摘要:

J . Chem. SOC., Faraday Trans. I, 1988, @(lo), 3567-3573 Conductance Stopped-flow Study of the Association Reaction of Colloidal Spheres with Poly(vinylpyrro1idone) Tsuneo Okubo Department of Polymer Chemistry, Kyoto Uniuersity, Kyoto 606, Japan The equilibrium and kinetic parameters of the association reaction of colloidal spheres with poly(vinylpyrro1idone) (PVP) have been estimated, for the first time, using the conductance stopped-flow (CSF) technique. Colloidal silica spheres (13.5 and 25 nm in diameter) and polystyrene latex spheres (85, 91, 109 and 455 nm in diameter) were used. Relaxation traces of the CSF measurements were obtained by mixing the spheres with PVP solutions. The binary association rate constants (k, = 103-106 dm3 mol-' s-') are much smaller than those of a diffusion-controlled reaction.Both k, and the equilibrium association constant decrease as the size of the spheres increases. Most kinetic studies on fast reactions using stopped-flow and temperature-jump techniques have been restricted to the spectrophotometric observation of the changes in concentrations of reactants, products and/or intermediates with respect to time. However, there are a variety of association processes and chemical reactions which are important systems but cannot be followed by spectrophotometric detection. In this report the association of a water-soluble linear polymer, poly(vinylpyrrolidone), with colloidal spheres of sizes ranging from 13.5 to 455 nm is discussed systematically. These association reactions do not show any clear change in their optical properties, although the changes in conductance that accompany association are expected to be relatively large.The electrical conductance method is most convenient for kinetic studies of macroions because the reactions are accompanied by a significant change in conductance. In this report we applied the conductance stopped-flow (CSF) technique to study association reactions between spherical macroions and linear polymers. The CSF apparatus can analyse fast reactions in the reaction-rate range of 1 ms to several seconds. In this instrument two kinds of solutions are mixed in a mixer cell and the reaction solution is then allowed to flow into an observation cell comprising a narrow tube usually 2 mm in diameter. The conductance change is recorded as a function of time immediately after the flow is stopped.The use of conductance for measuring rates of fast chemical reactions in solution was first studied by Saa1.l However, his technique was basically a flow method and required large amounts of reagents. Sirs2 and Prince3 constructed CSF apparatus independently, and measured the rates of the CO,+OH- reaction and hydrolyses of triphenyl- chlorosilane, respectively. The CSF technique has been studied further by Jaycock and O t t e ~ i l l , ~ Tregloan and La~rence,~ Corkill et aL6 and W ~ l f f . ~ Recently we constructed several types of CSF instruments and applied them to many chemical and physical reaction systems : micellar equilibria of ionic surfactants,8 polyelectrolyte complexa- tion~,~-ll inclusion reactions,12 enzymatic reactions,13 rotational and translational relaxation times of macromo1ecules14-17 and others.18-,l 3567 117-23568 Conductance Study of Colloids with PVP Table 1.Colloidal spheres used diameter charge number charge density sphere /nm per particle /pC cm-2 Ludox- AM 13.5f 1 19 0.54 Ludox-TM 25 f 2 41 0.33 DlC25 85 f 6 3 010 2.1 DlC27 91 _+ 6 4 550 2.8 D 1 B22 109 & 3 8 490 3.6 DlA84 455 + l o 73 600 1 .a Experimental Materials The characteristics (diameters and charge densities) of the colloidal spheres used in this work are listed in table 1. Monodisperse polystyrene spheres DlC25, DlC27, DIB22 and DlA84 were purchased from Dow Chemical Co. The diameters and their dispersions were measured using an electron microscope by the manufacturers.The charge numbers and charge densities of the strongly acidic groups of the spheres were determined by conductometric titrations with a Wayne-Kerr autobalance precision bridge, model B331 mark I1 (Bognor Regis, Sussex). Colloidal silica spheres, Ludox-AM and Ludox-TM, were kindly donated by Du Pont, Japan (Tokyo). Conductance and pH titration curves of the deionized colloidal silica spheres with NaOH solution showed that strongly acidic and very weakly acidic groups coexisted. 22, 23 Poly(vinylpyrro1idone) (PVP) was obtained from Tokyo Kasei Chemical Co., Tokyo (code PVP-K30, molecular weight = 40000). PVP was further purified by repeated precipitation using acetone. Conductance Stopped-flow (CSF) Measurements Details of the CSF apparatus have been described in previous The sample solution from the mixer, which was made of Teflon and a four-jet type, flowed through platinum plates.The platinum-plate electrodes (2 mm x 10 mm) were fixed on opposite walls (2 mm apart) inside the observation cell made of epoxy resin. Ca. 0.2 cm3 of solution was required for each run. A value of 1.30 cm-' was obtained for the cell constant. An ax. current of 50 kHz was given to the Wien bridge. The applied voltage across the cell was kept at 2 V (r.m.s.). The time change of the deviation of the solution conductance from its equilibrium value was amplified in two steps and monitored by a memoriscope and/or digital memory and an X-Y recorded after rectification. The dead time was 1 ms. Results and Discussion Conductance and pH Measurements of the Sphere-PVP Mixtures Fig.1-3 show the changes in conductance (open circles) and pH (crosses) of Ludox-TM, D1C25 and D1A84 spheres in the course of PVP addition. The conductance of the solutions increased and reached a constant value as the PVP was added, whereas the pH decreased. These results support the theory that protons are released in the course of association of the spheres with PVP molecules. It has been clarified that many counterions (protons) of spherical macroions condense on the surface of such a ~ p h e r e . ~ ~ - ~ ' Thus it is highly plausible that the condensed protons are released in the course of association of the sphere with PVP molecules by similar mechanisms to thoseT. Okubo 3569 I I I I i 0 1 2 3 4 5 [PIP]/ 10-3 dm3 Fig.1. Conductance (0) and pH ( x ) titrations of Ludox-TM spheres with PVP solution. Ludox- TM: 13.8 wt%, 0.01 dm3; PVP: 0.0369 mol dm-3. I I ,3.9 9.4 0 2 4 6 [Pw]/lo-3 dm3 Fig. 2. Conductance (0) and pH ( x ) titrations of DlC25 spheres with PVP solution. DlC25: 6.64 wt YO, 0.007 dm3; PVP: 0.0369 mol dm-3. 0 0.5 1 [PvP]/lo-3 dm3 Fig. 3. Conductance titration of DlA84 spheres with PVP solution. DlA84: 3.39 wt YO, 0.02 dm3; PVP : 0.0369 mol dm-3.3570 Conductance Study of Colloids with PVP I 1 Fig. 4. Typical traces of the relaxation of CSF measurements for Ludox-TM-PVP mixtures at 25 "C. [Ludox-TM] = 1.38 wt %; (1) [PVP] = 0.0129 mol dmP3, full scale time = 500 ms; (2) [PVP] = 0,00923 mol dm-3, time = 500 ms; (3) [PVP] = 0.00554 mol dmP3, time = 1 s.I 1 I 0 0.0 1 0.02 0.0 3 Fig. 5. Plots of reciprocal relaxation time of Ludox-AM + PVP (0) and Ludox-TM + PVP ( x ) mixtures against PVP concentration at 25 "C. [Ludox-AM] = 0.638 wt YO, [Ludox-TM] = 1.38 wt Yo. [PvP]/mol dm3 observed often for macrocation-macroanion association reactions.'l When the neutral polymers are added to the colloidal spheres the polymer molecules are adsorbed at the surface of the colloidal sphere, and an adsorbed layer is f ~ r m e d . ~ ~ - ~ ~ The driving force for adsorption is considered to be hydrophobic and dipoledipole interactions. The hydrophobicity of the PVP polymers has been discussed previously andT. Okubo 357 1 , 0 0.01 0.02 0.03 0.04 Fig. 6. Plots of reciprocal relaxation time of 0, D 1 C25-PVP ; x , D 1 C27-PVP ; A, D 1 B22-PVP and 0, DlA84-PVP mixtures against PVP concentration at 25 "C.[DlC25] = 0.366 wt YO, [PvP]/mol c h i 3 [DlC27] = 0.177 wt Yo, [DlB22] = 0.404 wt%, [DlA84] = 0.339 wt %. is very similar to that of other water-soluble neutral polymers such as poly(viny1 alcohol) and poly(ethy1ene glyc01).~~-~' The hydrophobicity of the polystyrene latex spheres is also very high.33*34 Thus it is highly plausible that the driving force for the attraction of latex spheres to PVP is the hydrophobic interaction. Recently the zeta potential of silica colloids was measured in the presence of several water-soluble and neutral polymers, including PVP.38 Strong association between the colloidal silica spheres and PVP molecules was found. The hydrophobicity of colloidal silica spheres is considered to be weak.However, there must be strong dipole-dipole interactions (hydrogen-bonding forces in many cases) between the acidic groups of the silica surface and the side chains of the PVP molecules. Fig. 1-3 show that the solution conductance increased sharply as the concentration of PVP increased, then began to saturate, and remained constant on further PVP addition as described above. By measuring the amount of PVP at which saturation began and the initial amount of colloidal spheres, the numbers of PVP molecules bound to each colloidal sphere in the presence of an excess of PVP molecules were evaluated as 0.30, 90 and 1400, respectively, in the equilibrium state. It is therefore highly plausible that there are multiple association steps, especially for large spheres.CSF Measurements and Kinetic Analyses Typical traces of the relaxation of Ludox-TM-PVP association from CSF measurements are shown in fig. 4. When aqueous solutions of colloidal silica spheres and an excess of PVP were mixed rapidly, a single and strong relaxation trace accompanied by a rapid increase in conductance appeared, i.e. log Alc increased with time. In the presence of an excess of PVP, the following reaction scheme is proposed: PVP PVP PVP (1) sphere sphere - PVP sphere (PVP)2 . . . . - .3572 Conductance Study of Colloids with PVP Table 2. Equilibrium and kinetic constants of association between colloidal spheres and PVP at 25 "C ~~ Ludox- AM 4.2 x 105 1.9 2.2 x 105 Ludox-TM 6.ox 105 4.9 1.2 x 105 DlC25 1.0 x 104 0.8 1.3 x 104 DlC27 1.3 x 104 0.6 2.4 x 104 D 1 B22 6.8 x lo3 0.9 7.0 x 103 DlA84 7.5 x 1 0 3 0.1 7.5 x 104 The most intensive relaxation step observed is assumed to be the first binary association step, z, and the other steps (z2, ..., 7,) are assumed to be much larger than 7,.These assumptions will be justified for colloidal silica spheres, since the equilibrium adsorption is only 0.3 PVP molecules per sphere, as described above. For polystyrene spheres, however, the equilibrium adsorption was much higher. The relaxation time for the first binary association step may contain a large experimental error because of the possible difficulty in deconvoluting subsequent association steps. In these assumptions the observed reciprocal relaxation time, z-l, is given by eqn (2) in terms of forward (k,) and backward (k,) reaction rate constants in the first step (z,), when an excess amount of PVP is present :39 (2) The slopes and intercepts of the linear part of the 7-l us.[PVP] plots give k, and k,, respectively. The equilibrium association constant, K,,,, is given by : z-' = k,[PVP] + k,. Typical plots of z-l us. [PVP] are shown in fig. 5 and 6 for colloidal silica spheres (Ludox- AM and Ludox-TM) and polystyrene latex spheres (DlC25, DlC27, DlB22 and DlA84), respectively. The linearities were satisfactory, and the k,, k, and K,,, values thus obtained are listed in table 2. Clearly, the k, values observed are in the range 103-106 dm3 mol-1 s-l and are much smaller than the rates of diffusion-controlled reactions (109-1010 dm3 mol-1 s-').This means that there is a high barrier for the formation of an activated complex. Note that both k, and Kass decreased as the size of the spheres increased, although the charge densities differ from each other; the Ludox samples have similar and much smaller charge densities than the corresponding values for polystyrene. The exact reason for this is not clear. However, the main reason may be correlated with the highly retarded translational and rotational movements of those large spheres. It is difficult to obtain a correlation between k, and the surface charge density: further study is in progress in our laboratory. I thank Du Pont, Japan (Tokyo) for supplying the colloidal silica samples. References 1 R. N. J. Saal, Recl. Trav. Chim. Pays-Bas, 1928, 41, 73; 264; 385.2 J. A. Sirs, Trans. Faraa'ay SOC., 1958, 54, 201. 3 R. H. Prince, Trans. Faraday Soc., 1958, 54, 838. 4 P. A. Tregloan and G. S. Laurence, J . Sci. Znstrum., 1965, 42, 869. 5 M. J. Jaycock and R. H. Ottewill, Proc. 4th Znt. Cong. Surface Act. (1964), vol. 2, p. 545. 6 J. M. Corkill, J. F. Goodman, S. P. Harrold and J. R. Tate, Trans. Faraday Soc., 1966, 62, 994. 7 M. A. Wolff, Chem. Znstrum., 1973, 5, 59. 8 T. Okubo, H. Kitano, T. Ishiwatari and N. Ise, Proc. R. SOC. London, Ser. A , 1979, 366, 81.T. Okubo 3573 9 T. Okubo, Biophys. Chem., 1980, 11, 425. 10 T. Okubo and A. Enokida, J. Chem. Soc., Faraday Trans. 1, 1983, 79, 1639. 11 T. Okubo, K. Hongyo and A. Enokida, J. Chem. SOC., Faraday Trans. 1, 1984, 80, 2087. 12 H. Kitano, J. Hasegawa, S.Iwai and T. Okubo, J. Phys. Chem., 1986, 90, 6281. 13 H. Kitano, J. Hasegawa, S. Iwai and T. Okubo, Polym. Bull., 1986, 16, 89. 14 T. Okubo, J. Am. Chem. Soc., 1987, 109, 1913. 15 T. Okubo, J. Colloid Interface Sci., 1987, 117, 165. 16 T. Okubo, J. Chem. Phys., 1987, 87, 3022. 17 T. Okubo, J. Phys. Chem., 1987, 91, 1977. 18 T. Okubo, and N. Ise, Polym. J., 1978, 1, 109. 19 M. Sawamoto, T. Higashimura, A. Enokida and T. Okubo, Polym. Bull., 1980, 2, 309. 20 T. Okubo, Dynamic Aspects of Polyelectrolytes and Biomembranes, ed. F. Oosawa (Kodansha, Tokyo, 21 T. Okubo, Makromol. Chem., Suppl., 1985, 14, 161. 22 T. Okubo, J. Chem. Phys., 1987, 87, 6733. 23 T. Okubo, J. Colloid Interface Sci., in press. 24 W. Vanderhoff, H. J. van de Hull, R. J. M. Tausk and J. Th. G. Overbeek, Clean Surfaces: Their Preparation and Characterization for Interfacial Studies, ed. G. Goldfinger (Marcel Dekker, New York, 1970). 25 S. Alexander, P. M. Chaikin, P. Grant, G. J. Morales, P. Pincus and D. Hone, J. Chem. Phys., 1984, 80, 5776. 26 T. Okubo, Ber. Bunsenges. Phys. Chem., 1987, 91, 1064. 27 T. Okubo, J. Colloid Interface Sci., in press. 28 R. R. Stromberg, D. J. Tutas and E. Possaglia, J. Phys. Chem., 1965, 69, 3955. 29 R. H. Ottewill and T. Walker, Kolloid 2. Z. Polym., 1968, 227, 108. 30 E. Killmann and R. Ec Kart, Makromol. Chem., 1971, 144, 65. 31 G. J. Fleer, L. K. Koopal and J. Lyklema, Kolloid 2. 2. Polym., 1972, 250, 689. 32 R. Varoqui and P. Dejardin, J. Chem. Phys., 1977, 66, 4395. 33 T. Okubo, J. Chem. SOC., Faraday Trans. I , 1987, 83, 2497. 34 T. Okubo, Colloid Polym. Sci., 1987, 265, 597. 35 T. Okubo, S. X. Chen and N. Ise, Bull. Chem. Soc., Jpn, 1973, 46, 397. 36 T. Okubo and N. Ise, J. Phys. Chem., 1969, 73, 1488. 37 T. Okubo and N. Ise, J. Phys. Chem., 1970, 74, 4284. 38 T. Okubo, J. Chem. SOC., Faraday Trans. I , submitted. 39 K. J. Laidler, Reaction Kinetics (Pergamon Press, London, 1963). 1982), p. 11. Paper 8/00053K; Received 4th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403567

