results fal! within the wide range of published data on the reduction of CO. An interesting catalytic feature of rhodium is the variability of its behaviour in the various catalysts tested. While some discussion could arise about the dissociation of CO under our conditions, it is reasonable to assume that the selectivity to oxygenated compounds is related to an active undissociated form of chemisorbed CO. Probably the first reaction step is the formation of formyl (M-CH=O) and/or hydroxymethyl (M-CH,OH) species.It is known that formyl species can be formed by the reaction of a carbonyl group with a hydridic hydrogenlg or by addition of a hydrogen atom to a stretched CO molecule. As no hydride can be formed on the surface of the catalyst, the second reaction path is more likely. In this case the chemisorbed CO must interact with a neighbouring active centre so as to lower the carbon-oxygen bond order to a value that allows the addition of a chemisorbed hydrogen atom. Metal formyl compounds are knownz0 but they are usually unstable. Thermodynamic data on the metal formyl formation reaction are not available; however, one can reasonably assume that it is thermodynamically unfavoured and that the formyl species is present on the surface only as an active intermediate that decomposes or reacts quickly with hydrogen to give more stable surface compounds such as methylene or hydroxymethyl groups.The lack of spectroscopic evidence for the presence of metal formyl species is indirect confirmation of their thermodynamic instability. The formation of methyl species from dissociated CO or hydroxymethyl species is promoted by a positive 6+ charge on the Rh atoms, as happens with highly dispersed Rh on ZrO, and A1,0,.21 Three reaction paths are possible for hydroxymethyl species : hydrogenation to methanol (or methane), dehydroxylation to methylene species, promoted by Bronsted surface acidity, and migration to carbon monoxide adsorbed on the same metal centre. Both the support and the morphology of the dispersed rhodium particles must control one or more reaction pathway. Therefore it is possible that dehydroxylation prevails with small rhodium metal particles dispersed on acidic alumina, producing only hydrocarbons, as has been found previously.16 With larger rhodium particles, where the boundary effects are smaller, or with less acidic alumina both the otherA.CERIOTTI et af. 1615 reaction pathways may occur. Both of them are also possible with rhodium supported on basic or neutral solids. The problem now arises of selectivity to methanol or to ethanol. The published datals indicate that the hydrogenation of CO over supported Rh catalysts prepared by decomposition of Rh carbonyls produces primarily methanol on basic oxides (MgO, ZnO, Be0 and CaO) and primarily ethanol on neutral supports (La,O,, ZrO,, TiO,, Tho, and CeO,). This confirms the change of product distribution in the different chemical environments of the supported rhodium.While methanol is formed by direct hydrogenation of the hydroxymethyl species, for ethanol the prior formation of a carbon-carbon bond must be considered. Let us assume, like many others,,, that a carbon+arbon bond is formed by the migration of a methyl group on chemisorbed carbon monoxide with formation of an acyl species, whose partial hydrogenation would give ethanol. Methyl formation being a necessary step for ethanol formation, the production of methane and higher hydrocarbons should occur when ethanol is produced but not when only methanol is produced. This occurs with basic oxides where the selectivity to methanol is > 90% and only traces of ethanol are present.Therefore methanol and ethanol are formed via two different reaction mechanisms: the first involving hydrogenation of a hydroxy- methyl species and the second the reaction of a methyl group with chemisorbed CO. These two paths need different chemical environments and both may occur on formally neutral supports. Small modifications to the surface of a neutral support can destabilise the hydroxymethyl species and hinder methanol formation, as happens with our Rh/ZrO, tested catalysts. We thank the Consiglio Nazionale delle Ricerche (C.N.R.) for financial support. P. D. Caesar, J. A. Brennan, W. E. Garwood and J. Ciric, J. Catal., 1979, 56, 274; C. D. Change, J. C. W. Kuo, W. H. Lang, S. M. Jacob, J. J. Wise and A.J. Silvestri, Ind. Eng. Chem., Proc. Des. Dev., 1978, 17, 255. W. Keim, M. Berger and J. Schlupp, J. Catal., 1980, 61, 359. J. W. Rathke and H. M. Feder, J. Am. Chem. Soc., 1978,100,3623. J. S. Bradley, in Fundamental Research in Homogeneous Catalysis, ed. M. Tatsui (Plenum Press, New York, 1979), vol. 3. 5 R. L. Pruett and W. W. Walker, U S . Patent, 3833634 (1974), 3 957857 (1976); J. L. Vidal, Z. C. Mester and W. E. Walker, US. Patent, 4 1 15428 (1978); L. A. Cosby, R. A. Fiato and J. L. Vidal, U S . Patent, 4 1 15433 (1978). H. F. Feder and J. W. Rathke, Ann. N. Y. Acad. Sci., 1980, 333,45. J. M. Dayle, A. P. Kouwenhoven, C. A. Schaap and B. van Dort; J. Organomet. Chem., 1979, 174, c 55. * G. C. Dimitries and E. L. Muetterties, J. Am. Chem. SOC., 1977, 99, 2796.P. C. Ellgen, W. J. Bartley and M. M. Bashin, J. Catal., 1978, 54, 120. lo M. Ichikawa, J. Chem. SOC., Chem. Commun., 1978, 466; M. Ichikawa, Bull. Chem. SOC. Jpn, 1978, 51, 2273; M. Ichikawa and K. Shikakura, New Horizons in Catalysis (Elsevier, Amsterdam, 1981), p. 925; Y. Iwasawa, T. Hayasaka and S . Ogasawara, Chem. Lett., 1982, 131; M. Ichicawa, Chem. Tech., 1982, 675. l1 A. Bossi, G. Carnisio, F. Garbassi, G. Giunchi, G. Petrini and L. Zanderighi, J. Catal., 1980,65, 16. l2 A. Bossi, G. Garbassi, G. Petrini and L. Zanderighi, J. Chem. SOC., Faruday Trans. 1, 1982,78, 1029. l3 R. Ercoli, P. Chini and M. Massi Mauri, Chim. Ind. (Milan), 1959,41, 132; S . Martinengo, P. Chini, V. G. Albano, F. Cariati and T. Salvatori, J. Orgunornet. Chem., 1973, 59, 379; S. Martinengo, P. Chini and G. Giordano, Inorg. Synth., 1980, 20, 209. l4 S. Cerny, V. Ponec and L. Hladek; J . Catal., 1977,5, 27. l5 E. Miyazaki, Surf. Sci., 1978, 71, 741. l6 J. A. Rabo, A. P. Risch and M. L. Poutsma, J. Catal., 1978, 53, 295. l7 F. Salymosi, I. Tombacz and M. Kocsis, J. Catal., 1982, 75, 78; M. L. Pontsma, P. A. Ibarbia, A. P. Risch and J. A. Rabo, J. Catal., 1978, 52, 157.1616 CO+H, REACTION OVER Rh AND Co CARBONYLS l8 E. Miyazaki, 1. Catal., 1980, 65, 84. Is J. A. Gladysz, G. M. Williams, W. Tam and D. L. Johnson, J. Organornet. Chem., 1977, 140, C 1; 2o E. L. Muetterties and J. Stein, Chem. Rev., 1979, 79, 479. J. A. Gladysz and W. Tam., J. Am. Chem. SOC., 1978, 100, 2545. G. V. Antoshin, E. Shpiro, D. P. Tkachenko, S. B. Nikishenko, M. A. Ryashentseva, V. I. Avaev and Kh. M. Minachev, New Horizons in Catalysis (Elsevier, Amsterdam, 1981), p. 302. 22 K. Moack and F. Calderazzo, J. Organomet. Chem., 1967, 10, 101; A. Davison and N. Marti, J. Organomet. Chem., 1974, 74, C 17. (PAPER 3/ 1684)

ISSN:0300-9599    

DOI:10.1039/F19848001605

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984, 80, 1617-1629 Kinetics of the Crystallisation of Calcium Oxalate Monohydrate BY EMIL N. RIZKALLA* AND MONA M. MOAWAD Faculty of Science, Ain Shams University, Abbassia, Cairo, Egypt Received 5th October. 1983 The kinetics of precipitation of calcium oxalate monohydrate have been studied conducto- metrically at 298 K for both spontaneous and seeded growth systems. The rate of growth follows a quadratic dependence upon the relative supersaturation, which suggests a surface- controlled growth mechanism. This rate equation holds fairly well for the various solid/solution ratios used in the range of 2.35-10.18 relative supersaturation and for seed concentrations of 0-200 mg dmP3. The effect of sodium tripolyphosphate and phosphonate additives on the precipitation kinetics of calcium oxalate in the absence and presence of well characterised seeds has been investigated at various levels of additive concentration.The inhibiting activity of these additives is discussed in relation to the surface characteristics of the inoculated seeds and structural features of the additive molecules. The fit of experimental data to the Langmuir adsorption isotherm supports a mechanism of inhibition through molecular adsorption of the foreign ions on the surface of the growing crystals. Numerous studies have been reported on the kinetics of crystal growth and dissolution of alkaline-earth oxalates. 1-4 Crystallisation of calcium oxalate is of particular interest, not only from the point of view of analytical chemistry, but also because of its biological importance as one of the main constituents of pathological deposits in the urinary tract.5 Earlier studies on the mechanism of renal stone formation have been concerned with the role of the organic matrix,6 the degree of urine supersaturation with respect to theconstituent ofinterest’ and theinfluence ofurinaryin- hibitors.8 At one stage the uncontrollable deposition of stone minerals was attributed to the role played by the organic matrix;6 however, Vermeulen et aL9 suggested that this mucoid organic matrix is a non-essential phase resulting from protein adsorption on the developing crystalline surface and can act only as a heterogeneous nucleation site for the precipitation of calcium oxalate.1° Previous studies of the precipitation of calcium oxalate have demonstrated the importance of kinetic factors in determining the phase formed during the crystallisation process.11-14 The observed stabilisation of the thermodynamically less stable dihydrate species over the more stable monohydrate phase was attributed to adsorption effects caused by polyphosphate anions on the calculus surface.15-16 On the other hand, Gardner17 claimed that the rate of crystal growth of calcium oxalate in the presence of low-molecular-weight additives such as pyrophosphate and organic phosphonates is independent of additive concentration, and only the induction period preceding growth was found to increase with increasing pyrophosphate concentration. Apparently, these results are in direct disagreement with the results of Robertson et aZ.,18 Sarig et al.19 and Crawford et a1.,20 who demonstrated the in vitro inhibitive effects of 1 -hydroxyethane- 1,l -diphosphonate (HEDP), pyrophosphate anions and polyelectrolytes, as well as the polyanions of heparin and condroitin sulphate.10171618 CRYSTALLISATION OF CALCIUM OXALATE MONOHYDRATE This paper presents a more detailed systematic investigation of the precipitation kinetics of calcium oxalate over a wide range of supersaturation ratios for both spontaneous and seeded growth systems using surface-characterised calcium oxalate seeds. The inhibitory effects of sodium tripolyphosphate and polyphosphonate anions have been studied at various levels of additive concentrations. EXPERIMENTAL REAGENTS AND SOLUTIONS Calcium chloride and potassium oxalate of AnalaR grade were supplied by B.D.H.Commercial-grade sodium tripolyphosphate (STP) was purified by recrystallisation four times from methanol + water mixtures.,l The purity of the final product (Na,P,O,, . 6H20) was checked by elemental analysis. The organic phosphonates (shown below), namely those of ethylenediaminetetra(methy1enephosphonic acid) (ENTMP), hexamethylenediaminetetra- (methylenephosphonic acid) (TENTMP) and 1 -hydroxyethane- 1,l -diphosphonic acid (HEDP), were kindly donated by Monsanto Industrial Chemicals, Louvain-La-Neuve, Belgium, and were used as such except for ENTMP, which was recrystallised as the tetrasodium salt following the procedure described before.22 Elemental analysis of the solid ENTMP and TENTMP and pH titration with standard sodium hydroxide solution (carbonate-free) indicated a purity > 99 % .The active acid content of HEDP was similarly determined by titrating a suitable aliquot of the reagent with NaOH solution, and its purity was calculated on the basis of the amount of base consumed per acid equivalent. CH,-N(CH,PO,H,), I OH CH,-N(CH,PO,H,), HEDP ENTMP TENTMP Stock solutions of the lattice ions and the additives were freshly prepared in conductivity water (doubly distilled, deionised water; conductivity < a) and diluted as required. The exact molarity of the calcium and oxalate ions was then checked using standard analytical methods. 23 Calcium oxalate seeds were prepared by three different methods: seed A was obtained as the spontaneously precipitated so!id formed upon the drop-wise addition of calcium chloride solution (0.01 mol drn-,) to a well stirred potassium oxalate solution (0.01 mol dm-9.The solid was filtered at once, washed well, dried overnight at 383 K and stored. Seed B was prepared essentially as described above, but in this case the precipitate was allowed to age with the mother-liquor for one month. Seed C was prepared by adding urea to an acidic solution of an equimolar mixture of calcium and oxalate ions. The mixture was allowed to stand for one week and the resultant oxalate salt was filtered, washed, dried at 383 K and stored. NITROGEN ADSORPTION MEASUREMENTS The surface areas of the solid materials were measured using a typical volumetric apparatus for nitrogen adsorption.The dead space of the sample bulb and its connecting tubes were calibrated using 99.9% pure helium. The adsorption isotherms were obtained by introducing known volumes of N, gas into the sample bulb and measuring the equilibrium pressures. The amount of adsorbed gas was then calculated assuming ideal behaviour. The saturated vapour pressure of the nitrogen gas was determined periodically during the adsorption run by measuring directly the pressure in equilibrium with liquid nitrogen. The area was calculated by using a value of 16.2 w2 per N, molecule. Table 1 summarises the relevant parameters calculated from the isotherms of the three solids. KINETIC MEASUREMENTS Details of the kinetic procedure are essentially the same as described in an earlier comrn~nication.