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1988, 84(10), 3575-3585 Solubilities and Vapour Pressures in the Quinquinary System NaC1-KCl-MgCl ,-CaC1,-H,O Part 1.-Predictions and Measurements at 25 "C Yizhak Marcus" and Neomi Soffer Department of Inorganic and Analytical Chemistry, The Hebrew University of Jerusalem, Jerusalem 91904, Israel Equations for the calculation of the solubilities of halite, sylvite and carnallite, individually or combined, have been set up for the quinquinary system NaCl-KCl-MgCI,-CaC1,-H,O at 25 "C. These equations also yield the vapour pressure of the water above both unsaturated and saturated solutions in this system. The equations take into account cation-cation interactions in mixed solutions having ionic strengths exceeding those of saturated solutions of individual salts.The calculated results are compared with measured values of the vapour pressures and with values of the solubilities reported in the literature. Solubilities of salt components in the common-anion quinquinary system NaCl- KC1-MgC1,-CaCl2-H2O or in ternary or quaternary subsystems thereof have been measured over the years by many authors over a wide range of temperatures.'-'' The tabulations of D'Ans ( 1933)4 and Linke ( 1965)13 summarized the results obtained up to those times. However, only in the last decade or so have serious efforts been made not only to report the measured solubilities in the form of phase diagrams, but also to correlate them by semiempirical expressions. 14, 16, l8 Such expressions naturally give rise also to the vapour pressures of both unsaturated and saturated solutions in these multicomponent systems. However, there are few measurements of such vapour pressures for systems higher than ternary ones with which the applicability of such expressions can be confirmed.lg It is the purpose of this paper (confined to data at 25 "C) and a forthcoming paper2' (dealing with data at 30-50 "C) to present theoretical expressions and their required empirical modifications for the calculation of vapour pressures simultaneously with solubility equilibria.Specifically, the solubilities of halite (NaCl), sylvite (KC1) and carnallite (KMgCl, - 6H,O) in the quaternary system NaCl-KC1-MgC1,-H,O (and in the quinquinary system including CaC1, in addition to the above at a constant mole ratio Ca:Mg = 1:3.5) are described by means of these equations.? Vapour pressures that have been measured for various saturated solutions consisting of these systems or sub- systems are compared with the predictions of these expressions.A special feature of the expressions that are derived is the recognition that the ionic strength of the saturated mixed salt solutions is generally above that of the saturated individual halite or sylvite solutions. Extrapolation of activity coefficient expressions, which are necessarily limited to concentrations of these salts up to saturation, is inadequate to describe accurately the solubility and vapour pressure of the mixed salt solutions at ionic strengths 3 or 4 times higher. Cation-cation interaction terms are required for dealing satisfactorily with these high concentrations.The required term can be provided by a single solubility datum of the relevant ternary system. t These systems are relevant to the Dead Sea and to processes for the production of potash therefrom. 35753576 Solubilities in the System NaCl-KCl-MgCl,-CaCl,-H,O Theory The solubilities of halite, sylvite and carnallite are given by the following solubility constant expressions : = Ks, ha1 ; asyl(sat) = Ks, s y l ; = Ks, car‘ (1) ahal = mNaf mCl- r: NaCl (2) asyl = mK+mCl- &KC1 ( 3 ) scar = mKi mMg2i m t l - r’, KCI r: MgCl, (4) where m is the molality, y + the mean ionic activity coefficient and aH the activity of water. The latter is related-to the vapour pressure, p , of the solution (po being that of pure water) : and to its osmotic coefficient, Qs The activities a for both saturated and unsaturated solutions are given by P = p0a€I2o ( 5 ) uHzO = exp[-O.O18015QsC mj(zi+ l)] i where 0.018015 kg mol-1 is the molar mass of water, z is the charge on the cation and the summation extends over all the salts j present in the solution.Expressions are therefore required that relate the activity coefficients y+ of the salts and the osmotic coefficient # of the solution to the composition and totalconcentration of the solution. Several authors have proposed sets of equations for this purpose, and some of these Sets14’21-28 have been evaluated and compared previously by one of us.19 Other Sets169 18,297 30 have also been examined as to their applicability to the present problem.The conclusion is that the approach of Lietzke and Sto~ghton~~y 28 is not only the simplest but is also readily amenable to application beyond the solubility limits of individual salts, as will be shown. The highly useful and widely employed approach of Pitzer and c o ~ o r k e r s ~ ~ * ~ ’ has been tested for the ternary sub-systems of the present quinquinary system (but not at all for the KCl-MgC1,-H,O subsystem) only up to ionic strengths < 8,26 which is not adequate for the present purposes. According to Lietzke and Sto~ghton,~’*~* the activity coefficient of a salt component i in our common ion mixture is In y+i = In yo+i+0.5 - C xi(ziz,” In y&-ln y”,) (7) i where y:i is the activity coefficient of the salt i (and similarly for salt J ) in its binary aqueoussolution at the same ionic strength I as that of the mixture: I = 0.5 C zj(z,+ l)mi i and xi is the ionic strength fraction of salt j in the mixture: (9) The summations in eqn (7) and (8) extend over all component salts j , including j = i [which cancels out in eqn (7)].The osmotic coefficient of the solution is given by the weighted mean : (10) xi = fzj(zj + 1) mJI. Qs = C (zj + 1) mj $p/C(zj + 1) mi i where Qs; is the osmotic coefficient of the salt j in its binary aqueous solution at the same ionic strength I as that of the mixture.Y. Marcus and N . Sofer 3577 Table 1. The parameters A;, [eqn (13)] for the activity and osmotic coefficients [eqn (1 1) and (12)] of the individual salts i at 25 "C k i = NaCl i = KC1 i = MgC1, i = CaCl, 0 1.7440 1.371 7 1.7452 1.6129 1 - 0.006 75 - 0.026 88 0.051 53 0.045 66 2 0.023 05 0.021 29 0.01078 0.008 57 3 - 0.002 794 - 0.003 025 -0.00643 - 0.000 274 4 1.095 6 x 7.900 x 1.265 x 0" " A value of 0 signifies a value < 1 x lo-'.The required Y?~(Z) and &(I) data can be inter- or extra-polated from any of several suitable expressions. Those of Meissner and are not sufficiently accurate, and those of Pitzer and Mayorga31 are limited to 4.5 and 2.5 mol kg-' for MgC1, and CaCI,, respectively, which are too low. The critically evaluated compliations by Hamer and Wu3, for 1 : 1 electrolytes and by Goldberg and N ~ t t a 1 1 ~ ~ for 2: 1 electrolytes, however, serve our purposes well (for 25 "C). These express the activity coefficient y:$ - by means of a Debye-Hiickel-type term and a power series in the ionic strength I : In Y:i = - Zi SZi( 1 + Aio C Aik[Z,(Z, -I- 1)/2]-k Ik ( 1 1) k where S = 1.1762 (at 25 "C), and Aio and the set of At, are salt-specific parameters.The osmotic coefficient q5: corresponding to y l i according to the Gibbs-Duhem relationship is q5; = 1 - zi SA;; I-'[( 1 + Aio 1;) - 2 In ( 1 + Aio Zi) - ( 1 + A , I;)-'] + c k(k + &[Zi(Z, + 1)/2]-k Zk. (12) k The quoted sources list Aik parameters up to k = 3 for NaCl and KCl,32 up to k = 4 for MgC1233*'i and up to k = 7 for CaC1,.33 For the sake of uniformity and comparisons with values at other temperatures,20 the selected activity coefficient values32. 33 have been refitted where necessary to eqn (1 1 ) with 5 parameters for each salt by means of a least- squares program,35 with the results presented in table 1 , where for k 2 1 Aik = k(k + [Zi(Zi -I- 1)/2IPk &.(13) If eqn ( l H 1 3 ) are used to calculate the solubilities of halite and sylvite in a solution containing, say, 3 mol kg-' MgCl,, the difficulty arises of the necessity to estimate y: of these salts at Z > 9 rnol kg-1, whereas the parameters are valid only up t o t h e saturation molality of the two 1 : 1 salts. When the parameters are used beyond these limits (4.8 mol kg-l for KCl and 6.1 mol kg-l for NaCl) the calculated solubilities differ by up to 10 %. It is necessary to take into account ' supersaturation ' values of the activity coefficients. An example of this, regarding the osmotic coefficients of NaCl in ternary solutions containing MgC1, or CaCl, at ionic strengths up to 15 mol kg-', is given by Teng and coworkers.s In the present case it has been found expedient to add a correction term that modifies Ai4 as follows: A:4 = Ai4 + Bijmi (14) where Bij is selected in such a manner to make the calculated solubilities of i andj 0' being the more soluble salt) to conform to the measured ones for the singular point of the ternary solution saturated in both salts.In multicomponent systems higher than ternary such correction terms are additive : A: = A; + Bij mi + Bih mh + . . . . In this manner both3578 Solubilities in the System NaC1-KCl-MgC1,-CaC1,-H2O Table 2. Molalities at the singular points in ternary systems at 25 "C and the derived interaction parameters Bi, [eqn (14)] 5.10 2.23 - 0" halite + sylvite 0.105 - 5.77 - 1.3 x halite + bischofite - 0.65 4.05 1.1 x sylvite + carnallite a A value of 0 signifies a value d 1 x lo-'.the activity coefficients and the osmotic coefficients can be corrected empirically for cation-cation interactions. For these calculations the vapour pressure of water and the solubility products Ks,i are still required. The former is p" = 3.168 kPa at 25 0C.37 The solubilities of halite, sylvite and bischofite (MgC12-6H,0) at 25 "C are 6.144, 4.827 and 5.810 mol kg-l, re- ~pectively,~, l3? 17,34 and the activity coefficients of these saturated solutions, according to eqn (1 1) and (1 3) and the parameters in table 1, are 1.004,0.574 and 32.7, respectively. The activity of water in a solution of MgCl, saturated with bischofite at 25 "C is 0.3278.34 Hence, according to eqn (1) and (2), log Ks,ha, = 1.580, logKs,syl = 0.885 and logK&, = 4.515.The solubility product of carnallite poses a more serious problem, since it dissolves incongruently. Solubility data for solutions that are in equilibrium with solid carnallite (and sylvite too, in most cases) are not very consistent: values of 4.0,14 4.1016 and 4.2218 have been reported for IogK,,,,,. Our solubility data support best the last value : 4.22 k 0.04. The singular points in the ternary solutions (not involving CaCl,) and the parameters Bij are shown in table 2. With these quantities as input parameters, solubilities have been calculated by ' computer experiments' in which water is ' evaporated ' isothermally from an unsaturated salt mixture until ai of one of the salts reaches its Ks,i (within a chosen small tolerance). This salt is then removed from the solution (along with the 6 mol of water per mol of salt in the cases of bischofite and carnallite) in appropriate amounts, as more water is 'evaporated' until the ai of the next salt to precipitate reaches its Ks,i value.38 Vapour pressures and saturation lines and saturation surfaces in the phase polyhedron result from these calculations, as presented in the Results section.Experimental A Texas Instruments quartz Bourdon precision pressure gauge (model 145) was used. Its readability was ca. 1 Pa, and it gave either absolute pressures or pressure differences, according to the Bourdon cell used. The gauge was connected to a vacuum system to which the measuring cell was also attached.The latter was equipped with a small glass- encased magnetic stirring bar, and could be surrounded either with a Dewar vessel for freezing its contents or with a thermostatic bath under which the magnetic stirrer could be placed. The temperature was controlled to k0.2 K. Solutions were made up by weight from analytical-grade salts and triply distilled water. A solution (with excess solid, if saturated) was placed in the cell and frozen under pumping, to remove dissolved air. It was then left to thaw in the thermostatic bath to equilibrate with the vapour and with any excess solid present. The vapour pressure was then measured as a function of time until a constant value was obtained. The reproducibility of the vapour pressure determinations was f 0.04 kPa, owing mainly to the uncertainty in the temperature.For instance, our measurements gave p" = 3.146+0.030 kPa (lit.37 3.168 kPa).Y . Marcus and N . Sofler 3579 Table 3. Measured and calculated vapour pressures at 25 "C of multicomponent solutions mNaCl mKCI mMgCI, saturated with obsd" calcd 1.75 0.98 0.61 0.34 0.47 0.32 0.19 0 0.17 0.34 0.38 0 0.39 0.25 0.17 0.05 1.56 3.01 3.42 0.99 3.64 3.9 1 4.20 4.98 0.43 0.86 0.98 0 1.04 1.12 1.20 0 none ha1 ha1 ha1 ha1 + car ha1 + car ha1 +car syl +car 2.40 1.70 1.58 1.35 1.36 I .34 1.28 1.48 2.40 1.75 1.58 1.33 1.48 1.38 1.26 1.38 a The uncertainty is estimated at f 0.04 kPa. Fig. 1. The vapour pressure, p , in kPa, of aqueous solutions of magnesium chloride: (---) alone;' 0, saturated with sylvite; a, saturated with halite, at 25 "C.Curves were calculated from eqn (9, (6), (10) and (12). Vapour Pressures The vapour pressures of certain saturated multicomponent solutions at 25 "C were measured, and are compared in table 3 with the values calculated from eqn (9, (6), (10) and (12). In addition, the vapour pressures of ternary solutions containing MgC1, and saturated with either halite or sylvite were measured at 25 "C, and are compared in fig. 1 with the values calculated from the abovementioned equations.3580 Solubilities in the System NaCl-KCI-MgC1 ,-CaCl ,-H,O KCl NaCl Fig. 2. Phase diagram for the solubilities of halite and sylvite in the ternary system NaCl-KC1- H,O at 25 "C; the calculated curve [eqn (1)-(3), (7)-(9), (1 1) and (1 3)] is compared with data of d'Ans4 and of Osokorova et aZ.3 Solubilities The calculated solubilities of halite and sylvite in the ternary system NaC1-KCl-H,O at 25 "C are compared with the available experimental values in fig.2, in terms of the mass fractions w (wt %) of the component salts. The mass fraction of salt i is related to the molalities by wi = lOOrn,M, 1 +c rniMi (15) i'( j 1 ( where M is the molar mass of a salt (in kg mol-l) and the summation extends over all salts, including j = i. The converse relation gives (16) mi = wi/Mi 100-C wi . The data can be summarized by the following quadratic solubility curves: wNaCl = 26.64 0.08 - (0.689 k 0.045) wKCl + (0.0126 0.0037) Wkcl (174 wKCl = 26.35 k 0.05 - (0.935 & 0.01 5 ) wNaCl + (0.0098 If: 0.0007) wkaCl.