~~ Supersaturated solutions of calcium oxalate were prepared by adding smallE.N. RIZKALLA AND M. M. MOAWAD 1619 Table 1. Relevant parameters calculated from the isotherms of solids A, B and C affinity constant, surface area average pore seed C /m2 g-' radius, s;/A A B C 3.0 1 .o 3.7 90.2 137.3 133.4 39.3 38.5 28.0 volumes of a pre-thermostatted potassium oxalate solution to the calcium chloride solution so as to make the final concentration of both reactants the same. The rate of growth was then monitored conductometrically. Phosphonate additives and/or calcium oxalate seeds were always added to the calcium solution and the volume of the reacting mixtures was kept constant at 250 cm3. The delivery time of the oxalate solution was determined to be 30 s or less.The temperature in all experiments was maintained constant at 298.00+0.01 K by means of a circulating constant-temperature bath. During mixing of the reactants stirring was achieved mechanically, then the mixture was left static. RESULTS In the absence of additives the following equilibria may be considered: Hf + C20i- HC,O, ; K? (1) Caf&) + C20i~q) * CaC204(aq) ; P I (2) CaC204(aq) + C20i~q) * c(c204)t~q) ; P 2 (3) Ca!zq) + C2Oiraq) CaC20,* H,O,,) ; Ksp. (4) Under the working experimental conditions (pH > 6.5; Ca2+/C20i- = 1 .OO) equilibria (1) and (3) may be n e g l e ~ t e d . ~ ~ . ~ ~ Concentrations of the ionic species in the supersaturated solutions were calculated from the changes in the specific conductance and from the expressions for the total metal and total ligand and the mass-balance equations.Details of the calculations are given by Nancollas and Gardner. l1 Computer-assisted iterative algebraic methods were used which involve successive approximations to the ionic strength, I . The activity coefficients for a z-valent ion, yz, were obtained using the extended Debye- Huckel equation proposed by Davies -log yz = 0.51 15 z2[h/(1 + P ) - 0 . 3 I]. ( 5 ) By analogy with the results of previous crystallisation studiesll the rate of crystal growth, R, in the absence and the presence of seeds and/or additives were analysed using the following relationship : R = - d[Ca2+]/dt = - d[C20,2-]/dt (6) = kobs{([Ca2+t]t [C2042-1t)' - ( g p / , i ) a ) 2 = kobsA2 where K& is the thermodynamic solubility product, [XI, is the ionic concentration of species X at time t and kobs is the observed rate constant.For all calculations the value of K& was taken as 2.00 x The results of crystal-growth experiments from pure solutions are summarised in mo12 dm-6.281620 CRYSTALLISATION OF CALCIUM OXALATE MONOHYDRATE Table 2. Crystallisation of calcium oxalate from supersaturated solutions at 298 K in the absence of additivesa (A) spontaneous growth expt. &a(= Tox) kobsl lo2 no. / mol dm-3 S dm3 mol-1 min-' 1 2.50 2 3.00 3 3.50 4 4.00 5 5.00 4.59 5.71 6.83 7.94 10.18 1.15 1.95 2.69 4.02 6.07 (B) seeded growth k / 102 expt. Tea( = Tax) seed seed dm6 mol-l no. /lo-* mol dm-3 S type conc./mg dm-3 min-l m-2 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 1 S O 2.00 2.50 1 S O 2.50 1 S O 2.00 2.50 2.50 2.50 2.50 2.50 2.50 2.50 2.50 2.50 2.50 2.35 3.47 4.59 2.35 4.59 2.35 3.47 4.59 4.59 4.59 4.59 4.59 4.59 4.59 4.59 4.59 4.59 A A A B B C C C A A A B B B C C C 200 200 200 200 200 200 200 200 80 120 160 80 120 160 80 120 1 60 0.30 0.70 1.04 0.78 1.07 0.58 1.20 1.80 1.10 1.05 1.12 1.07 0.98 1.08 1.70 1.71 1.81 a Tc, and To, represent the initial calcium and oxalate concentrations, respectively, and S is the relative supersaturation given by s = Wa2+l, [C2o:-I,Y - (Gp/YW(g!p/Y3+.table 2, and typical crystallisation rate curves are shown in fig. 1. At all seed concentrations and in the range of supersaturation ratio, S, used in these experiments for spontaneous growth, precipitation commenced immediately, with no evidence of an induction period. The integrated form of eqn (6) A;' - AT' = kobs t (7) where At1 and A;' represent the concentration functions at time t and t = 0, respectively, is plotted in fig.2 and clearly shows an initial slow growth surge, the duration of which is a function of the supersaturation ratio and which is undetectable for solutions with S 2 10 or in the presence of seeds. The effect of additives on the kinetics of growth was studied at S = 10.18 forE. N. RIZKALLA AND M. M. MOAWAD 50 100 3 3t 1621 10 20 t/min Fig. 1. Growth curves for calcium oxalate monohydrate in the absence of additives for spontaneous-growth systems (0, expt. I ) and seeded-growth systems (0, expt. 19; A, expt. 16; 0 ; expt. 22). 4 8 12 I I I I I I 20 40 60 t/min Fig. 2. Plots of the integrated form of eqn (6) for the growth of calcium oxalate at different supersaturation ratios in the absence and in the presence of different seeds: A, expt.1 ; ., expt. 2; A, expt. 3; 0, expt. 4; 0, expt. 16; x , expt. 19; 0, expt. 21. spontaneous-growth experiments and at S = 4.59 and in the presence of 200 mg dm-3 for seeded-growth runs. The results obtained are listed in table 3. Plots of calcium concentration as a function of time in the presence of STP or other phosphonate additives are illustrated in fig. 3-6. DISCUSSION Agreement of the growth-rate data with a parabolic rate law [eqn (7)] over the range of supersaturation ratios and solid/solution ratios used (as seen in fig. 2) rules out bulk diffusion of electrolyte to the crystal surface as the rate-determining step and is indicative of a surface-controlled growth mechanism.1622 CRYSTALLISATION OF CALCIUM OXALATE MONOHYDRATE Table 3.Effect of additives on the rate of growth of calcium oxalate at 298 K (A) spontaneous growth: (Tea = To, = 5.00 x mol dm-3) exp t . additive conc. kobs no. additive / 1 OPs mol dm-3 / 1 O2 dm3 mol-1 min-I 24 25 26 27 28 29 30 31 32 33 34 STP STP STP TENTMP TENTMP ENTMP ENTMP ENTMP ENTMP HEDP HEDP 4.0 10.0 20.0 2.0 4.0 0.8 4.0 10.0 20.0 2.0 4.0 1.18 0.59 0.20 0.94 0.58 2.92 1.29 0.88 0.30 0.86 0.68 (B) seeded growth: (Tea = To, = 2.50 x mol dm-3; seed conc. = 200 mg dmP3) expt. additive conc. k'/ 1 O2 no. seed additive / lop6 mol dm-3 dm6 mol-l rnin-' m-2 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 A A A A A A A A A A A A A A B B B B B B B B B B B B B C C C C C C STP STP STP TENTMP TENTMP TENTMP ENTMP ENTMP ENTMP ENTMP HEDP HEDP HEDP HEDP STP STP STP TENTMP TENTMP TENTMP ENTMP ENTMP ENTMP ENTMP HEDP HEDP HEDP STP STP STP STP TENTMP TENTMP 2.0 4.0 10.0 0.4 2.0 4.0 0.4 2.0 4.0 10.0 0.4 2.0 4.0 10.0 2.0 4.0 10.0 0.4 2.0 4.0 0.4 2.0 4.0 10.0 0.4 2.0 4.0 0.4 2.0 4.0 10.0 0.4 1.2 0.26 0.12 0.02 0.60 0.38 0.20 0.61 0.1 1 0.05 0.02 1.23 0.99 0.56 0.54 0.28 0.16 0.02 0.46 0.18 0.15 0.50 0.11 0.04 0.02 0.75 0.38 0.32 1.20 0.35 0.16 0.03 0.59 0.27E. N.RIZKALLA AND M. M. MOAWAD 1623 10 20 t/min Fig. 3. Effect of STP on the crystal growth of calcium oxalate: 0, expt. 24; 0, expt. 36; A, expt. 50; 0, expt. 64. I I I 1 I i 10 20 tlmin Fig.4. Effect of TENTMP on the crystal growth of calcium oxalate: 0, expt. 27; A, expt. 39; 0, expt. 53; 0, expt. 68. Attempts to use Nielsen's chronomal analysis29 to interpret the kinetic results were successful only over a limited range of 0, where 0 is the degree of reaction. Thus for a diffusion-controlled or surface-controlled reaction of order P, the growth rate can be expressed by the following integrals: ID=K,t=j~~e-I(l-e)-'de (8 4 Ip = KPt = j:B-$(l -B)-PdB (8 b) where K,, and K p are constants related to the final particle size and the initial concentrations of the reactants. Plots of I (obtained through 0) against time are expected to be linear for a particular mechanism. The results obtained here are1624 CRYSTALLISATION OF CALCIUM OXALATE MONOHYDRATE 10 20 t/min Fig.5. Effect of ENTMP on the crystal growth of calcium oxalate: 0, expt. 30; 0, expt. 43; 0, expt. 57. 