(1 7 b) Calculated solubilities of halite and sylvite in the ternary systems KCl-MgC1,-H,O and NaCl-MgC1,-H,O and in the quaternary system NaCl-KCl-MgC1,-H,O at 25 "C are compared with the available experimental values in fig. 3 (in terms of molalities). Note that the data for systems involving KCI extend only up to the carnallite point (see table 2), beyond which the solutions become saturated in carnallite. The solubility curves for the system KCl-MgC1,-H,O around the carnallite point and at higher MgCl, molalities are shown in fig. 4 (cf the corresponding figure of Wood)." For quaternary systems it is necessary to define solubility surfaces, and again quadratic equations are found to be sufficient, but in this case a cross-term is required.Y.Marcus and N . Sofer 358 1 I I 0 2 4 mMgCl, Fig. 3. The solubilities of halite [(a) and (b)] and sylvite [(c) and (43 in aqueous MgCl, solutions at 25 "C. Filled points, each of the former salts separately; empty points, each in the presence of a saturated solution of the other. Data: A, Takegami;, a, W. B. Lee and A. C. Egerton, J. Chem. SOC., 1923,123,706; V, H. Precht and B. Wittgen, Chem. Ber., 1981,14, 1672; 0, G. Ron, Dead Sea Works Ltd, unpublished report, 1976. Curves are calculated. 1 .oo 0.75 2 0.50 E 0.25 0 1 I I 1 1 3.5 4 .O 4 . 5 5.0 mMgCI, Fig. 4. The solubilities of sylvite (a) and carnallite (b) as a function of the MgCI, molality in the vicinity of the carnallite point at 25 "C. Data: 0, Zdanovskii, quoted by Lehman14 and V d ~ v e n k o ; ~ ~ 0, D ' A ~ s ; ~ A, Osok~rova;~ 0, Serowy;12 V, Wood.ls3582 Solubilities in the System NaCl-KC1-MgC1,-CaC1,-H,O Table 4.Parameters for the solubility surface [eqn (18)] for quaternary system NaC1-KC1- MgC1,-H,O and the quinquinary system NaCl-KCl-MgC1,-CaC1,-H,O (at a constant molar ratio Ca: Mg = 1 : 3.5) at 25 "C halite saturationa sylvite saturationa no Ca with Ca no Ca with Ca ___ w:a 26.64 20.33 26.68 20.46 WE 26.35 11.42 26.32 11.38 bNa,, 0.0123 0.0114 0.0121 0.017 b,,, 0.0115 0.0064 0.0103 0.0059 'NaMg -1.236 1.001 -1.200 1.000 aKMR -1.202 0.199 -1.117 -0.146 aNaK -0.689 - -0.712 - aKNa -0.935 - -0.929 - 0.0098 - bNaK 0.0126 - 0.0146 - bKN, 0.0094 - cNaKMg 0.0114 - 0.0118 - cKNaMg 0.0238 - 0.0236 - a In each column the left-hand numbers concern saturation with the designated salt only, the right- hand numbers concern simultaneous saturation with both this and the other salt.KC I 20 40 60 80 Fig. 5. Janecke projection for the mutual precipitation of halite and sylvite at 25 "C from aqueous MgC1, solutions (full data points and solid calculated line) and from aqueous MgCl,-CaCI, solutions (Ca:Mg = 1 :3.5, empty data points and dashed calculated line). 0, Calculated carnallite precipitation point, the other data points from A, O~okorova;~ 0, 0, D'Ans;* T7, V, Oka;I5 0, Menczel" and ., Ron (see fig. 3 for reference). For the three salts k, i and j , the solubility of k is given in terms of the mass fractions (wt%) of i and j as (18) where wi is the solubility of the salt k in water, and the other five parameters pertain to ternary solutions [(a) and (b)] and to the quaternary solution (c).The values of these parameters are shown in table 4. W k = WE -k a k i Wi -k b k i W t -k a k j Wi -k bki Wj' -k Ckii W i WiY. Marcus and N . Sofler 3583 For the quinquinary system NaC1-KC1-MgC1,-CaC1,-H,O it is best to present the results in terms of a Jaenecke projection (superscript J) where the equivalent of the CaCl, present is added to that of MgCl, so that, e.g. and (the numbers are the molar masses). Such a projection is shown in fig. 5, where the calculated curve for the precipitation of both halite and sylvite, and the point beyond which (at higher wkgCl ) carnallite is also precipitated is shown for two cases. In one case CaCl, is absent, and in the other it is present at a constant molar ratio of 1 : 3.5 of CaCl, to MgCl,.The calculated curves are compared with data from the literature. For the latter, if CaCl, is present, ratios not too distant from 1 : 3.5 have been selected as far as possible. Discussion The novel feature of the present work is the binding together of vapour pressure data and multiple solubility data of multicomponent systems by a single set of equations. Vapour pressure data of ternary subsystems of the present quinquinary system have been presented previously by many authors, and need not be reviewed here. However, these studies concerned unsaturated solutions (sometimes rather concentrated ones), they employed mainly the isopiestic comparison method, and their main purpose was to obtain the activity coefficients and hence the interaction coefficients in these systems.In the present study emphasis has been placed on saturated solutions (fig. 1 and table 3) and on the quinquinary system (table 3), but the main feature is the representation of the vapour pressures of the ternary and higher systems by means of equations derived from the binary solutions and solubility data for the singular points of the ternary systems. Thermodynamics requires, of course, a definite relationship between the two phenomena. The present demonstration of its applicability, however, is in terms of the very simple equations (1HlO) and the minimal number of input data: PO, the relevant K, values and the compositions of the ternary singular points. The ability to predict (and to control) the vapour pressure of multicomponent systems is important in such applications as the precipitation of salt components from brines in evaporation basins.38 Translated into water activity values, eqn (9, this is significant in other kinds of processes, involving, e.g.selective membranes for the enrichment of desirable salt component^.^' The correction parameters B;,, eqn (14), are empirical quantities found to be able to describe the singular points of ternary phase diagrams by means of eqn (1HlO). They may, however, have a significance beyond this, since they permit the calculation of cation-cation interaction parameter^.^' For an amount of solution containing 1 kg of water, the excess Gibbs free energy of a ternary solution containing the salts i and j is given by GE/RT = [(zi + 1) mi + ( z j + 1) m,] (1 - q5) + (zi + 1) mi In y k i + (z, + 1) rn, In yki.(19) For the present level of approximation a symmetrical cation-cation interaction parameter gij is defined by4' GE/RT = xi(GFo/RT)+x,(G,"O/RT)+xix, 12gij (20) (21) where for each salt of our chloride mixture GEo/RT = ( z + 1) mo( 1 - q5' + In YO+) - the subscript O designating the pure salt at the (extrapolated, if necessary) ionic strength of the mixed solution. A combination of eqn (19H21) yields gij = ('/xi xj I ) [(xilzi) In (Y+~/YO+J + (xj/zj) In ( ~ + j / ~ " , ) l (22)3584 Solubilities in the System NaCl-KC1-MgCl,-CaC1,-H2O noting that rn = 2xI/z(z+ 1) for a salt in the mixture, and that the terms in the osmotic coefficients cancel out in view of eqn (10).Eqn (19)-(22) are valid (at the stated level of approximation) for any ternary solutions, both saturated and unsaturated. However, the singular point of the ternary system permits the evaluation of y+z and y k j in the mixed solution at this point from the imposition of the two solubirity product equations (23) where Ks, m and z take the subscripts i or j . The activity coefficient for the salt i in the mixed solution, where therefore becomes z + l S(l-1) Ks = mm& y + 'H,O K 1)-(4)1: mc1 = 21[x,/(z, + 1) + x,/(z, + I)] y k Z = K]-/(zz+l) 5, [z,(z, + 1)]1'(2z+1) (21)-l [x,/(z, + 1) (24) + x,/(z, + l)]-z&+l) akf$-l)/(2,+l) and similarly for the salt j . Eqn (22) may now be solved, using eqn (24) for y + with the I , x, and x3 data of the mixture at the singular point and the relevant K, valuesand with the extrapolated values of y", based on the empirical Bi, correction parameters.The resulting equations are, for halite and sylvite, gNaK = (2/xNaCl xKCl I ) [xNaCl In (@, h a l / y ~ NaCl xkaCl I ) +xKCl In (@,syl/y:Kcl xicl I)] = 0.0005 (at I = 7.33 mol kg-l) (25) for halite and bischofite, gNaMg = (2/xNaCl XMgC1, I ) (XNaC1 In {@, h a l / Y l N a C l XNaCl [(2 + xNaC,)/31t '1 + o.5xMgClz In [@, b l S / ~ ~ M g C l , xbgcl, x (3 - xMgcl2)~(I/3) a&,,]) = 0.146 (at I = 17.42 mol kg-') (26) and for sylvite and carnallite, g K M g = (2/xKCl XMgCl, I ) (xKCl In {@,, s y l / Y i KCl + xK,1)/3]' I > + 0.5xMgC12 In [@, C!d,/Y: MgCl, xLgc1,(3 - x M g c i j x (Z/3)a&20&,1]) = -0.0183 (at I = 12.80 mol kg-l).(27) Eqn (26) and (27) involve the water activities, since both bischofite and carnallite crystallize with six mol of water per mol of salt, and the relevant values are taken from fig. 1 or table 3 : a,,,(hal + bis) = 0.332 and aH,O(syl +car) = 0.587 (the former being insignificantly different from that of pure saturated aqueous bischofite). The values of the interaction parameters g,, mean that there are no significant interactions of Na+ and K+ (either repulsive or attractive) to form ion triplets Na+-Cl--K+, but there is fairly strong repulsion of Na+ and Mg2+, and fairly strong attraction of K+ and Mg2+ to form ion triplets K+-C1--Mg2+. The latter are probably a step in the direction of the formation of carnallite, and involve the hydration shells in some manner.Note that as the ionic strength of the mixed solution decreases, the attractive interactions of K+ and Mg2+ turn into repulsive ones. These points have been discussed previously. 38 It is interesting to note the small, but significant, differences between the quinquinary system and the quaternary subsystem without calcium shown by fig. 5. Replacement of some of the magnesium by calcium seems to ameliorate the severe shortage of 'free' water in the very concentrated solutions, because of a smaller overlap of the inner hydration spheres. This point, again, has been discussed before.38 E. Pross participated in the vapour pressure measurement. The Dead Sea Works Ltd are thanked for support of this study.Y. Marcus and N . Sager 3585 References 1 J.H. van’t Hoff, Sitzungsber. Preuss. Akad. Wissen., 1905, 252. 2 S. Takegami, Mem. Coll. Sci. Kyoto Imp. Univ., 1921, 4, 317. 3 N. A. Osokorova, M. A. Opikhtina, D. N. Shoikhet, E. F. Plaksina and A. I. Zaslavskii, Trans. State 4 J. D’Ans, Die Losungsgleichgewichte der Systeme der Salze Ozeanischer Salzablagernugen (Berlin, 5 I. Igelsrud and T. G. Thompson, J . Am. Chem. Soc., 1936, 58, 2004. 6 N. S. Kurnakov and A. V. Nikolaev, Izd. Akad. Nauk SSSR, Chem. Ser., 1938, 403. 7 0. K. Yanatieva, Zh. Obshch. Khim., 1947, 17, 1039. 8 W. J. Lightfoot and C. F. Prutton, J. Am. Chem. Soc., 1946, 68, 1001; 1948, 70, 4112; T. A. Meyer, W. J. Lightfoot and C. F. Prutton, J. Am. Chem. Soc., 1949, 71, 1236; W. J. Lightfood and C. F. Prutton, J. Am. Chem. Soc., 1949, 71, 1238.Inst. Appl. Chem. SSSR, 1932, 16, 24. 1933). 9 A. B. Zhdanovskii, Zh. Fiz. Khim., 1948, 22, 1486. 10 G. 0. Assarson, J. Am. Chem. Soc., 1950,72, 1437; J. Phys. Chem., 1953,57,717; G. A. Assarson and A. Balder, J. Phys. Chem., 1954, 58, 253, 416; 1955, 59, 613; 1956, 60, 1436. 11 A. P. Perova, Zh. Neorg. Khim., 1957, 2, 2789. 12 F. Serowy, C. Dahne, G. Doring and G. Malyska, Wissensch. Z., 1964, 6, 328. 13 Seidell’s Solubilities of Inorganic and Metal Organic Compounds, ed. W. F. Linke (Van Nostrand, 14 A. Lehrman, Geochim. Cosmochim. Acta, 1967, 31, 2309. 15 M. Motoyama, N. Kadota and S. Oka, Nippon Kaisui Gakkaishi, 1972, 26, 16; 173; K. Majima, K. Katsuki, M. Tejima and S. Oka, Nippon Kaisui Gakkaishi, 1972,26, 199; 205; 1973,27, 164; 315; 321 ; M. Motoyama, M.Kadota and S. Oka, Nippon Kaisui Gakkaishi, 1974, 28, 146; 327; 1975, 30, 26; S. Oka, M. Kadota and M. Motoyama, Nippon Kaisui Gakkaishi, 1976, 30, 142; M. Motoyama, M. Kadota and S. Oka, Toyo Daigaku Kogakubu Kenkyu Hokoku, 1975, 11, 15. Princeton, 4th edn, 1965). 16 J. R. Wood, Geochim. Cosmochim. Acta, 1975, 39, 1147; 1976, 40, 1211. 17 R. W. Potter and M. A. Clynne, J . Res. US. Geol. Survey, 1978, 6, 701. 18 E. Menczel, A. Apelblat, A. Roy and E. Korin, Rev. Chim. Miner., 1980, 17, 508. 19 Y. Marcus, Geochim. Cosmochim. Acta, 1977, 41, 1739. 20 Y. Marcus and N. Soffer, to be published. 21 G. Scatchard, J. Am. Chem. Soc., 1961, 83, 2636. 22 V. M. Vdovenko and M. A. Ryazanov, Radiokhimiya, 1965, 7, 39; 442. 23 R. M. Rush and J. S. Johnson, J. Phys. Chem., 1968, 72, 767.24 G. Scatchard, R. M. Rush and J. S. Johnson, J. Phys. Chem., 1973, 74, 3786. 25 P. J. Reilley, R. H. Wood and R. A. Robinson, J. Phys. Chem., 1971, 75, 1305. 26 K. S. Pitzer and J. J. Kim, J. Am. Chem. Soc., 1974, 96, 5701. 27 M. H. Lietzke and R. W. Stoughton, J. Solution Chem., 1972, 1, 299. 28 M. H. Lietzke and R. W. Stoughton, J. Inorg. Nucl. Chem., 1974, 36, 1315. 29 J. V. Leyendekkers, Anal. Chem., 1971,43,1835; J. Phys. Chem., 1971,75,946; J. V. Leyendekkers and M. Whitfield, J. Phys. Chem., 1971, 75, 957. 30 H. P. Meissner and J. W. Tester, Ind. Eng. Chem., Proc. Des. Deu., 1972, 11, 128; H. P. Meisner and C. L. Kusik, AIChE J., 1972, 18, 294; H. P. Meissner, C. L. Kusik and J. W. Tester, AIChE J., 1972, 18, 661; C. L. Kusik and H. P. Meissner, Ind. Eng. Chem., Proc. Des. Dev., 1973, 12, 112; H. P. Meissner and C. L. Kusik, Ind. Eng. Chem., Proc. Des. Dev., 1973, 12, 205. 31 K. S. Pitzer, J. Phys. Chem., 1973, 77, 268; K. S. Pitzer and G. Mayorga, J. Phys. Chem., 1973, 77, 2300. 32 W. J. Hamer and Y-C. Wu, J. Phys. Chem. Ref. Data, 1972, 1, 1047. 33 R. N. Goldberg and R. L. Nuttall, J. Phys. Chem. Ref. Data, 1978, 7, 263. 34 J. A. Rard and D. G. Miller, J. Chem. Eng. Data, 1981,26, 38; this is a fit with 8 parameters, including 35 B. R. Staples and R. L. Nuttall, NBS Tech. Note, 1976, 298. 36 J. Sangster, T. T. Teng and F. Lenzi, Can. J. Chem., 1973, 51, 2626; F. Lenzi, T. T. Tran and T. T. Teng, Can. J. Chem., 1975,53,3133; D. T . Au, T. T. Teng and J. M. Sangster, Can. J . Chem., 1978,56, 1853. half-integral values of k , and has not been employed here, for the sake of uniformity. 37 J. A. Riddick and W. B. Bunger, Organic Solvents (Wiley-Interscience, New York, 1970). 38 Y. Marcus, in Ionic Liquids, ed. D. Inman and D. G. Lovering (Plenum, London, 1981), pp. 97-115. 39 Y. Marcus and T. Nakashima, J. Phys. Chem., 1983, 87, 794. 40 H. L. Friedman, J. Chem. Phys., 1960, 32, 1351. Paper 8/00063H; Received 5th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403575

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1988, 84(10), 3587-3596 Excess Enthalpies of in the Supercritical Region Christopher J. Worrnald,* Nabil Al-Bizreh and Trevor K. Yerlett Department of Physical Chemistry, University of Bristol, Bristol BS8 1 TS Excess molar enthalpies HE of acetone-n-hexane mixtures have been measured at the critical temperature 493.2 K of the x = 0.5 mixture, and at the supercritical temperatures 510.2, 523.2 and 548.2 K. The HZ measure- ments for {x(CH,),CO + (1 - x)C,HI4} were made at fixed mole fractions x at pressures up to 7.94 MPa. At 510.2 K measurements were made at x = 0.25, x = 0.50 and x = 0.75, but at all other temperatures the measure- ments were made at x = 0.50 only. The HE values measured at 493.2 K have been combined with the enthalpies of acetone and n-hexane to obtain the enthalpy us.pressure isotherm at the critical temperature T,(x = 0.5) of the mixture. The HZ values in the supercritical region exhibit double maxima similar to those observed for carbon dioxide-ethane mixtures at supercritical temperatures. The double maxima fade as the temperature increases. The shape of the H: against p graphs is well reproduced by the Patel-Teja equaiion of state using k,, = 0.14 in the combining rule a12 = (1 - kI2)(alla2,)~. Density-dependent local composition mixing rules were found not to improve the fit to the measurements. The Patel-Teja equation has been used to explore the shape of the HZ(x,p,T) surfaces in the supercritical region. The surfaces are similar to those for carbon dioxide-ethane mixtures. Two mixtures which are of interest because the critical temperatures of their components are close together are acetone-n-hexane and carbon dioxide-ethane. For acetone (T, = 508.1 K, pc = 4.70 MPa) and n-hexane (T, = 507.4 K, pc = 2.97 MPa) the difference between the T, values (0.7 K) is smaller than that (1.12 K) between the T, values for carbon dioxide (T, = 304.21 K, pc = 7.38 MPa) and ethane (T, = 305.33 K, p , = 4.87 MPa).Kay1 has measured the critical temperature TJx) and critical pressure p,(x) of {x(CH,),CO + (1 -x)C,Hl,} over a range of x. Ratzch and Strauch2 have measured T,(x) close to x = 0.5 and agree with Kay's value T,(x = 0.5) = 493k0.4 K, 14.7 K below the mean of the pure-component T, values. Kay's value of p,(x = 0.5) is 3.78 & 0.03 MPa.Coincidentally Tc(x = 0.5) for carbon dioxide-ethane mixture^^*^ is 14k0.5 K below the mean of the pure-component T, values. The ( p , T ) projection of the ( p , T, x) surface for acetone-n-hexane mixtures in the critical region is shown in fig. 1. The shape of the ( p , T ) projection is almost the same as that of carbon dioxide-ethane shown in fig. 1 of ref. (5). Like carbon dioxide-ethane, acetone-n- hexane has a positive azeotrope which lies above the vapour-pressure curves for the pure substances. We recently reported measurements of the enthalpy of n-hexane,, acetone' and of the (OS(CH,),CO + 0.5C6H,,) mixture.' The measurements were made at temperatures up to 573.2 K and at pressures up to 11.3 MPa, and supercritical isotherms at 510.2 K were measured for the mixture and both components.Excess molar enthalpies HE(x = 0.5) calculated from the measurements showed an unusual double maximum when plotted against the pressure.' We have also recently reported measurements of HZ for (0.5c0, +0.5C2H,) in the supercritical r e g i ~ n . ~ These measurements, made using a flow 35873588 HZ for Acetone-n-Hexane Mixtures 4 I I do0 500 I 1 TIK Fig. 1. Phase diagram of (x(CH,),CO + ( 1 - x)C,H,,} in the critical region : (-) the upper solid curve is the vapour pressure of acetone and the lower solid curve is the vapour pressure of n-hexane. The curves are joined by the ( p , T ) projection of the ( p , T , x ) locus. (---), dew and bubble lines. Vertical broken lines indicate temperatures at which H: was measured over a range of pressure.0, [T,(x),p,(x)J measurements of Kay.2 mixing calorimeters rather than a total enthalpy calorimeter, exhibit similar double maxima to those obtained for (OS(CH,),CO + 0.5C,H1,}. Uncertainties in the HZ values for this latter mixture depend upon the uncertainties of the pure component and mixture enthalpies. When these were combineds the uncertainty in the HE values was estimated to be > 0.3 kJ mol-', which in the region of the double maximum observed at 510.2 K and x = 0.5 is ca. 8%. The uncertainty in the HZ values for (0.5CO2+0.5C,H,) obtained using a flow mixing calorimeter was > 2%. To facilitate better comparison of the two mixtures we have now made direct measurements of HZ for (x(CH,),CO + (1 - x)C,HI4) using a high-temperature flow mixing calorimeter.Experimental The flow mixing calorimeter was the same as that used previously" to make measurements of H i for binary mixtures containing steam. Acetone and n-hexane were purified as before.6' ' At 298.15 K the density of the purified acetone was 784.5 1 kg m-3 (lit.," 784.40 kg m-3) and the density of the purified n-hexane was 655.05 kg mW3 (1it.,l2 654.84 kg rn-,). The impurity in both fluids was < 0.1 mol YO. The pure liquids, drawn from reservoirs or from calibrated burettes, were pumped into flash boilers. Flow rates, obtained by timing the rate of drainage of the burettes, were accurate to 1 YO. The liquids were flash-vapourised at 5 15 K, at which temperature decomposition of the acetone was negligible. Vapours were mixed in a calorimeter containing a reverse flow labyrinth with a heater in the centre at the point of mixing.Power, supplied to the heater to offset the endothermic mixing process, was adjusted until the temperature of the outflowing mixture was within 0.01 K of the temperature of the inflowing pure components. The fluidised alumina bath containing the calorimeter was controlled to 0.1 K. The steel bomb containing the calorimeter smoothed temperature fluc- tuations to c 0.01 K. Mixture leaving the calorimeter was condensed and collected in a receiver back pressured with nitrogen controlled to +_20 kPa. Mixture removed from the receiver was analysed by g.1.c. ; negligible amounts of decomposition products were detected. The overall uncertainty in the HZ measurements is 0.1 kJ mol-' except at peak maxima, where it is 0.2 kJ mol-l.C.J . Wormald, N . Al-Bizreh and T. K. Yerlett Table 1. Excess molar enthalpies, HE, of (x(CH,),CO + (1 - x)CBHI1) 3589 p/MPa H z / J mol-' p/MPa Hz/J mol-1 p/MPa HE/J rnol-' x = 0.25 2.00 2.50 2.80 2.89 2.96 3.05 3.10 3.17 3.25 3.31 3.38 3.58 3.8 1 3.94 4.07 4.24 4.41 4.72 4.83 4.94 5.14 5.34 5.58 5.86 x = 0.50 1.58 2.34 2.82 3.07 3.17 3.25 3.34 3.53 3.70 4.05 4.20 4.37 4.45 4.55 4.63 4.70 4.74 4.82 4.88 4.92 5.06 T = 510.2 K 632 1065 1594 1840 2316 3120 4495 6533 6842 683 1 666 1 6149 4998 4025 2522 1293 460 162 880 1480 1478 1153 1037 896 T = 510.2 K 553 1049 1693 2865 4393 520 1 5290 5405 528 1 4848 4574 4146 3942 3571 3215 2856 2649 3015 3305 3882 3795 5.19 5.63 5.96 6.40 6.96 7.78 x = 0.75 2.27 2.67 2.89 3.05 3.12 3.24 3.38 3.45 3.57 3.71 3.85 4.08 4.25 4.42 4.52 4.58 4.67 4.76 4.84 4.89 4.96 5.05 5.21 5.36 5.50 5.71 6.0 1 6.36 x = 0.50 1.83 2.23 2.33 2.38 2.48 2.54 2.62 2.79 2.91 3462 2957 2694 2504 2350 2197 T = 510.2 K 64 1 962 1329 1832 2838 3043 3129 3129 3 149 31 13 3017 300 1 2919 2880 2872 2880 2977 3534 489 1 5285 5276 4677 3668 3252 2952 2654 2196 1709 T = 493.2 K 844 1417 1648 6364 7130 7008 7062 6887 6740 3.1 1 3.21 3.38 3.47 3.69 3.72 3.73 3.8 1 3.83 3.86 3.95 4.13 4.48 4.86 5.3 1 5.91 x = 0.50 2.49 2.91 3.32 3.80 4.07 4.1 1 4.25 4.53 4.90 5.22 5.49 5.90 6.26 6.52 6.95 7.3 1 7.94 x = 0.50 2.7 1 3.30 4.29 4.57 5.00 5.74 6.00 6.38 7.25 7.91 643 1 6323 5953 5580 5243 5 142 4916 7204 5855 5171 3932 3183 2742 2555 2442 2137 T = 523.2 K 630 899 1340 2522 3667 3864 3973 3789 348 8 3195 3010 3176 3334 3387 3120 2853 2520 T = 548.2 K 822 1201 1755 2013 2540 2708 2702 2614 2486 23503590 HE for Acetone-n- Hexane Mixtures Results and Discussion Results of HE measurements at x = 0.25, 0.50 and 0.75 are listed in table 1.The uncertainty in x is kO.01. Experiments at fixed composition were made over a range of pressure at 493.2, 510.2, 523.2 and 548.2 K. These temperatures are marked by vertical broken lines in fig. 1. Results of measurements at 510.2 K and x = 0.25, 0.50 and 0.75 are plotted against pressure in fig. 2(a)-(c). Solid lines through the points were drawn with a flexicurve. The 6.8 kJ mol-' peak at 3.3 MPa and x = 0.25 shown in fig. 2(a) becomes a rounded shoulder of 3.2 kJ mol-' at x = 0.75 [fig.2(c)], while the 1.5 kJ mol-' peak at 5.1 MPa and x = 0.25 becomes a sharp 5.4 kJ mol-1 peak at x = 0.75. The height of the peaks at x = 0.5 are almost the mean values of the peak heights at x = 0.25 and x = 0.75. The maxima in H: are at pressures ca. 0.25 MPa greater than the critical pressures of the pure components. As the pressure increases from 2 to 6 MPa there is clearly a very strong composition dependence of HZ(x). This is shown in fig. 3(a)-(c), where HZ(x) graphs have been constructed at pressures corresponding to the two peak maxima (3.3 and 5.0 MPa), the pressure corresponding to the minimum between the two peaks (4.7 MPa), and at 4.2 MPa, at which pressure the HE(x) curve is almost parabolic. The solid lines in fig. 3(a)-(c) were obtained by plotting H z ( x ) / 4 x ( I-x) and drawing the best straight line through the three points.The results of HZ(x = 0.5) measurements at 493.2, 523.2 and 548.2 K listed in table 1 are shown in fig. 4(a)-(c). Solid lines linking the experimental points were drawn with a flexicurve. At 493.2 K both components are below their critical temperatures. The two vertical parts of the curve are at the saturated vapour pressures of n-hexane (2.41 MPa) and acetone (3.81 MPa). Below 2.41 MPa both fluids enter the mixing calorimeter as vapours, and the first section of the HE curve corresponds to the change of enthalpy as gaseous mixture is formed. Just above 2.41 MPa n-hexane enters the calorimeter as a liquid, acetone enters as a vapour and a vapour mixture is formed. The height of the discontinuity at 2.41 MPa corresponds to the enthalpy of evaporation of n-hexane.It might be expected that above 2.41 MPa H z would continue to increase with increasing pressure. However, this effect is masked by a decrease in the enthalpy of evaporation of n-hexane as the density of the vapour phase increases. Above 3.81 MPa acetone enters the calorimeter in the liquid phase and there is a discontinuity in the HZ curve due to evaporation. Now the minimum in the critical locus of the mixture is at T,(x = 0.5) = 493 0.4 K, 0.2 K below the experimental temperature. The fluid mixture formed is just supercritical, and the density increases very rapidly with increasing pressure. HZ(x = 0.5) therefore falls sharply and continues to diminish with increasing pressure as the density of the mixture increases and approaches that of the pure components. At 523.2 and 548.2 K acetone, n-hexane and the mixture are supercritical under all conditions.The double maxima at 523.2 K are smaller than those at 510.2 K shown in fig. 3 (b), and are shifted to higher pressures. The maxima at 548.2 K are smaller still, and are again shifted to higher pressures. This shift to higher pressures was expected.13 The Critical Isotherm at 493.2 K Our HZ measurements at 493.2k0.1 K are so close to the critical temperature (493 k 0.4 K) of the mixture that they can be regarded as being at T,(x = 0.5). Enthalpy increments AHm of acetone and n-hexane have been measured as a function of pressure at 473.2, 490.2, 498.2 and 505.2 K.6*7 The increments are relative to 298.15 K and the saturated vapour pressure at this temperature.Interpolation of the enthalpy isotherms to 493.2 K is straightforward. The molar enthalpy of the mixture AH,(x) can be calculated from the equation AH,(x) = H:(x) + xAH,( 1) + (1 - X ) AHm(2) (1)C. J. Wormald, N. Al-Bizreh and T. K. Yerlett 359 1 2 0 2 4 6 P M a Fig. 2. Excess molar enthalpies, H:, of (x(CH,),CO + (1 - x)C,H,,) at 510.2 K plotted against pressure. (a) x = 0.25, (b) x = 0.50, (c) x = 0.75. 