1 I I I I 10 20 t/min Fig. 6. Effect of HEDP on the crystal growth of calcium oxalate: 0, expt. 34; x , expt. 47; 0, expt. 61. illustrated in fig. 7(A) and (B) for spontaneous-growth [run (l), S = 4.591 and seeded-growth [run (15), S = 4.59, concentration of seed A = 80 mg dmP3] experi- ments, respectively. The linearity conditions appear to be limited to the late stages of growth (6 = 0.3-0.8) in both cases. Also, it is quite difficult to distinguish between I, and I3 operating mechanisms. Nielsen4 concluded from his results on spontaneous precipitation measurements that the growth reaction is controlled by an interphase process for the concentration range 0.3-1 .O mmol dm-3 and is diffusion-limited at higher concentrations.It was also shown that the value of P is apparently constant, 2 < P < 4, in limited concentration ranges. P reaches the limiting value of 4 as the solution becomes infinitely dilute. Reported plots of I, and I4 chronomals for solutions of initial concentration ca. 0.8-1.5 mmol dm-3 show a discontinuity below 6 = 0.25-0.35 with either positive or negative intercepts at t = 0.30 The fractional- order behaviour was attributed to the incompleteness of dissociation of calciumE. N . RIZKALLA AND M. M. MOAWAD 1625 1 tlmin 24 l6 tlmin 8 Fig. 7. Plots of ID and I , ( P = 1-4) chronomals as a function of time for calcium oxalate precipitation; (a) in the absence of seed (expt.1) and (b) in the presence of seed A (expt. 15). oxalate and the true order of reaction was believed to be 4. On the other hand Gardner,13 applying seeded-growth conditions ([Ca] = [Ox] = 0.5 mmol dmP3, seed concentration = 50 mg dm-3) and correcting for the incomplete dissociation of the complex species, concluded that the growth reaction over the range (8 z 0.2-0.9) is controlled by an I, mechanism. In the latter case a negative intercept was also observed. This non-zero intercept might suggest a different operating mechanism for growth in the low-8 range. Such compound-growth mechanisms have been proposed for many other Other factors which might lead to the discontinuity of the Ip against t plots are (1) the precipitation of different hydrate forms and (2) the hypothesis behind the chronomal analysis which presumes (a) homogeneous nucleation and (b) the formation of uniformly spherical particles.29 The first two possibilities are ruled out, since in the first case the discontinuity should be reflected in the plots of (Arl-A~l) against t , whereas in the second case plots for seeded and unseeded conditions are expected to be different. Recent electron-microscopy meas~rements~~1626 CRYSTALLISATION OF CALCIUM OXALATE MONOHYDRATE d 40 20 60 40 20 ~ ~~ " 0 0.2 0.4 0.6 0.8 1.0 PIP O Fig.8. Nitrogen adsorption isotherms for solids A (O), B (0) and C (a). revealed that calcium oxalate can be precipitated in any of the forms monoclinic monohydrate, tetragonal dihydrate or triclinic trihydrate, which are not spherical in shape as chronomal analysis presumes.Increasing the supersaturation ratio corresponds to an increase in the observed rate constants. The logarithmic relationship between S and kobs is satisfactorily linear with slope equal to 2.0-2.2, which is consistent with the former rate equation. For any particular supersaturation the observed rate constant when solutions are inoculated with seeds is two orders of magnitude higher than that observed for a spontaneous-growth experiment, and the extent of this enhancement is a function of the seed weight in solution. According to classical nucleation theories, in pure solutions a fraction of the lattice ions is consumed in forming the embryos which are used as a base for the actual growth stage, whereas in seeded systems the added solid presents suitable sites for direct growth.This also accounts for the slow surge observed in the integrated rate plots in the case of spontaneous growth which disappears upon impregnation with seeds or with increasing S. Accordingly, eqn (7) may be rewritten in the form (9) A t 1 - A71 = k'st where s is a parameter related to the surface area of the solid matrix. At low supersaturation ratios the value of k' was found to be an increasing function of the surface area of the solid (cf. table 1); however, at higher supersaturations solid B behaved in an opposite manner to that expected relative to seed C . This may be assessed by investigating the adsorption isotherms of the three solids. As is seen in fig. 8, solid B shows a type-3 isotherm with a hysteresis loop, indicating that it has high affinity towards capillary condensation.The low value of the constant C for this solid also suggests that the solid has low affinity towards nitrogen adsorption. On the other hand, solids A and C showed an intermediate type of isotherm between 2 and 3, with solid C much closer to type 2. With solid B a decondensation process is expected to take place at low ionic concentrations. Accordingly, more active sites will be obtained, aiding the acceleration of the rate of precipitation.E. N. RIZKALLA AND M. M. MOAWAD 1627 4 8 12 1.8 51 4 b -u" .y W \ 1.4 1.0 1 I I I I 2 4 6 [Add] -'/I O5 mol dm-3 Fig. 9. Langmuir adsorption isotherms in the presence of STP (-) and ENTMP (---): 0, 0, unseeded growth; x , A, in the presence of seed A; 0, in the presence of seed C.Experimental results obtained for calcium oxalate precipitation in the presence of STP and phosphonate additives shown in table 3 for spontaneous- and seeded-growth systems clearly indicate a marked inhibitory effect on the overall precipitation process. The term 'inhibitory activity' will be referred to in this context to describe the reduction in the rate of nucleation of new crystals of calcium oxalate and/or the rate of growth of new or added seed crystals of the salt. If we assume that the retarding action of the additive anions is a result of their adsorption at growth sites on crystal surfaces32 or that prevents their further growth, then the Langmuir adsorption treatment should be valid. If the adsorbed additive of concentration [Add] covers a fraction a of the total available surface, then the rate of adsorption may be expressed as kA[Add] (1 -a) and the rate of desorption as kda, where kA and k, are the corresponding rate constants.At equilibrium, it can be shown that where ko and kAdd are the growth rate constants in the absence and presence of additive, respectively. Plots of eqn (10) are illustrated in fig. 9 for STP and ENTMP additives. The inhibitory action of STP in unseeded systems was attributed to a dynamic adsorption mechanism based on the probability of collision between embryos and STP ions33 which leads to the former's disintegration before reaching a critical size. G l a ~ s n e r , ~ ~ on the other hand, suggested that the active minor component acts as a nucleator either in the form of a complex or a polymeric species,35 leading to the stabilisation of the supersaturated solution.The other class of additives which has proved to be most effective in inhibiting calcification is the organic polypho~phonates.~~ These compounds are characterised1628 CRYSTALLISATION OF CALCIUM OXALATE MONOHYDRATE by the relatively inert P-C-P and P-C-N-C-P linkages which are substituted for the hydrolysable P-0-P bond of the polyphosphates. The relative abilities of the additives in retarding the precipitation process vary with varying experimental conditions and the texture of the inoculated seed. Thus at a level of 4 pmol dm-3 additive concentration, the orders TENTMP > HEDP > ENTMP and ENTMP > TENTMP > HEDP hold for spontaneous and seeGed growth, respectively. In the light of the foregoing discussion, it is reasonable to interpret the observed results in terms of both adsorption and complexation phenomena. Adsorption-rate measurements of 32P-labelled tripolyphosphate on 3-5 pm particles of strontium sulphate precipitated from ‘pure ’ solutions showed that adsorption of STP takes place soon after the birth of 17 A nuclei in a solution containing p~lyphosphate.~~ If we extend these conclusions to other phosphonate systems, it would be expected that adsorption of the larger molecules ENTMP and TENTMP is likely to take place after the birth of larger aggregates of the host lattice, whereas with the HEDP molecule nuclei of smaller dimensions would fulfil the adsorption re- quirements.This sequence holds fairly well for the spontaneous-growth results except for TENTMP. The enhanced reactivity of TENTMP is ascribed to other structural factors. It is generally accepted that one molecule of ENTMP, HEDP or STP is capable of interaction with one active metal site, whereas with TENTMP it is quite difficult to group the four phosphonate ligands simultaneously around one metal site; most probably each -N(CH,PO,H,), group will interact independently. This might suggest that at the same molar concentration TENTMP is capable of inhibiting a greater number of small embryos in comparison with ENTMP. In the presence of seeds, other structural factors such as surface area, affinity towards adsorption of polar molecules, pore sizes and their distribution, rate of adsorption and desorption of the additives from the solid matrix, and cross-sectional areas of the additive molecules are equally important.For one and the same batch of seeds the rate of adsorption is related to the number of binding groups, and the inhibiting activity would be expected to follow the order ENTMP > TENTMP 2 HEDP. A comparison of the k’ values obtained at a constant additive concentration and in the presence of various seeds shows that the degree of inhibition decreases in the order C > B > A with TENTMP, ENTMP and HEDP and A 2 B > C with STP. If the surface area of the solid is the dominant factor, the sequence A 2 C x B should hold. Apparently, the pore radius is an equally important controlling factor. Large molecules such as TENTMP (cross-sectional area estimated to be 55 A2) are expected to block small capillaries, as is the case with solid C.Increasing the pore radius of the solid matrix will allow some calcium and/or oxalate ions to diffuse inside the pores and allow growth to proceed. Introduction of a second molecule of TENTMP will be sterically hindered by the presence of the first. Using smaller molecules such as STP (cross-sectional area ca. 37 A2) an interesting sequence of inhi- bition is displayed. Solid C behaves as the least inhibited seed, whereas solids A and B are retarded almost to the same extent. Following the same argument, surfaces of wide pores are capable of accepting more than one molecule of STP, resulting in a higher degree of retardation, whereas surfaces of narrow pores are not sufficiently blocked by these small molecules in spite of the differences in surface area.We thank Prof. S. A. Selim for her help with the nitrogen-adsorption measurements.E. N. RIZKALLA AND M. M. MOAWAD 1629 G. L. Gardner and G. H. Nancollas, J. Znorg. Nucl. Chem., 1976, 38, 523. S. T. Liu and G. H. Nancollas, J. Znorg. Nucl. Chem., 1976, 38, 515. G. H. Nancollas and N. Purdie, Trans. Faraday Soc., 1961, 57, 2272. A. E. Nielsen, Acta Chem. Scand., 1960, 14, 1654. (a) E. L. Prien and C. Fondel, J. Urol., 1947, 57, 949; (b) C. Langren, Acta Radiol. Suppl., 1956, 1, 133; (c) W. H. Boyce, Am. J. Med., 1968, 45, 673. (a) J. S. King and W. H. Boyce, Ann. N. Y. Acad. Sci., 1963,104,579; (b) B. Finlayson, C. W. Vermeulen and E. J.Stewart, J. Urol., 1961, 86, 355. (a) J. S. Elliot, Surg. Clin. North Am., 1965, 45, 1393; (6) W. G. Robertson, M. Peacock and B. E. C. Nordin, Clin. Sci., 1968, 34, 579; (c) 1971, 40, 365; ( d ) W. G. Robertson, M. Peacock, R. W. Marshall, D. H. Marshall and B. E. C. Nordin, N. Engl. J . Med., 1976, 294, 249. (a) J. E. Howard, U. C. Thomas, L. M. Barker, L. M. Smith and C. L. Wadkins, Johns Hopkins Med. J., 1967, 120, 119; (b) H. Fleisch and S. Bisaz, Am. J. Physiol., 1962, 203, 671; (c) W. G. Robertson, D. S. Scurr and C. M. Bridge, J. Cryst. Growth, 1981, 53, 182. C. W. Vermeulen, E. S. Lyon and W. B. Gill, Invest. Urol., 1964, 1, 370. 10 W. H. Boyce, Proc. Conf: Urolithiases, Physical Aspects (U.S. Natl Acad. Sci., 1972), p. 97. l1 G. H. Nancollas and G. L. Gardner, J.Cryst. Growth, 1974,21, 267. l 2 A. D. Randolph and G. W. Drach, J. Cryst. Growth, 1981, 53, 195. l3 G. L. Gardner, J. Cryst. Growth, 1975, 30, 158. l4 B. Tomazic and G. H. Nancollas, J. Colloid Interface Sci., 1980, 75, 149. l5 C. E. Dent and D. J. Sutor, Lancet, 1971, 7728. l6 H. Fleisch and S. Bisaz, Experienta, 1964, 15, 276. G. L. Gardner, J . Phys. Chem., 1978,82, 864. la W. G. Robertson, M. Peacock and B. E. C. Nordin, Clin. Chim. Acta, 1973,43, 31. l9 S. Sarig, M. Raphael and A. Ron, Zsr. J. Chem., 1973, 11, 635. 2o J. E. Crowford, E. P. Crematy and A. E. Alexander, Aust. J. Chem., 1968, 21, 1067. 21 Q. T. Quimby, J. Phys. Chem., 1954,58, 603. 22 E. N. Rizkalla and G. R. Choppin, Znorg. Chem., 1983, 22, 1478. 23 A. I. Vogel, A Text Book of Quantitative Inorganic Analysis (Longmans, London, 1966). 24 E. N. Rizkalla, J. Chem. Soc., Faraday Trans. I , 1983, 79, 1857. 25 L. G. Sillh and A, E. Martell, Stability Constants of Metal-ion Complexes (Special Publication no. 25, The Chemical Society, London, 1971). G. M. Armitage and H. S. Dunsmore, J. Znorg. Nucl. Chem., 1972, 34, 281 1. 27 G. H. Nancollas, Interactions in Electrolyte Solutions (Elsevier, Amsterdam, 1966). 28 D. J. White and J. H. Nancollas, J , Cryst. Growth, 1982, 57, 267. 2B A. E. Nielsen, Kinetics of Precipitation (Pergamon Press, Oxford, 1964). 30 A. E. Nielsen, J. Colloid Sci., 1955, 10, 576. 31 (a) A. E. Nielsen, Acta Chem. Scand., 1958, 12, 951; (b) R. S. Lee and I. D. Robb, J. Chem. SOC., 32 J. L. Meyer and J. H. Nancollas, Calc$ Tissue Res., 1973, 13, 295. 33 H. Naono, Bull. Chem. Soc., Jpn, 1967, 40, 1104. 34 (a) A. Glassner and S. Skurnik, Zsr. J . Chem., 1968,6,69; (6) A. Glassner, Zsr. J. Chem., 1969,7,633. 35 A. Glassner, S. Sarig and Y. Lamed, J. Cryst. Growth, 1972, 12, 173. 36 M. D. Francis, Calcif. Tissue Rex, 1969, 3, 151. Faraday Trans. I , 1979, 75, 21 16. (PAPER 3/ 1768)

ISSN:0300-9599    

DOI:10.1039/F19848001617

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984,80, 1631-1643 Photoreduction of Methyl Orange Sensitized by Colloidal Titanium Dioxide BY GRAHAM T. BROWN AND JAMFS R. DARWENT* Department of Chemistry, Birkbeck College, University of London, Malet Street, London WClE 7HX Received 17th October, 1983 Colloidal TiO, can sensitize the photoreduction of Methyl Orange to a hydrazine derivative. In steady-state experiments the maximum rate occurs at pH 4.7 in the absence of 0,, but shifts to pH 3.0 when 0, is present. Similar behaviour is observed with Methyl Red despite the difference in pK, values for the two dyes. This suggests that the maximum rate results from changes in the surface composition of TiO,. Flash-photolysis experiments show that electron transfer from TiO, to Methyl Orange depends on the concentration of protonated dye (DH) and the potential of an electron in the conduction band (eCB) of TiO,, so the rate of electron transfer = 8.3 x lo7 [DH][e,,]/ [H+]0.38 (mol dm1-3)-0.62 s-l.Electron transfer to 0, was also measured indirectly in this work and showed a similar logarithmic dependence on pH, rate = 5.5 x 103[0,][ecB]/ [H+]o.42 (mol dm-3)-0.58 s-l. These results can be described by a Taffel relationship when the overvoltage is controlled by the pH of the solution. Charge-transfer