0, This work table 1 ; (-) the solid line through the points was drawn with a flexicurve; (---) the upper, middle and lower broken lines in (b) were calculated using k,, = 0.20, 0.14 and 0.10, respectively. The broken lines in (a) and (c) were calculated using k,, = 0.14. k,, is the adjustable parameter in eqn (7) used to fit the Patel-Teja equation of state to the measurements. where x = 0.5, and (1) and (2) refer to acetone and n-hexane, respectively. Enthalpy increments calculated from eqn (1) are listed in table 2.At 493.2 K the ideal-gas enthalpy of acetoneI4 is 49.37 kJ mol-l, while that of n-hexane12 is 66.61 kJ mol-I. Subtraction of the mean ideal-gas enthalpy from AH,(x = 0.5) yields the residual enthalpy H:(x) of the mixture. The critical residual enthalpy isotherm for the mixture is shown as the solid curve in fig. 5. Also shown in the figure are the residual enthalpy isotherms of acetone and n-hexane at 493.2 K. The close proximity of the solid curve to the residual enthalpy of acetone accounts for the sharp spike on the H:(x = 0.5) curve at 3.81 MPa shown in fig.4. The uncertainty of the table 2 enthalpy increments is estimated to be + 1 YO.3592 HE for Acetone-n-Hexane Mixtures X I I I I 1 6 -0 0.5 1 Y A Fig. 3. Excess molar enthalpies, H: of (x(CH,),CO + (1 - x)C,H,,) at 510.2 K plotted against mole fraction x of acetone. (-) Solid lines drawn through the experimental points as described in the text; (---) calculated from the Patel-Teja equation of state using k,, = 0.14; ( 0 * .) calculated from the Patel-Teja equation of state using the density-dependent local composition combining rules of Matthias and Copeman.16 (a) 3.3 MPa; (b) the parabolic curve is at 4.2 MPa and the skewed curve is at 4.7 MPa, where there is a valley between the two peaks shown in fig. 2(b); ( c ) 5.0 MPa.Comparison with the Patel-Teja Equation of State Comparison of the saturated liquid and vapour densities for acetone with four recent cubic equations of state' showed that of those tested, the Patel-Teja (PT) equation15 was superior. The equation fits acetone enthalpy increments to within 2-3% except in the critical region, where the fit is to within 5 % . The fit to n-hexane enthalpy increments6 is similar. The residual molar enthalpy, H z , for the one-fluid model may be written as where H: = p V- R T - (a - Tda/d T ) F( V ) (2) F( V ) = fN-l In[( V+ Q ) / ( V+ M ) ] (3) and N = [bc + 0.25(6 + c)'$ (4) M = $(b+c)-N ( 5 ) Q = i(b + C ) + N . (6) (7) a,, = (1 - U ( a 1 1 a22k To calculate cross-terms the arithmetic mean rule was used to obtain 6,, and c,,, and a,, was calculated fromC.J . Wormald, N . Al-Bizreh and T. K. Yerlett I I I I I 3593 4 8 p/Wa Fig. 4. Excess molar enthalpies of {x(CH,),CO + (1 - x)C,H,,) at T = (a) 493.2, (b) 523.2 and ( c ) 548.2 K. 0, This work, table 1 ; (-) solid lines were drawn with a flexicurve; (---) calculated from the Patel-Teja equation of state using k,, = 0.14; (. * . .) calculated from the Patel-Teja equation of state using the density dependent local composition combining rules of Matthias and Copeman. l6 Table 2 Enthalpies of fO.S(CH,),CO+0.5 C,H,,}, at 493.2 K" p/MPa AHJJ mol-' p/MPa AH,/J mol-' p/MPa AH,/J mol-I 0.1 57.8 3.7 49.5 4.2 41.6 1 .o 56.5 3.8 48.8 4.4 41 .O 2.0 54.8 3.9 43.6 4.8 40.5 3 .O 52.2 4.0 42.5 5.8 39.6 3.6 50.0 4.1 42.1 7.0 39.2 a The molar enthalpy increments, AH,, were calculated from the excess molar enthalpies, HZ, listed in table 1 together with the enthalpies of acetone and n-hexane listed in ref.(6) and (7).3594 Hz for Acetone-n- Hexane Mixtures plMPa Fig. 5 The residual molar enthalpy HZ of {0.5(CH3),CO+0.5 C,H,,}, (CH3),C0 and C,H,, at 493.2 K. This temperature is the critical temperature of the x = 0.5 mixture. (-) Calculated from HE values and pure component enthalpies as described in the text; (---) HZ for acetone (upper curve) and n-hexane (lower curve). The mixture parameter a was calculated from (8) and similar expressions were used to obtain b and c for the mixture. As the PT equation does not fit the residual enthalpies of the pure fluids to better than 5 % in the critical region it was not expected that the fit to HE would be anything but approximate.Fig. 2(b) shows the fit to HE(p,x = 0.5) obtained using k,, = 0.10,0.14 and 0.20. Atp = 3.3 MPa the PT equation with k,, = 0.2 gives a peak which is below that obtained from experiment by 0.6 kJ mol-', ca. three times the size of the uncertainty in H i . The choice k,, = 0.14 fits the peak at 5.0 MPa almost exactly, but falls short of the peak at 3.3 MPa by ca. 1 kJ mol-'. It is interesting to see how well the PT equation fits the Hg(p, x ) measurements at x = 0.25 and x = 0.75. The broken lines shown in fig. 2(a) and (c) calculated using k,, = 0.14, are of similar shape to the experimental curves but fall short of fitting the peaks by ca. 1 kJ mol-l. The fit to the composition dependence is shown in fig.3(a) and (c). The fit to H,(p,x = 0.5) at 498.2, 523.2 and 548.2 K is shown in fig. 4 (a)-(c). Again the calculated curves are similar to those obtained from experiment. Matthias and Copemanl' proposed density-dependent local composition (DDLC) mixing rules for the Peng-Robinson17 (PR) equation of state which greatly improve calculated liquid-vapour and liquid-liquid equilibria. At low densities the mixing rules reduce to the one-fluid model. To investigate whether these mixing rules would improve the fit in the supercritical region, and in particular to the curves shown in fig. 3(a) and (c), we used the truncated DDLC model together with the PT equation. Matthias and Copeman express the attractive molar Helmholtz energy Aatt of a mixture in the form where a is given by eqn (8), unc is a non-conformal attractive energy parameter and F( V ) is given by eqn ( 3 ) .The mixing rule for anc is where a; and u;l are obtained from criticality conditions and t,, and t,, are adjustable parameters. The non-conformal term vanishes when t,, = t,, = 0. The residual molar enthalpy HZ obtained using (9) is a = xf a,, + 2x,x, a,, + x i a,, Aatt = - aF( V ) - ancF( V),/2RT ( 9 ) anc = xf( 1 - x , ) ait,, + xi( 1 - x,) a;l t,, (10) H; = HZ(one.f,uid) - anc F( V),/ R T .C. J . Wormald, N . Al-Bizreh and T. K. Yerlett 3595 plMPa Fig. 6 The HE(x,p) surface for (x(CH,),CO + (1 - x)C6H14} at 510.2 K. (-) Calculated from the Patel-Teja equation of state using k,, = 0.14. Fig. T/K 7 The H;(x, T) surface for (x(CH,),CO + (1 - x)C,H,,} at 4.8 MPa.(-) Calculated from the Patel-Teja equation of state using k,, = 0.14. A range of values oft,, and t,, was tested. The dotted curves shown in fig. 3 and 4 were calculated using t,, = 0.046 and t,, = 0.285. The fit to the subcritical results at 498.2 K shown in fig. 4(a) is improved, but the fit to the results at 523.2 and 548.2 K shown in fig. 4(b) and (c) is worse. The fit to the skewed peaks shown in fig. 3(a) and (c) is not improved. The model seems to offer no improvement in the supercritical region. As the PT equation reproduces the shapes of the HE curves in the supercritical region reasonably well it is interesting to calculate HE over a wider range of x, T and p than was covered by the experiments. The HZ(p, T = 510.2 K, x) surface is shown in fig.6. Particularly interesting are the curves calculated at low mole fractions. As x-0 the peak sharpens and moves close to p,(n-hexane). At x = 0.1 the peak maximum is 4.8 kJ mol-1 at 3.25 MPa, and at x = 0.01 the maximum is 1.46 kJ mol-' at 3.12 MPa. At x = 0.9 the peak maximum is 3.95 kJ mol-1 at 5.0 MPa. At x = 0.99 there is a peak not shown in the figure similar to that at x = 0.01. The peak maximum is3596 HE for Acetone-n- Hexane Mixtures 1.19 kJ mol-l at 4.88 MPa, close top,(acetone). At temperatures above 510.2 K the peak heights are smaller, and the peak maxima are at higher pressures. Another particularly interesting surface is HE(p = 4.8 MPa, T, x). This surface, shown in fig. 7, is just 0.1 MPa above p,(acetone). At x = 0.99 the I mol% of n-hexane in the mixture produces a sharp narrow peak at a temperature ca.1 K above TJacetone). The peak rises to a maximum of 6.2 kJ mol-' at x = 0.9 and diminishes with decreasing x. The shoulder of a second peak emerges at x = 0.7 and is seen more clearly at x = 0.5. At x = 0.3 a double maximum appears. At x = 0.25 the two peaks are almost the same height, although this curve is not shown. As our measurements at x = 0.25 were made at a single temperature it was not possible to use our results to verify this double maximum. Below x = 0.25 the second peak becomes more pronounced. At x = 0.01 the peak is broad with a maxinium at 550 K, ca. 42 K above T,(n-hexane). At pressures > 4.8 MPa the peak heights are smaller, and the peak maxima are at higher temperatures.The HK(x,p) surface of {x(CH,),CO + (1 - x)C,H,,} at 5 10.2 K is similar to that of (xC0, +(1 -x)C,H,} at 308.40 K shown in fig. 5 of ref. ( 5 ) . For both mixtures the critical temperatures of the components are almost the same. The shape of the Ht(x,p, T ) surfaces depends on the relative position of the critical points on the phase diagram, and the amplitude of the maxima observed depends on how far apart the critical points of the components are. The values of HZ for carbon dioxide-ethane were best fitted using k,, = 0.1325, close to the value k,, = 0.14 chosen for acetone-n-hexane mixtures. Considering the very different nature of the molecules in the two mixtures these values of k,, are surprisingly close. As for carbon dioxidexthane mixtures, the double maxima in the acetone-n-hexane H:(x,p, T ) surfaces arise not from any unusual physical interactions in the fluid mixture, but from the irregular variations of the mole- fraction-weighted mean of the residual functions of the pure components. This is explained in ref. (5). References 1 W. B. Kay, J. Phys. Chem., 1964, 68, 827. 2 M. T. Ratzsch and G. Strauch, 2. Phys. Chem. (Leipzig), 1972, 249, 243. 3 N. E. Khazanova, L. S. Lesnevskaya and A. V. Zakharova, Khim. Prom., 1966, 4, 364. 4 K. Ohgaki and T. Katayama, Fluid Phase Equilibria, 1977, 1, 27. 5 C. J. Wormald and J. M. Eyears, J . Chem. Soc., Faraday Trans. 1, 1988, 84, 000. 6 C. J. Wormald and T. K. Yerlett, J. Chem. Thermodyn., 1985, 17, 1171. 7 T. K. Yerlett and C. J. Wormald, J . Chem. Thermodyn., 1986, 18, 371. 8 C. J. Wormald and T. K. Yerlett, J. Chem. Thermodyn., 1987, 19, 215. 9 C. J. Wormald and J. M. Eyears, J. Chem. Thermodyn., 1987, 19, 845. 10 C. J. Wormald and C. N. Colling, J. Chem. Thermodyn., 1983, 15, 725:. 11 A. L. Lydersen, Univ. Wisconsin Coll. Eng., Eng. Exp. Stn. Rep. 3, 1955. 12 A. P.I. Research Project No. 44: Selected Values of Properties of Hydrocarbon and Related Compounds 13 C. J. Wormald, Fluid Phase Equilibria, 1986, 28, 137. 14 J. Chao and B. J. Zwolinsky, J . Phys. Chem. ReJ Data, 1976, 5, 322. 15 N. C. Pate1 and A. S. Teja, Chem. Eng. Sci., 1982, 37, 463. 16 P. M. Matthias and T. W. Copeman, Fluid Phase Equilibria, 1983, 13, 91. 17 D. Y. Peng and D. B. Robinson, Ind. Eng. Chem. (Fundam.), 1976, 15, 59. (Texas A and M University, 1976). Paper 8/00125A; Receiued 14th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403587

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1988, 84(10), 3597-3603 Theoretical Interpretation of the Heats of Immersion of Lower n-Alcohols on Faujasites and Pentads Wolf-Dietrich Einicke, Ulf Messow, Rolf Schollner and Gernot Zahn Department of Chemistry, Karl Marx University, Liebigstr. 18, Leipzig 7010, German Democratic Republic On the basis of calorimetrically measured heats of immersion of lower alcohols on faujasite and pentasil zeolites, interaction energies between the zeolites and the adsorbed alcohols have been calculated. It is shown that the adsorption complexes of Na+-adsorbed molecules are the same for both families of zeolites. Several attempts have been made to calculate adsorbent-adsorbate interactions on the basis of the heats of immersion on Gal et a1.l interpreted the heats of immersion in terms of electric-field and London-type interactions. In this study we use the Gal model to interpret the heats of immersion of lower n-alcohols on several zeolites with varying silicon-to-aluminium ratios. Theory Gal et al.proposed a model in which all interactions contributing to the heat of immersion AHi can be described by average, pairwise and mutually independent interaction energies (1) where #Na and #o are the interaction energies of the sodium ion and the oxygen ion of the zeolitic lattice with the adsorbed molecule, and nNa and no are the numbers of interactions. All adsorbate-adsorbate interactions are described by #sp and q5' refers to the adsorbate zero energy. AhL is the heat of condensation. The last three terms of eqn (1) are independent of the zeolite and can be combined into the value AE AHi = nNa #Na 4- no $0 4- n$hSp i- n#' i- nAh, The interaction energy #Na can be resolved into five terms: (a) the disperion term -ANa/r6, where r is the effective distance between the interacting sodium ion and the adsorbed molecule.AN, can be calculated according to where El and E, are the ionization energies of the adsorbed molecule and the sodium ion and ai are the corresponding polarizabilities; (b) the repulsion term BNa/r12, where B,, can be calculated from equilibrium conditions; (c) the interaction of the field F of the sodium ion with the induced dipole moment of the adsorbed molecule (a,F2)/(8n&,), where E, is the electric permittivity of a vacuum; ( d ) the interaction of the field F o f the sodium ion with the permanent dipole moment ,u of the adsorbed molecule (@)/ (471~~); (e) the interaction of the field gradient of the sodium ion with the quadrupole moment, which can be neglected in the case of alcohol adsorption. 