reactions at a junction between a semiconductor and an electrolyte may provide a route for the conversion of solar energy.l? Recent research has shown that dispersions of oxide semiconductor particles can photodissociate water into H, and 023-5, and liquid-junction photovoltaic cells could prove to be an attractive alternative to conventional solar cells.6*7 These reactions also play a central role in photography and the photo-oxidation of pigments.In the last two years a number of research groups have studied the electron-transfer reactions in colloids of Ti0,,8-10 CdS2* and CdSe.ls The colloids are transparent and stable for a period of weeks and so the reactions can be monitored directly by conventional photochemical techniques such as flash p h o t o l y ~ i s ~ - ~ ~ ~ l8 and resonance Raman spectroscopy.16~ l 7 Aromatic azo compounds are one of the most important classes of commercial dyes. In general, these dyes have very short excited-state lifetimes because of rapid trans-cis photoisomerization.19 They are stable to visible and near-u.v.light2O? 21 but can be photoreduced in the presence of sensitizers such as chlorophyll22+ 23 or mandelic and provide a useful probe for photoredox reactions. We have carried out a detailed study of the reduction of Methyl Orange photo- sensitized by colloidal TiO, using time-resolved and steady-state techniques. The results provide an insight into the photoreduction of aromatic azo compounds and the effect of pH on interfacial electron-transfer reactions from TiO,. Methyl Orange was chosen since its chemical structure is simpler than that of most commercial azo dyes. As a result the products were more easily analysed. Similar results were found with Direct Red 81 and we expect our observations t o apply to many commercial azo dyes.16311632 PHOTOREDUCTION OF METHYL ORANGE EXPERIMENTAL Steady-state irradiations were carried out using an Applied Photophysics clinical irradiator employing a 900 W xenon lamp and monochromator. Microsecond flash-photolysis experiments were performed with an Applied Photophysics K200 system using copper sulphate filter solutions to remove excitation wavelengths below 290 nm. Visible absorption spectra were recorded with a Perkin-Elmer 554 and 402 spectrophotometer for 1 and 10 cm quartz cells, respectively. Samples for anaerobic experiments were outgassed by repeated freeze-thaw cycles to at least lo-, Torr. This treatment did not affect the absorption spectra of colloidal TiO,. Unless otherwise stated all of the solutions contained 5 x lop3 mol dm-3 acetate or 5 x mol dm-3 borate as pH buffers and 0.1 % polyvinyl alcohol (PVA) as a support for the TiO,. The extinction coefficients taken for protonated and unprotonated Methyl Orange were 4.57 x lo4 and 2.68 x lo4 dm3 mol-l cm-l, re~pectively.~~ PVA (molecular weight 72000), Methyl Orange, Methyl Red, TiC1, (98.5%) and all other chemicals were supplied by B.D. H. (A.R. grade). Water was double distilled. The TiO, sol was made by the method described by Henglein.12 Laser light scattering showed that the average diameter of the particles was 40 nm. When these colloids were stabilized with PVA (0.1 % ) they were transparent and stable for months. RESULTS AND DISCUSSION STEADY-STATE IRRADIATION When an anaerobic solution of Methyl Orange and TiO, (1 0-4 mol dm-3, pH 9.25) is illuminated with ultra-band-gap radiation (1 = 310 nm), the visible absorption band of Methyl Orange disappears and a new peak grows in at 247 nm (see fig.1). No bleaching is observed in the absence of TiO, and visible radiation (A > 400 nm) does not result in the disappearance of dye, which indicates that only light absorbed by TiO, will lead to destruction of Methyl Orange. Indeed, fig. 2 shows that the initial rate of dye loss is directly proportional to the number of photons absorbed by TiO, (5 x 10-5-5 x mol dm-3). On increasing the concentration of TiO, above 5 x mol dmP3 the rate levels off, since most of the incident light is absorbed in a smaller volume close to the window of the absorption cell. Methyl Orange is photoreduced rapidly in this region until the rate is controlled by diffusion of dye from the bulk of the solution.This does not apply to low concentrations of TiO,, where light is absorbed more evenly throughout the 1 cm cell. The reaction obeyed first-order kinetics with respect to Methyl Orange. The spectra in fig. 1 suggest that Methyl Orange (I) is photoreduced to a hydrazine derivative (11): The product (11) absorption maximum at 247 nm was also observed when Methyl Orange was chemically reduced (Zn/AcOH) and disappeared over a period of hours when air was admitted to the solution. From fig. 1 the extinction coefficient of (11) at 247 nm is 2.1 x lo4 dm3 mo1-l cm-l, which is comparable to the extinction coefficient of 1,2-diphenylhydrazine in ethanol ( E , ~ ~ = 2 x lo4 dm3 mo1-1 cm-1).26 Below pH 5 , loss of Methyl Orange was accompanied by spectral changes consistent with the production of protonated hydrazine.After irradiation, addition of base showed that the same product was formed over the pH range 1.5-9.5. Gratzel et al. have shown that the potential of an electron in the conduction bandG. T. BROWN AND J. R. DARWENT 1633 t \ 300 500 h/nm Fig. 1. Photosensitized reduction of Methyl Orange. (TiO,, loF4 mol dmp3; PVA, 0.1 %; borate buffer, 5 x lop3 mol dm-3; R = 310 nm). 0 1 .o d I v) 3 0.75 2 2 \ az" 0.5 0.25 0 [Ti02]/10-3 mol dm-' I5 0.1 0.2 0.3 0.40.5 1.02.0 I I I 1 1 1 I I Fig. 2. Rate of photosensitized reduction of Methyl Orange against TiO, concentration and fraction of light absorbed by TiO, (unless otherwise stated, conditions as in fig.1).1634 PHOTOREDUCTION OF METHYL ORANGE I 1 I 0 2 6 10 PH Fig. 3. (a) Initial rate of Methyl Orange (a) and Methyl Red (0) reduction against pH in anaerobic conditions [TiO,, 5 x mol dm-3; acetate buffer (pH < 6), 5 x mol drnv3; otherwise conditions as in fig. I]. (b) Initial rate of Methyl Orange reduction against pH in aerobic conditions. of TiO, particles shifts linearly with pH to more cathodic 27 Similar effects are well known for oxide electrodes28 and If this was the controlling factor in the reduction of Methyl Orange, the rate could be expected to increase steadily with pH. Fig. 3(a) shows the rate of reduction does increase up to pH 4 but then passes through a maximum at pH 4.6. It would be tempting to explain this on the basis of the equilibrium: + R-NzN-R’ R--N=N-R’ + H+ (2) H pK, = 3.5 (Methyl Orange)3o pK, = 4.8 (Methyl Red).30 The maximum rate, however, occurs at the same pH for Methyl Orange and Methyl Red, despite their significantly different pK, values (the pK, of Methyl Orange was not appreciably affected by colloidal TiO,).Also, varying the concentration of acetate buffer from 0 to mol dm-3 had no effect on the rate of reduction. This suggests that the maximum rate does not result from the pK, of Methyl Orange or catalysis by acetate ions. It may, however, reflect changes in the surface composition of theG . T. BROWN AND J. R. DARWENT 1635 m 0.4 0.3 2 - 0.2 K 3 '0 00 g 0.1 - 0 n A 0 - 1 I 1 I Y 0' 2 4 6 8 10 , t - A s PH Fig. 4. Yield of Methyl Orange reduced per photoflash against pH for aerobic (0) and anaerobic (0) conditions [Methyl Orange, 5 x lop6 mol dm-3; otherwise conditions as fig.3 (a)]. TiO, particles. Schindler et ~ 1 . 