118 3597 FAR I3598 Heats of Immersion in Alcohol-Zeolites The interaction energy, $Na, per adsorbed molecule can be calculated according to (4) when the field, F, is defined and the values of B N a and r are derived.Concerning the field F, Gal et al. made the following assumptions. The dipoles -O-Na+ are replaced by point dipoles located in the middle of the dipole length 1. The field of the dipole is defined as - 2el ( 5 ) s3 (r+0.5 where e is the elementary charge and s is the electrostatic interaction distance related to the London-type interaction distance, r, by (6) s = r+0.5 1.Furthermore, only the interactions with the neighbouring adsorbed molecule were taken into account. Long-range interactions were neglected. In the case of X- and Y-zeolites, the adsorbed molecules can only interact with the sodium ions located in the large cages. Therefore the cation-distribution model of Mortier et a1.6 was used. The differentiation of eqn (4) leads to the following equation which allows the minimization $Na with respect to r, and the calculation of BN, F = - = 2P where rNa is the equilibrium distance of r. Eqn (5)-(7) give 0.25rNa (0.5 - ANa e2Fa2 $ =--- 2rka m0(rN,+0.5 I)6 rNa+0.5 I Na - ( 1 . 0 - nco(rNa + 0.5 I)3 The interaction energy of the second term $o is defined as A0 $0 = (9) but ro is not available.Because the third term is also not calculable, Gal et al. tested the model indirectly. They assumed that the limiting values of adsorption are the same for all zeolites of a family. The values of ethanol adsorption on our zeolites are shown in table 1 . It follows that the terms no g50 and AE are the same for one zeolite family. Eqn (2) can therefore be applied to two different zeolites and a substraction of the two equations leads to AHi(zeolite 1) - AHi(zeolite 2) = CnNa(zeolite 1) - nNa(zeolite 2)1 #Na' This means, in our case, the variation in the heats of immersion is a result of the different numbers of interactions, but not of different average energy g5Na. Experimental The chemical composition of the zeolites is given in table 1.The pentasil zeolites from template synthesis were activated for 8 h in a stream of air at 873 K before use to remove all organic material from the zeolite.W-D. Einicke et al. 3599 Table 1. Characterization of the molecular sieves investigated mg EtOH mol. sieve Si : A1 chemical composition (ao/bo/co)/A g-1 zeolite silicalite - NaZSM-5 50 NaZSM-5 19 NaZSM-5 13.5 NaY 5 NaY 2.6 NaX 1.4 20.074 (1 1) 'I 121.4 19.869 (8) 13.359 (14) ,I 19.883 (5) 124.4 13.358 (13) J (si02)96. 0 Na1.9(A102)1.9(Si02)94. 1 20.101 (5) 1 13.389 (15)) Na6.6(A102)6.6(si02)89.4 20.128 (5) 1 19.913 (1 1) 132.3 13.396 (13)) Na32(A10,)32(Si0,)160 - 214.5 - 232.8 Na80(A102)8,(Si0,)11, - 255.3 Na,3(A102),3(Si02)139 Table 2. Physical parameters and basic data of the alcohols used dipole moment8 polarizability" Ei8 4, / m3 /1O-l8 J /lo-'* m6 J liquid molecule-I molecule-l molecule-' molecule-' C m methanol 5.478 8.232 1.74 1.618 ethanol 5.544 12.917 1.68 2.452 n-propanol 5.577 17.532 1.62 3.208 n-butanol 5.610 22.154 1.60 4.004 n-pentanol 5.643 26.832 1.59 4.818 - Na+ in zeolitesg - 0.190 7.57 x 10-'O " Calculated using the refractive index given in ref.(8). The characterization of the zeolites was carried out by X-ray diffraction using Ni- filtered Cu, radiation over a 29 range of 7 to 44" at a scanning rate of 0.5" min-l. The zeolites were degassed in a batch vessel at 693 K and a pressure of less than 0.01 Pa. The heats of immersion of the alcohols were measured calorimetrically at 30 "C by means of an isothermal LKB sorption microcalorimeter.More detailed information about the heat measurements is given in ref. (7). Physical parameters and basic data of the alcohols used for theoretical considerations are given in table 2. Results and Discussion A comparison of the X-ray diffraction patterns with data given in the literature,1° the lattice parameters and the uptake of ethanol of the samples led us to the conclusion that the crystallinity of the molecular sieves is quite high. The heats of immersion of the lower alcohols on pentads and faujasites are demonstrated in fig. 1 and 2 which shows the dependence of the zeolites on the A1 : A1 + Si ratio. In both families of molecular sieves this dependence is linear, which enables us to interpret the experimental heats of 118-23600 Heats of Immersion in Alcohol-Zeolites 9 0.8 0- " I M 2 70. 2- 6 0. 5 0. 2 1 2 3 4 chain length Fig. 1. Heats of immersion of n-alcohols with pentasils showing dependence on the alcohol chain length: a, silicalite; A, NaZSM-5 (50); W, NaZSM-5 (19); 0, NaZSM-5 (13.5). 0.1 a2 a3 A1 : (A1 + Si) Fig. 2. Heats of immersion of n-alcohols on faujasites showing dependence on A1 : (A1 + Si) ratio : 0, methanol; a, ethanol; 0, propanol; a, butanol. immersion by means of the model mentioned above. Such behaviour was also reported by Tsutsomi and Takahashi5 for the immersion of polar molecules on the zeolites NaY and Cay, and by Hagiwara et aZ." for the heats of immersion of ethanol and butanol on carbon black, which was dependent on the concentration of the active H-centres on the surface.Other properties such as hydrophobicity, catalytic activity or methanol adsorption also vary linearly with the A1 content of the pentasil samples as shown by Olson'2 and Nakam0t0.l~ The values of the heats of immersion calculated from eqn (8)W-D. Einicke et al. 3601 I y. _---_ y ----_ y ---- -x---b-x 1 2 3 4 chain length Fig. 3. Contribution of different interaction energies to the heat of immersion and dependence of alcohol chain length on NaZSM-5 (19) : 1, dispersion interaction : lattice-molecule ; 2, field-dipole interaction ; 3, field-induced dipole interaction ; 4, dispersion interaction : Na+-molecule. 1 , . 0.1 02 0.3 A1 : (A1 + Si) Fig. 4. Contribution of different interaction energies to the heat of immersion of methanol and its dependence on the Al:(Al+Si) ratio of the faujasites.For the key see fig. 3. and (10) are connected by the dashed lines. The differences in the experimental and the calculated heats of immersion are < 5 J g-l in the case of faujasite zeolites and < 2 J g-l with the pentasil family. Taking into account an error in the estimation of the heats of immersion of 1.5 J g-l and the error in the Si:Al ratio determined by conventional chemical analysis, the values are in a good agreement. Fig. 3 shows the contribution of the different interaction energies to the heats of alcohol immersion on NaZSM-5 (Si:Al = 19), from which can be interpreted the increasing values with increasing carbon number by means of the higher dispersion3602 Heats of Immersion in Alcohol-Zeolites I I , 1 2 3 4 chain length Fig.5. Interaction distance rNa and dependence on alcohol chain length and zeolite family: 0, pentasils ; 0, faujasites. interaction of the oxygen wall with the alcohols (polarizability and ionization energy also increase). The plateau and the decrease of the heats of immersion at higher carbon numbers reported by Messowl* should be affected by the more complicated arrangement of the alcohol molecules in the channels of the pentasils. The strong dependence on the Al: Si ratio is a result of the significant change in the contribution of the non-specific and specific interactions. This is demonstrated for the immersion of faujasite zeolites in methanol in fig. 4. The non-specific : specific interaction ratio varies from 1.75 (Nay; Si:Al = 5) to NaX (Si:Al = 1.4) with 0.74.The contribution made by specific interactions to the heats of immersion of the alcohols are in good agreement with those calculated by Kiselev15 from the heats of adsorption at a coverage of one molecule per cavity on an NaX zeolite. Furthermore, the values for the contribution of the field-induced dipole interaction (ca. 12%) are the same as those reported by Stachl' for the part of the polarization to the heats of adsorption of n- alkanes on Nay- and Nay-zeolites. The best-fit parameter rNa of eqn (8), as regards dependence on the number of carbons and the kind of zeolitic family, is given in fig. 5 . It is interesting to note that the values of the interaction distance are nearly the same for each alcohol. It seems that the properties of the adsorption complex Na+-adsorbed alcohol molecule are independent of the nature of the zeolitic oxygen wall.With increasing chain length the interaction distance also increases. In the case of methanol adsorption a model based on the van der Waals radii in which a methanol molecule is in direct contact with the zeolitic sodium ion, indicates that the interaction distances calculated by means of eqn (8) are physically realistic. References 1 J. Gal, M. V. Radak and R. V. Hercigonja, Proc. 5th Int. Conf. Zeol., ed. L. V. C. Rees (Heyden, 2 R. V. Hercigonja, B. B. Radak and V. M. Radak, Proc. Int. Conf. Properties and Applications of 3 B. Coughlan and W. H. Carrol, J. Chem. Soc., Furuday Trans. I , 1976, 72, 2016. 4 K. Tsutsumi and H. Takahashi, ref. (I), p. 483. 5 K. Tsutsumi and H. Takahashi, J . Phys. Chem., 1972, 76, 710. London, 1980), p. 516. Zeolites, ed. R. P. Townsend (The Chemical Society, London, 1979), p. 198.W-D. Einicke et al. 3603 6 W. J. Mortier, H. J. Bosmans and J. B. Uytterhoeven, J. Phys. Chem., 1972, 76, 650. 7 H. Herden, W-D. Einicke, U. Messow, K. Quitzsch and R. Schollner, Chem. Tech., 1982, 34, 364. 8 Handbook of Chemistry and Physics (C.R.C. Press, Boca Raton, Florida, 64th edn, 1984). 9 Handbook of Chemistry and Physics (CRC Press, Boca Raton, Fla, 54th edn, 1973). 10 G. T. Kokotailo, S. L. Lawton, D. H. Olson and W. M. Meier, Nature (London), 1978, 272, 437. 11 S. Hagiwara, K. Tsutsumi and H. Takahashi, Carbon, 1981, 19, 107. 12 D. H. Olson, W. 0. Haag and R. M. Lago, J. Catal., 1980, 61, 390. 13 H. Nakamoto and H. Takahashi, Zeolites, 1982, 2, 67. 14 U. Messow, K. Quitzsch and H. Herden, Zeolites, 1984, 4, 255. 15 A. V. Kiselev, Molecular Sieve Zeolites II, A.C.S. Symp. Ser. 102, p. 37. 16 H. Stach, H. Thamm, K. Fiedler and W. Schirmer, Adsorption of Hydrocarbons in Zeolites, preprints of the Workshop (Berlin, 1979). Paper 8/00210J; Received 11th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403597

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J . Chem. Soc., Faruday Trans. I, 1988, 84(10), 3605-3613 Infrared Study of Ammonia-Carbon Monoxide Reactions on Silica-supported Iron Catalysts Colin Johnston, Norman Jorgensen and Colin H. Rochester* Chemistry Department, The University, Dundee DDl 4HN Infrared spectroscopy has been used to study the reactions of carbon monoxide and ammonia over Fe/SiO, catalysts at cu. 298-723 K. Adatoms of nitrogen resulting from the dissociative adsorption of ammonia reacted with CO to form surface isocyanate on iron at 298-523 K. At high temperatures spillover of isocyanate to the silica surface occurred. An accompanying decomposition reaction led to the appearance of intense infrared bands at 2120 and 2055 cm-l, which are ascribed to surface cyano complexes of iron. Hydrogen cyanide was adsorbed on iron at ca.298 K as Fe=C=NH. At 573 K complete oxidation of bulk iron occurred, giving bulk cyanide and surface isocyanate groups on the silica support. Subsequent standing at ca. 293 K in vacuum led to the formation of HCN polymerisation products. Infrared bands appearing in spectra of ammonia adsorbed on Fe/SiO, have been assigned to vibrations of NH(ad~),l-~ NH2(ads)24 and NH,(ads),'I5 all of which are believed to be intermediates in the catalytic synthesis or decomposition of ammonia.6 Four infrared bands, two due to linearly adsorbed CO and two ascribed to bridge- bonded CO, have been o b ~ e r v e d ~ * ~ for CO adsorbed on Fe/SiO, samples which in the presence of adsorbed ammonia gave infrared bands assigned to vibrations of N H , ( ~ ~ S ) .~ The present experiments involving competitive adsorption of CO and NH, and reactions of CO-NH, mixtures over Fe/SiO, were aimed at gaining further information about sites for the adsorption of ammonia on small iron particles. A possible product from the reactions was adsorbed isocyanate NCO(ad~),~-ll which has been reported as a product from the reaction of NO+CO over Fe/Si0,.5 Cyanide groups, CN(ads), were thought to be responsible for a band at 2060 cm-l in spectra of MgO after treatment with CO-NH, at high temperature,' and therefore HCN has been adsorbed on Fe/Si02 in order to give an indication of infrared band positions for CN groups on the surface of iron. Experimental Self-supporting discs of 5 wt % Fe/Si02 were prepared from iron(rr1) nitrate and aerosil silica (surface area 176 m2 g-') by reduction in hydrogen as before' with a final reduction temperature of 628 5 K.Research-grade CO (99.997 YO) was passed through a 77 K trap before use. Ammonia (99.98%) was purified by treatment with KOH and then Na before removal of permanent gases by a series of freeze-thaw cycles. Hydrogen cyanide gas in argon diluent was prepared from the reaction between KCN and concentrated sulphuric acid under argon. Spectra were recorded with a Perkin-Elmer 68 1 spectrometer in conjunction with a 3600 infrared data station. All spectra were recorded with discs at the ambient temperature (ca. 298 K) in the spectrometer compartment. 36053606 CO-NH, on Fe/SiO, Results Heat treatment of silica in CO-NH, mixtures only gave infrared bands due to SiNH, groups, which were also formed when silica was heated in ammonia a10ne.