3 ~ studied the protonation of TiO, and found two equilibria : --TIOH: -TIOH + H+, pK, = 4.95 (3) -TIOH -TiO- + H+, pK, = 7.8. (4) The steady-state reduction of Methyl Orange will depend on the rate of reaction of photogenerated electrons and holes in TiO,. The fate of these reactants will be influenced by the surface of the particles. When equilibria (3) and (4) are compared with the experimental results, it appears that the optimum surface for reduction of Methyl Orange, in this system, contains both -TiOH and --TIOH; groups. When oxygen is present the maximum rate of reduction occurs at the pK, of Methyl Orange [fig.3 (b)] and the rate drops rapidly to zero as the pH is increased. This results from competition between oxygen and Methyl Orange for reducing species. The following flash-photolysis results show that only protonated Methyl Orange can compete with oxygen for interfacial electron transfer from TiO, and that the rate constant for electron transfer to protonated dye increases with pH. DYNAMICS OF INTERFACIAL ELECTRON TRANSFER TO METHYL ORANGE ANAEROBIC CONDITIONS When samples of Methyl Orange (< 6 x lop6 mol dm-3) and colloidal TiO, (5 x lop4 mol dmp3) were studied using microsecond flash photolysis, loss of dye was followed at the maxima of the visible absorption bands (510 and 470 nm for the protonated and unprotonated species). In anaerobic conditions, the amount of dye reduced per flash was independent of the dye concentration (above 10-smol dm-3) and showed a small decrease on going from pH 1.5 to 8.5 (fig.4). Since the steady-state results show that a hydrazine was the main product at all pH values, fig. 4 indicates that the yield of reducing species, i.e. electrons trapped in the con- duction band of TiO, after the photoflash, is essentially independent of pH. In contrast the dynamics of dye reduction show a much more complex pH dependence, as illustrated in fig. 5 and 6.1636 PHOTOREDUCTION OF METHYL ORANGE Fig. 5. (a) Decay and (b) recovery of Methyl Orange after photoflash (pH 1.45; A = 510 nm). (c) Slow decay of Methyl Orange after photoflash (pH 8.6; A = 470 nm). At pH 1.5 the initial rapid decay of the dye [fig.5 (a)] was followed by a slow partial recovery (ca. 30%) [fig. 5(b)]. Below pH 6 the initial decay obeyed pseudo-first-order kinetics with respect to the Methyl Orange concentration. Fig. 6(a) shows that the observed first-order rate constants (kobs) are proportional to the dye concentration, so that the reaction is overall second order with: rate = kobs[eCB] = k,[Methyl Orange][e,,] (0 where [e,,] represents the concentration of electrons trapped in the conduction bandG. T. BROWN AND J. R. DARWENT 80 6 0 - I v) I 0 - - E “E 4 0 - -a A? \ n 1637 - - 200 - I m . D & 30 I 0 1 2 3 L [Methyl Orange] /pmol dm-’ - I m . III 0 -r“ 200 100 0 i 1 I I I 1 2 3 G [Methyl Orange] /pmol dm-’ Fig. 6. (a) kobs for Methyl Orange reduction against dye concentration under anaerobic conditions [PH 2.2 (O), 3.2 (0) and 5.75 (a)]. (b) kobs under aerobic solutions [PH 2.3 (O), 3.2 (0) and 4.2 (a)].2ot 0 2 4 6 PH Fig. 7. Second-order rate constants for electron transfer from TiO, to Methyl Orange against pH (k, calculated for total Methyl Orange concentration).1638 PHOTOREDUCTION OF METHYL ORANGE of TiO, after the photoflash and kD is the rate constant for electron transfer to Methyl Orange. The effect of pH on k , is shown in fig. 7, where the rate rises to a maximum at pH 3.0 and then decreases with increasing pH. EFFECT OF OXYGEN When samples were equilibrated with air the rate of electron transfer from TiO, to Methyl Orange appeared faster but plots of kobs against dye concentration had the same slope as found for anaerobic solutions and the intercept when no dye is present was no longer zero [see fig.6(b)]. When the samples were purged with oxygen there was a five-fold increase in the intercept whereas the slope of the graph remained the same. Purging the solution will increase the oxygen concentration by a factor of five so this suggests that the intercept is due to electron transfer from TiO, to oxygen. In aerobic conditions, eCB can react with 0, and Methyl Orange, which are both in large excess compared with [e,,], so that removal of eCB is controlled by two parallel first-order reactions : -d[ecB]/dt = (k,[0,] + k,[Methyl Orange])[e,,] = ko bs Ee CBI (ii) where k, is the rate of electron transfer to oxygen. For reduction of Methyl Orange : -d[Methyl Orangelldl = k,[Methyl Orange][e,,] = kD[Methy1 Orange][ecB], exp (- kobs t ) (iii) and hence [Methyl Orange] - [Methyl Orange], = ([Methyl Orange], - [Methyl Orange],) exp (-kobs t ) (iv) and the amount of dye reduced per flash is controlled by the relative magnitude of k,[0,] and k,[Methyl Orange] : kD[Methyl Orange][e,,], k,[0,] + k,[Methyl Orange] (v) [Methyl Orange], - [Methyl Orange], = where the subscripts 0 and a refer to the concentrations immediately after the photoflash (0) and after complete removal of eCB (a).The observed rate constant (kobs), which was calculated from the changes in Methyl Orange concentration, thus results from electron transfer to 0, and Methyl Orange [eqn (ii)] and by extrapolating the data to zero dye concentration it is possible to determine k,.Oxygen also has a marked effect on the amount of dye reduced per flash, as would be expected from eqn (v). This is shown in fig. 4. In anaerobic conditions the yield is effectively constant but with aerobic solutions it falls to zero at pH 5 , since k,[0,] is then greater than kD[Methyl Orange]. The flash-photolysis results are consistent with the following scheme, where TiO, particles are considered as short-circuited photochemical electrodes and reduction of Methyl Orange occurs via electron transfer from TiO, to protonated dye molecules (DH): TiO, + hv + TiO,(h++cB) (light absorption) ( 5 ) TiO,(h+-ecB) + TiO, (charge recombination) (6)G . T. BROWN AND J. R. DARWENT 1639 h+ + PVA -+ H+ +oxidized PVA (hole removal) (7) eCB + 0, -+ 0;- (oxygen reduction) (8) eCB + DH -+ DH ' - (electron transfer to dye) (9) DH'- + eCB + H+ -+ DH; (hydrazine formation) (10) 2DH .- -+ DH; + D- (disproportionation) (1 1) where D- represents ( ~ ~ , ) , ~ ~ , ~ , ~ = ~ ~ , ~ , ~ ~ ~ .Reactions (5)-(7) will occur during the 10 ,us photoflash. (The lifetime of a photogenerated ecB-h+ pair is thought to be of the order of nanoseconds in moderately doped Ti0,.27) Oxidation of PVA and water will trap electrons in the conduction band of TiO,. In the absence of oxygen these reducing species have a lifetime of several hourss and are able to reduce Methyl Orange in a dark reaction. The sharp decrease in k , above the pK, of Methyl Orange (fig. 7) suggests that the rate-determining step is electron transfer to the protonated dye in this pH range.When oxygen is present electrons may be transferred to produce Og-, which is thought to be formed as a chemisorbed species on the surface of Ti0,.32 Below pH 5 the semi-reduced dye will be protonated, since its pK, should be similar to that reported for semi-reduced azobenzene (pK, = 7.1).33 This radical can accept a second electron from TiO, [reaction (lo)] or undergo disproportionation [reaction (1 l)], which will regenerate some of the dye and account for the partial recovery shown in fig. 5 (b). The recovery obeys a mixture of first- and second-order kinetics from which 64 k,, z 2 x lo8 dm3 mol-l s-l assuming that the absorption of DH'- is negligible compared with that of DH at 510 nm. The second-order rate