~ All other bands appearing when Fe/SiO, discs were heated in C S N H , mixtures may be attributed to the presence of iron in the catalyst.Spectra of CO on Fe/SiO, contained bands at 2040cm-l (high coverages) and 2020 cm-' (low coverages) due to linearly adsorbed CO and at 1970 and 1880 cm-l due to bridge-bonded C0.7*8 Subsequent addition of ammonia at ca. 293 K resulted in the disappearance of infrared bands due to linearly adsorbed CO and a slight increase in the intensities of the bands due to bridge-bonded CO. New maxima at 3380, 3290 and 1610 cm-l ascribed to NH,(ads), were also present in spectra of Fe/SiO, exposed to NH, without preadsorption of C0.5 Experiments in which non-dissociatively adsorbed (infrared-detectable CO) was completely removed by evacuation7 before the admission of ammonia showed that the intensities of the bands due to NH,(ads) were more intense after CO treatment of Fe/SiO, and subsequent adsorption of NH, than after NH, adsorption without preadsorption of CO.Dissociative adsorption of C08 generated surface species which were not removed by evacuation at ambient temperature and which promoted the adsorption of ammonia. The adsorption of ammonia on Fe/SiO, followed by evacuation at ambient temperature and admission of CO gave bands due to adsorbed CO which were weaker particularly for linearly adsorbed CO than those appearing if NH, was not preadsorbed. However, a new band due to surface isocyanate FeNcO(ad~)~ appeared at 2220 cm-'. This band also appeared if both NH, and CO together were left in contact with catalyst no matter whether the NH, or CO was adsorbed first.If NH, was added first, and not removed by evacuation before addition of CO, no bands due to CO(ads) were observed. The reaction of NH, and CO to form FeNCO(ads) was promoted by increasing temperature up to 523 K [fig. 1 (bHk)]. At higher temperatures spillover of NCO(ads) from the metal surface to the silica support12*13 led to the appearance of bands at 2320 [fig. 1 (Z)-(r)] and 1465 cm-l due to SiNCO(ads).14-16 The maximum due to FeNCO(ads) decreased in intensity with increasing temperature in the range 523-603 K [fig. 1 (Z)-(n)], an effect which was accompanied by the concomitant growth of a pair of intense maxima at 2120 and 2055 cm-l.The band due to SiNCO(ads) decreased in intensity in the higher temperature range 673-723 K [fig. l(q) and (r)]. A disc of Fe/SiO, was heated in a CO-NH, mixture and cooled to ambient temperature with the gases still present to give an infrared spectrum similar to that shown in Fig. 1 (Z). Subsequent removal of the gases by evacuation revealed bands due to NH,(ads) at 3380, 3290 and 1610cm-l which were otherwise obscured by intense maxima due to gas-phase ammonia. Species responsible for the maxima at 2320, 2220, 2120 and 2055 cm-' were also not desorbed by evacuation at ca. 293 K. A desorption experiment involving a disc on which only NH,(ads) and FeNCO(ads) had been generated [cf. fig. 1 (k)] showed that the intensity of the band at 2220 cm-l was ca.25, 85 and 97% reduced in intensity after evacuation at 400, 477 and 490 K, respectively. Surface FeNCO groups were also destroyed by the addition of water vapour at ca. 293 K. No infrared bands due to reaction products could be detected. The addition of CO at ca. 293 K to Fe/SiO, exposed to ammonia enhanced the intensities of the infrared bands due to NH,(ads). The subsequent formation of FeNCO(ads) was rapid. For example, the maximum intensity of the band at 2220 cm-l was attained in less than 1 min for Fe/SiO, at 500 K. No further change in spectrum occurred over a subsequent period of 25 min with the disc at 500 K. An Fe/SiO, disc which had been heated in NH,-CO at 723 K was left in the gaseous mixture for 16 h at 298 K.A weak residual band at 2375 cm-' was accompanied by an intense absorption envelope containing the maxima at 2120 and 2055 cm-' [fig. 2(a)]. AC. Johnston, N . Jorgensen and C. H. Rochester 2400 2000 4 0 0 2 000 I 1 goo 2000 3607 w avenumtm/m-' Fig. 1. Spectra of 5 % Fe/SiO, exposed to ammonia (10 kN m-2) and CO (10 kN m-2) mixtures (a) before addition of gases, (b)-(r) after consecutive heat treatments (15 min) in the gaseous mixture at temperatures of (b) 298, (c) 323, ( d ) 348, (e) 373, cf) 398, (g) 433, (h) 458, (i) 473, (1') 503, (k) 523, (I) 553, (m) 573, (n) 603, (0) 623, ( p ) 648, (4) 673 and ( r ) 723 K. band at 1655 cm-' disappeared on evacuation which also promoted the reappearance of a maximum at 2236 cm-l ascribed to FeNCO(ads) [fig. 2(d)].Transient species present during removal of surface species by evacuation gave maxima at 1735 cm-' [fig. 2(b)] and 2320 cm-' [fig. 2(c)] and the band at 2375 cm-' was briefly enhanced in intensity before its complete removal. A strong band at 1465 cm-l disappeared on evacuation and is probably attributable to a vibration of NCO(ads) groups." The intensities of the main bands displayed in fig. 1 are plotted in fig. 3 as a function of the temperature of Fe/SiO, during exposure to NH,-CO. The extent of formation of FeNCO(ads) was enhanced in the two temperature ranges 298-348 and 458-523 K [fig. l(a)], perhaps suggesting the involvement of two types of adsorption site. The absorbance data emphasize the correspondence between the loss of FeNCO(ads) and the appearance of the species responsible for the infrared bands at 2120 and 2055 cm-' [fig.3 (c)], the relative intensities of which remained constant within experimental error. These two bands weakened in the temperature range 623-673 K when no FeNCO(ads) remained on the iron surface. However, a further increase in intensities occurred at 723 K in parallel to a decrease in the concentration of SiNCO(ads) on the surface of the silica support [fig. 3 (b)]. The parallel between the loss of band intensity at 2320 cm-l and the growth of the maxima at 2120 and 2055 cm-9 is further reflected in the difference between fig. l(r) and fig. 2(a). The effects of hydrogen on the adsorbed products of reaction of CO-NH, over Fe/SiO, have been studied. A disc of Fe/SiO, was treated in order that its spectrum3608 CO-NH, on Fe/SiO, 'r ' I I I 1 2400 2000 1800 1600 1400 wavenumber/cm-' Fig.2. (a) Fe/SiO, [after fig. 1 (r) was recorded] left in the CO-NH, mixtures (298 K, 16 h) and (b)-(d) subsequently evacuated (298 K) for (b) CQ. 1 min, (c) ca. 2 min and (d) to a pressure of 2 x lo-, N m-,. TIK Fig. 3. Intensities of bands at (a) 2220, (b) 2320 and (c) 2120 and 2055 cm-' as a function of the temperature of Fe/SiO, in a CO-NH, mixture.3609 C. Johnston, Iv. Jorgensen and C. H. Rochester 10% I - 2 3 0 0 2 0 0 0 - 2 3 0 0 200C w avenumber/cm-' Fig. 4. Spectra of Fe/SiO, after (a) reduction, (b) heating in NH, (10 kN m-2wO(10 kN m-,) at 500 K (1 h) and 600 K (1 h) and evacuation (300 K, 30 min), (c) addition of hydrogen (101 kN m-,, 300 K) and heating (15 min) at ( d ) 400, (e) 500 and (f) 600 K.1 I I I 2300 2000 1800 1600 wavenumber/cm-' Fig. 5. Spectra of Fe/SiO, as for (a) fig. 4(b), (b) fig. 4(c) and (c)-(e) after heating in hydrogen (101 kN mP2, 473 K) for ( c ) 15, ( d ) 30 and (e) 60 min.3610 CO-NH, on Fe/SiO, I I I I 3 500 3000 2400 2100 wavenumber/cm-' I 1700 1600 Fig. 6. Spectra of Fe/SiO, (a) reduced disc, (b)-(h) after contact with HCN at (b) 303 K, ( c ) 573 K ( 5 min), ( d ) 573 K (25 min), (e) 573 K (55 min), (f) 573 K (1 15 min), ( g ) 673 K (15 min) and (h) 673 K (35 min), (i)-(k) after evacuation for (i) 5 min, (I] 1 h and ( k ) 5 days. contained maxima at 2320, 2220, 2120 and 2055 cm-' [fig. 4(b)]. The addition of hydrogen at 300 K caused a large increase in the intensity of the band at 2220 cm-' due to FeNCO(ads) [fig.4(c)]. Subsequent heat treatment at 400-600 K fig. 4 ( + 0 reversed this effect and led to the removal of the species responsible for all the infrared bands except that at 2320cm-' due to SiNCO(ads). An experiment in which spectra of Fe/SiO, in hydrogen at 473 K were recorded after various times gave similar results and established that the removal of the surface species was slow (fig. 5). A band at 1655 cm-l (cf. fig. 2) disappeared when the disc was heated in hydrogen [fig. 5(b)-(d)], although a species giving a weak band at 1570 cm-l was more resistant to decomposition. Spectra of HCN in contact with Fe/SiO, at 303 K contained weak bands at 3340 and 3280 cm-' due to the gas phase and a maximum at 1685 cm-' ascribed to an adsorbed species [fig.6(b)]. Subsequent treatments at 573 and 673 K produced time-dependent effects [fig. 6(c)-(h)], which resulted in the disappearance of the band at 1685 cm-' and the appearance of three bands at 2362, 2320 and 2225 cm-'. The latter appeared first at 573 K followed by the maximum at 2320 cm-l, both bands growing in intensity with time. A general increase in absorption intensity in the range 300&3700 cm-' suggested that dissociative adsorption of HCN was generating surface hydroxyl groups on the silica support. The band at 2362 cm-' became distinct when the temperature was raised to 673 K, which also resulted in a weakening of the bands at 2225 and 2320 em-'. However, brief evacuation restored the 2225 cm-l band to more than its original intensity whilst removing the band at 2362 cm-' [fig.6(i)]. Further evacuation for 1 h had little additional effect [fig. 601, but after 5 days new spectral features includedC. Johnston, N . Jorgensen and C. H. Rochester 361 1 bands at 35 15, 3408 (NH stretching vibrations), 297 1 (CH stretching vibrations), 1671, 1627 and 1468 cm-l [fig. 6(k)]. The band at 2225 cm-l also shifted to lower wavenumbers. Heating an Fe/SiO, disc in HCN at 573 K was accompanied (ca. 5 min) by a colour change from black to yellow. Discussion The band at 2220 cm-l, ascribed to a vibration of NCO(ads) on iron, corresponds to similar bands in the range 2200-2270 cm-l which have been attributed to NCO(ads) resulting from the reaction of CO and NH, over MgO,' MgO-Coo solid solutions1' and supported rhodium.1° Isocyanate complexes of iron typically give an infrared band around 2180-2216 cm-l.17 The formation of NCO(ads) on iron may involve reaction between CO and N adatoms, which are generated by the dissociative adsorption of NH, at 290 KlS and are more readily formed at higher temperature^.'^ The appearance of FeNCO(ads) after exposure of Fe/SiO, to ammonia, evacuation at ca.293 K and admission of CO shows that the reactive nitrogen-containing species was retained on the iron after evacuation. The only infrared-detectable bands after evacuation were at 3380, 3290 and 1610 cm-l, and these have been ascribed to N H , ( ~ ~ s ) . ~ The bands underwent no detectable change in intensity when CO was admitted, suggesting that NH-containing species were not directly involved in NCO(ads) formation.The assignment of infrared bands at 2320 and 1465 cm-l to SiNCO(ads) is consistent with the similar attribution of bands at 2300 and 1460 cm-l in spectra of Rh/SiO, heated in CO-NO mixtures.13 Spillover of NCO(ads) from the transition-metal surface to the oxide surface in oxide-supported metal catalysts is well established. l3 Dissociative adsorption of isocyanic acid20 and ethyl isocyanate21 on silica gave infrared bands at 23 13 and 2308 cm-', respectively, due to SiNCO(ads). The high-temperature stability of SiNCO(ads) either in CO-NH, mixtures (fig. 3) or in hydrogen (fig. 4 and 5 ) is consistent with the previously noted stability of this species in Rh/SiO, ~ata1ysts.l~ Eley et aL2' observed that SiNCO(ads) formed from the reaction between ethyl isocyanate and silica was not destroyed by evacuation at 973 K and was only partially desorbed even at 1118 K.The promoting effect of dissociatively adsorbed CO on the subsequent associative adsorption of NH, may be attributed to the capacity of oxygen adatoms to impart cationic character on adjacent surface iron atoms, which therefore became strengthened as potential sites for interaction with lone-pair electrons in ammonia molecules. The results of the studies of the competitive adsorption of CO and NH, at ca. 293 K suggests that there were at least two forms of NH,(ads) on iron, one strongly adsorbed (bands at 3380,3290 and 1610 cm-l after evacuation) and a weakly adsorbed species for which the infrared bands were obscured by bands due to gas-phase ammonia.The weakly held species was desorbed by evacuation at room temperature and occupied sites which were available in the absence of ammonia gas for the adsorption of CO in the bridge bonded form. In the presence of ammonia gas the subsequent admission of CO failed to generate the bridged CO species on iron, although bands characteristic of bridged CO did appear if ammonia was desorbed by evacuation before the admission of co. Linearly adsorbed CO occupied sites which were also available for the strong adsorption of ammonia. Strongly adsorbed ammonia displaced linear CO from the surface. Dissociative adsorption of ammonia to give N adatoms probably occurred at sites on which CO could be linearly adsorbed in the absence of ammonia.The N adatoms were the precursors of adsorbed NCO groups formed by reaction with CO. The intense bands at 2120 and 2055 cm-l cannot be unambiguously assigned. Carbonyl complexes of iron typically give infrared bands in this region,,, although this is an unlikely attribution because of the absence of bands below 2000 cm-l (due to bridged3612 CO-NH, on Fe/SiO, carbonyl groups) the high-temperature stability of the species and the fact that the bands did not appear if Fe/SiO, discs were heated in CO alone. Nitrogen and azido complexes of iron also give bands in the same spectral regi~n.,~-,~ However, the failure to generate the bands by heat treatment of Fe/SiO, in nitrogen, nitrous oxide, ammonia or nitrous oxide-ammonia mixtures suggests that the bands were not due to N, or N; species.The most likely explanation of the bands would be that they were due to cyano complexes formed at the iron surface. The two bands appeared at positions typical of those reported for cyano complexes of iron.27 By analogy with the present assignment, a band at 2060 cm-' in infrared emission spectra of magnesium oxide which had been heated in CO-ammonia mixtures was attributed to cyanide ions.9 Borello et d9 have given detailed mechanisms for CO-ammonia reactions leading to adsorbed formamide ions, formate ions, isocyanate derivatives and isocyanate ions on magnesium oxide. The appearance of infrared bands at 1735, 1655, 1600 and 1570cm-l in the present study may also be attributed to formamide (1735 cm-1)15728 or similar species to those formed on