constants (k,) in fig.7 result from electron transfer to the protonated dye but k , refers to the total dye concentration. To correct for this (vii) where k , is the rate constant for electron transfer from TiO, to DH and K, is the equilibrium constant for reaction (2). Fig. 8 shows a plot of log k , against pH, which shows that log k, varies linearly with pH: k, = 8.3 x 107/[H+]0.38 (mol dm-3)-0.62 s-1. Similarly log k,, the rate constant for reduction of oxygen, also increases linearly with pH as shown in fig. 8: k, = 5.5 x 103/[H+]0.43 (mol dm-3)-0.57 s-l. This may reflect changes in the conduction band potential of TiO, as previously observed for reduction of methyl viologen.lo The conduction band potential of TiO, (ECB) is known to shift cathodically with ECB = EEB-0.059 pH (viii) where EEB is the value of ECB at pH 0.Since the standard potential of the couples DH/DH'- and O,/O;- will be independent of pH, the rate of interfacial electron transfer from TiO, to DH and 0, should increase with pH. Such behaviour has been demonstrated by Gratzel et al. for the photoreduction of methyl viologen by colloidal PH:lo7 27-291640 PHOTOREDUCTION OF METHYL ORANGE 8 .O 6.0 m A? M - 5.0 4 .O 1 .o 2 .o 3.0 Fig. 8. Plot of log k, (e) and log k, (0) against pH, where k, and k, are second-order rate constants for electron transfer to 0, and DH at a given pH. PH TiO,. They found that the rate constant varied with the overvoltage ( V ) , and hence pH, in accord with the Taffel equation:l07 27 log (k,,/k",) = - ( 1 - a) nFV/2.303 RT = - (1 - a) nF(E, - Ec,)/2.303 RT hence log (ket/kzt) = - n( 1 - a)(E& - E, -0.059 pH)/0.059 = n( 1 -a) pH - n( 1 - a)(E& - E0)/o.059 (4 where n is the number of electrons in the transfer step, F is the Faraday constant, kEt is the rate constant for electron transfer at the standard potential for the redox couple, Eo is the standard redox potential for DH/DH'- and O,/O;l- and a is the transfer coefficient.In the present case there is reasonable agreement, over the limited pH range available, between eqn (ix) and fig. 8. From these a = 0.62 for DH and 0.57 for O,, which suggests that a symmetrical transition state is involved in both reactions. HIGH pH Above pH 7 the kinetics of anaerobic dye reduction show a distinctiy different behaviour, which was illustrated by fig.5 (c). The reaction occurs over several secondsG. T. BROWN AND J. R. DARWENT 1641 2 .o 1.5 " I wl I - - 0 E E a 1.0 2 *O m e =I 0.5 0 1 2 3 4 5 [Methyl Orange] 11O-j mol dm-3 Fig. 9. Plot of kobs (second order) against dye concentration at pH 8.6 [conditions as in fig. 3 (41- and can be described equally badly by second-order kinetics or a mixture of two first-order reactions. At high pH cis-trans isomerization is also observed, but the photostationary state is re-established within milliseconds of the photoflash so that it does not interfere with the kinetics of dye reduction. Cis-trans isomerization was not observed below pH 7 since the reaction is acid ~atalysed~~ and at low pH it occurs too rapidly to be detected by microsecond flash photolysis. At pH 8.6 the disappearance of dye can be fitted to second-order kinetics and the observed second-order rate constants again increase linearly with dye concentration (fig.9), showing that at high pH the rate can be formulated as: rate = k,,,[Methyl Orange] [eCBl2 ( 4 and the reaction is overall third order. At pH 8.6 interfacial electron transfer will involve unprotonated dye (D-) and the unprotonated semi-reduced dye (D2. -), which both have more negative reduction potentials than DH or DH-. After the photoflash the following reactions could result in the observed kinetic behaviour : D- + eCB D2 * - (12) D2'-+eCB+2H+ + DH;. (13) A disproportionation reaction [similar to reaction (1 l)] would lead to overall second- order kinetics and does not seem to be involved at high pH, since the reaction is third 54 FAR 11642 PHOTOREDUCTION OF METHYL ORANGE order overall.This may reflect the low concentrations of D2'- and the high degree of charge repulsion between the two anions. If reaction (1 3) is slow and rate determining, then fast reduction of D- and reoxidation of D2'- could establish the equilibrium reaction (1 2), in which case rate = K,,k,,[Methyl Orange][e,,12 (xii) kobs = K12k13 = 4.8 x lo6 rnol-, dm6 s-l (xiii) at pH 8.6, where K,, is the equilibrium constant for reaction (12) and k,, is the rate constant for electron transfer from TiO, to the unprotonated dye radical D2' -. CONCLUSIONS Colloidal TiO, can photosensitize the reduction of aromatic azo compounds such as Methyl Orange and Methyl Red. For Methyl Orange and Methyl Red steady-state reduction occurs with a maximum rate at pH 4.7 in the absence of oxygen. This may reflect surface protonation of TiO,.In aerobic solutions the maximum rate shifts to lower pH, since 0, is reduced more rapidly than the unprotonated dye. Flash photolysis shows that the rate of electron transfer from TiO, to Methyl Orange depends on the concentration of protonated dye and the potential of an electron in the conduction band of TiO,. Consequently the rate passes through a maximum at pH 3.0 (fig. 7). This is lower than the pH at which the maximum rate is observed in the steady-state experiments. The origin of this difference is not clear; however, the rate of steady-state dye reduction will also depend on the rate at which the hole (h+) reactions occur, i.e.light absorption, charge recombination and oxidation of PVA and water. These reactions are not reflected in the flash-photolysis measurements, which were made > 10 ps after the photoflash. The rate constant for electron transfer from TiO, to oxygen was measured indirectly in these experiments and also showed a logarithmic dependence on pH, which can be attributed to changes in the potential of the TiO, conduction band. This work was supported by the S.E.R.C. and the London University Central Research Fund. A. J. Bard, J. Phys. Chem., 1982, 86, 172. M. Gratzel, Acc. Chem. Res., 1981, 14, 376. A. V. Bulatov and M. L. Khidekel, Izv. Akad. Nauk SSSR, Ser. Khim. Nauk, 1976, 1902. J. M. Lehn, J. P. Sauvage and R. Ziessel, Noun. J. Chim., 1980, 4, 623.M. Gratzel, E. Borgarello, J. Kiwi, E. Pelizzetti and M. Visca, J . Am. Chem. SOC., 1981, 103, 6324. M. J. Wrighton, Ace. Chem. Res., 1979, 12, 303. A. Heller, Ace. Chem. Res., 1981, 14, 154. M. A. Fox, B. Lindig and C. Chen, J. Am. Chem. SOC., 1982, 104, 5828. A. Henglein, Ber. Bunsenges. Phys. Chem., 1982,86, 241. lo D. Duonghong, J. Ramsden and M. Gratzel, J. Am. Chem. SOC., 1982, 104, 2977. l1 J. Kuczynski and J. K. Thomas, Chem. Phys. Lett., 1982, 88, 445. l2 A. Henglein, Ber. Bunsenges. Phys. Chem., 1982, 86, 301. l 3 A. Henglein, J. Phys. Chem., 1982, 86, 2291. l5 K. Kalyanasundaram, E. Borgarello, D. Duonghong and M. Gratzel, Agnew. Chem., Int. Ed. 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White and A. J. Bard, J. Am. Chem. SOC., 1983, 105, 27. 30 D. D. Perrin, in Dissociation Constants of Organic Bases in Solution (Butterworths, London, 1965). 31 P. W. Schindler and H. Gamsjager, Discuss. Faraday SOC., 1971, 52, 286. 32 M. V. Rao, K. Rajeshwar, V. R. Pai and J. DuBow, J. Phys. Chem., 1980, 84, 1987. 33 E. Hayon and P. S. Rao, Anal. Chem., 1976,48, 564. 34 R. Lovrien, P. Pesheck and W. Tisel, J. Am. Chem. SOC., 1974, 96, 244. (PAPER 3/ 1843) 54-2

ISSN:0300-9599    

DOI:10.1039/F19848001631

出版商:RSC    

年代:1984

数据来源: RSC