MgO.The species responsible for the infrared band at 2375 cm-l is more likely to be a silyl isocyanate than a silyl cyanide,14-16 its existence on the silica rather than the iron surface being confirmed by the absence of the band in spectra of alumina-supported iron which had been heated in CO-ammonia mixtures.,' The formation of surface isocyanate on silica at high temperatures possibly14~ l6 involves reaction between surface oxygen atoms on silica and cyano species generated at the iron or silica surface. In contrast the SiNCO species responsible for the infrared band at 2320 cm-' contains oxygen atoms derived from CO molecules. Surface species on iron after heat treatment in CO-ammonia mixtures include cyano and isocyanate groups, carbon and oxygen atoms derived from the dissociative adsorption of CO and N adatoms derived from ammonia.Hydrogen is dissociatively adsorbed on iron at ambient temperature.6 The enhancement in the band due to FeNCO(ads) when discs treated with CO-ammonia and evacuated were exposed to hydrogen (fig. 4) must arise because of displacement and reaction of surface species induced by the adsorption of hydrogen. The absence of other appreciable changes in the infrared spectra suggest that the species involved were infrared inactive in the spectral region studied. A tentative proposal would be that hydrogen was adsorbed at sites occupied by dissociatively adsorbed CO which was thereby displaced to an adjacent N adatom to form NCO. The formation of NCO(ads) from N(ads) and CO(g) may be unfavourable for N adatoms at iron sites adjacent to sites occupied by 0 adatoms derived from dissociatively adsorbed CO.However, conversion of N(ads) to NCO(ads) may become favourable when the adjacent sites are occupied by H adatoms. The infrared band at 1685 cm-l in spectra of HCN on Fe/SiO, is assigned to the C=N stretching vibration of an Fe=C=NH surface complex on iron. A similar structure has been proposed for HCN adsorbed on nickel and tungsten.,O Subsequent heat treatment of Fe/SiO, in HCN did not help in the assignment of the bands at 2120 and 2055 cm-' for Fe/SiO, heated in CO-ammonia mixtures. The rapid colour change of Fe/SiO, from black to yellow showed that oxidation of iron in the bulk phase had gone to completion. The band at 2225 cm-l may be assigned to C=N stretching vibrations of bulk cyanide.Bands at 2320 and 2362 cm-l may be ascribed to isocyanate probably on the silica s u r f a ~ e . ~ ~ ? ~ ~ The spectrum of Fe/SiO, which had been treated with HCN at high temperature and left in vacuum at room temperature for five days showed that polymerisation of HCN had o c ~ u r r e d . ~ ~ * ~ ~ One possible product is the HCN tetramer diamin~maleionitrile,~~~ 31-33 although the band at 297 1 cm-l shows that a product containing a CH group was also formed. We thank the S.E.R.C. for two CASE studentships, ICI Agricultural Division for collaboration and financial support, and Drs J. R. Jennings and S. A. Topham for helpful discussions.C. Johnston, N. Jorgensen and C. H. Rochester 3613 References 1 T.Nakata and S. Matsushita, J. Phys. Chem., 1968, 72, 458. 2 D. V. Pozdnyakov and V. N. Filimonov, Kinet. Catal., 1972, 13, 475. 3 T. Nakata, J. Chem. Phys., 1982, 76, 6328. 4 T. Nakata and S. Matsushita, J. Chem. Phys., 1982, 76, 6335. 5 C. Johnston, N. Jorgensen and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1988, 84, 2001. 6 M. Grunze, in Chemical Physics of Solid Surfaces and Heterogeneous Catalysis, ed. D. A. King and 7 N. Jorgensen and C. H. Rochester, Appl. Catal., 1986, 25, 69. 8 C. Johnston, N. Jorgensen and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1988, 84, 309. 9 E. Borello, S. Coluccia and A. Zecchina, J. Catal., 1985, 93, 331. D. P. Woodruff (Elsevier, Amsterdam, 1982), p. 143. 10 D. Gutschik and H. Miessner, React. Kinet. Catal.Lett., 1983, 22, 221. 11 A. Zecchina and G. Spoto, J. Catal., 1985, %, 586. 12 J. Rasko and F. Solymosi, J. Chem. SOC., Faraday Trans. I , 1980, 76, 2383. 13 W. C. Hecker and A. T. Bell, J. Catal., 1984, 85, 389. 14 F. E. Ruttenberg and M. J. D. Low, J. Am. Ceram. SOC., 1973, 56, 241. 15 D. D. Eley, G. M. Kiwanuka and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1973, 69, 2062. 16 B. A. Morrow and I. A. Cody, J. Chem. SOC., Faraday Trans. I . 1975,71, 1021. 17 R. A. Bailey, S. L. Kozak, T. W. Michelson and W. N. Mills, Coord. Chem. Rev., 1971, 6, 407. 18 K. Kishi and M. W. Roberts, Surf. Sci., 1977, 62, 252. 19 M. Weiss, G. Ertl and F. Nitschke, Appl. Surf. Sci., 1979, 2, 614. 20 F. Solymosi and T. Bansagi, J. Phys. Chem., 1979, 83, 552. 21 D. D. Eley, G. M. Kiwanuka and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1973, 69, 2062. 22 K. Nakamoto, Infrared Spectra of Inorganic and Coordination Compounds (Wiley, New York, 1963), 23 W. Beck, W. P. Fehlhammer, P. Pollmann, E. Schuierer and K. Feldl, Chem. Ber., 1967, 100, 2335. 24 A. Sacco and M. Aresta, Chem. Commun., 1968, 1223. 25 K. B. Yatsimirskii, V. V. Nemoshkalenko, Y. P. Nazarenko, V. G. Aleshin, V. V. Zhilinskaya and Y. D. Taldenko, Dokl. Phys. Chem., 1975, 217, 835. 26 V. V. Zhilinskaya, Y. P. Nazarenko, Y. I. Bratushko and K. B. Yatsimirskii, Russ. J. Inorg. Chem., 1974, 19, 1197. 27 E. F. Herrington and W. Kynaston, J . Chem. SOC., 1955, 3555. 28 J. C. Evans, J. Chem. Phys., 1954, 22, 1228. 29 C. Johnston, unpublished results. 30 J. R. Anderson and N. J. Clark, J. Catal., 1966, 6, 20. 31 M. J. D. Low, N. Ramasubramanian, P. Ramamurthy and A. V. Deo, J. Phys. Chem., 1968, 72, 32 M. J. D. Low and P. Ramamurthy, J. Res. Inst. Catal., Hokkaido Univ., 1968, 16, 535. 33 R. L. Webb, S. Frank and W. C. Schneider, J. Am. Chem. SOC., 1955, 77, 3491. pp. 177, 181. 2371. Paper 8/002781; Received 25th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403605

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1988, 84(10), 3615-3623 Infrared Study of the Adsorption of Ammonia, Pyridine and Hydrogen Chloride on Barium Sulphate William Neagle and Colin H. Rochester* Chemistry Department, The University, Dundee DDl4HN Infrared spectra are reported of ammonia, pyridine and hydrogen chloride adsorbed on barium sulphate in the presence and absence of surface hydroxyl groups and adsorbed water. Surface hydroxyl groups do not exhibit Brsnsted acidity but form hydrogen bonds with adsorbed acceptor molecules. Exposed Ba2+ cations constitute Lewis-acidic adsorption sites. Surface sulphate anions act as hydrogen-bond acceptor sites in the presence of hydrogen- bond donor molecules. The dissociative adsorption of water on barium sulphate generates hydroxyl groups at Ba2+ sites but does not lead to the concomitant formation of HSO; anions.The latter may be generated by the dissociative adsorption of hydrogen chloride and exhibit Brsnsted acidity in the presence of adsorbed pyridine. An infrared study of water adsorption on barium sulphate showed that dissociative adsorption to form surface hydroxyl groups was accompanied by non-dissociative adsorption probably of two types involving, first, ligation to weakly Lewis-acidic Ba2+ cations and, secondly, hydrogen bonding with exposed sulphate anions.' The present study of ammonia and pyridine adsorption was aimed at confirming the existence of Lewis-acidic surface sites and at probing the acidic character of surface hydroxyl groups. Hydrogen chloride was adsorbed on barium sulphate in an attempt to generate and characterise exposed HSO; anions.Experimental Details of the preparation and bulk characterisation of barium sulphate are given elsewhere. Experimental procedures for infrared study of self-supporting pressed discs or for gravimetric study of broken discs were as bef0re.l Samples were heated at 773 K under vacuum to remove bulk water and carbonate impurities before adsorption experiments were carried out.' Ammonia was purified by distillation on to sodium and subsequent distillation of the resulting blue solution. AnalaR pyridine was doubly distilled. Hydrogen chloride was prepared from sodium chloride and concentrated sulphuric acid. Results There are no surface hydroxyl groups on barium sulphate which had been heated in vacuum at 773 K.Subsequent exposure to ammonia at 298 K gave infrared spectra exhibiting bands at 3400, 3300, 3240 and 1625 cm-' [fig. 1 (a)], which were not present in spectra of ammonia alone. Evacuation of the cell with discs at 298 K led to the disappearance of the bands, showing that ammonia was only weakly adsorbed on the barium sulphate surface. Gaseous ammonia gives infrared bands at 3414 (v3 vibration), 3336/3338 (v,) and 1628 cm-' whereas the spectrum of solid ammonia contains corresponding bands at 3378, 3223 and 1646 ~ m - l . ~ Although not clear from fig. 1, several spectra recorded here for ammonia adsorbed on barium sulphate showed 36153616 Adsorption of Bases on BaSO, 3500 3000 1600 1400 wavenumber/cm-' Fig. 1. Spectra of barium sulphate evacuated at 773 K and (a) exposed to ammonia (332 N m-2) at ca.298 K and (b j(g) exposed to increasing pressures of water vapour. evidence for a shoulder at ca. 3365 cm-' on the broadest maximum at 3400 cm-'. The results therefore suggest that ammonia retains its molecular integrity when adsorbed on barium sulphate and that there are two types of adsorbed molecule characterised by v3 and v1 bands at 3400 and 3300 cm-' and at 3365 and 3240 cm-l. The lack of splitting of the maximum at 1625 cm-' due to v4 vibrations is not surprising since the v4 frequency is least sensitive to variation with changing electronic environment of the ammonia molecule. Addition of water vapour to ammonia adsorbed on barium sulphate gave no new infrared bands which could be ascribed to species deriving from ammonia.In particular there was no evidence for ammonium cations which would have been formed if water molecules had gained Brernsted-acidic character by adsorption at Lewis-acidic surface sites. Infrared bands due to adsorbed ammonia were apparently diminished in intensity, although this effect was obscured by the growth of an intense broad band envelope centred at 3450cm-' due to vibrations of adsorbed water molecules (fig. 1). The deformation vibration of water gave a maximum at 1630cm-' which obscured any effects involving the band at 1625 cm-' (due to adsorbed ammonia). All spectral features reported here in the 1200-1 700 cm-' region are presented with the background spectrum of barium sulphate' subtracted from the spectra. Shoulders at 3680 and 3580 cm-' on the broad band envelope ascribed to water may be assigned to OH stretching vibrations of surface hydroxyl groups which are also formed when water is adsorbed on barium sulphate in the absence of ammonia.' Evidence for the interpretation of the bands due to v, and v, vibrations of adsorbed ammonia in terms of two modes of adsorption was provided by the relative extents of depletion of the four infrared bands in the presence of water.The band at 3240 cm-l had almost disappeared from spectra at a stage [fig. 1 ( d ) and (e)] when the maxima at 3400 and 3300 cm-' remained prominent. TheW. Neagle and C. H. Rochester 3617 n E 3 2 c .9 b 4000 3500 3000 I0 Fig. 2. Spectra of barium sulphate (a) after evacuation at 773 K, (b) exposed to water vapour (2.39 kN m-2) and evacuated at 298 K (72 h), (c)-(i) in the presence of pressures/N m-' of ammonia of (c) 5.3, ( d ) 10.6, (e) 21.2, (f) 53.1, (g) 79.7, (h) 266 and (i) 2660.band at 3400 cm-lbecame depleted at the low-wavenumber side, supporting the suggestion of a shoulder at 3365 cm-' and its pairing with the band at 3240cm-l. Adsorbed ammonia giving the infrared bands at 3365 and 3240 cm-l was more readily displaced from the surface by water than was adsorbed ammonia, giving the infrared bands at 3400 and 3300 cm-l. The latter was present even after multilayer adsorption' of water [fig. 1 (f) and (g)]. Fig. 2 shows spectra resulting from the addition of ammonia to an hydroxylated barium sulphate surface which also has a residual coverage of associatively adsorbed water molecules.Spectra in the 1400-1 700 cm-l range are shown relative to the spectrum (6) for the hydrated/hydroxylated surface in order to be able to distinguish the growth of the band at 1625 cm-l due to the v, vibration of adsorbed ammonia molecules in the absence of the band at 1630 cm-l due to adsorbed water.' The bands due to the NH stretching vibrations of adsorbed ammonia molecules also grew in intensity although the species giving the bands at 3400 and 3300 cm-' was apparently more readily formed at low ammonia pressures than the species giving the bands at 3365(sh) and 3240 cm-l. This effect was particularly noticeable for barium sulphate which had been exposed to water vapour (2.39 kN m-2) and only evacuated for 30 min (cf. 72 h in fig.2) at 298 K before the addition of ammonia. Adsorbed water is more effective at blocking sites with which ammonia interacts to give the 3365 and 3240 cm-' bands than at blocking sites on which ammonia is adsorbed to give the bands at 3400 and 3300 cm-l. The band at 3680 cm-l due to isolated hydroxyl groups gradually disappeared from spectra with increasing pressures of ammonia in contact with barium sulphate (fig. 2). The band reappeared to its initial intensity when the disc was evacuated at 298 K, suggesting that the hydroxyl groups were perturbed by hydrogen-bonding interactions with adsorbed ammonia molecules. There was no evidence for the formation of ammonium ions. Exposure of a deuteroxylated/deuterated barium sulphate surface' to ammonia gave results compatible with those for the hydroxylated surface, although isotopic H/D exchange caused the replacement of OD groups by OH groups in the presence of an excess of ammonia.3618 Adsorption of Bases on BaSO, I 9 9 :s- A 3200 3000 A L 1600 1500 3200 3000 wavenumber/an- 10% I 8 1600 1500 Fig.3. Spectra of barium sulphate after evacuation at 773 K and A (aHc) exposure to increasing pressures of pyridine followed by evacuation (298 K, 30 min), A(d) and B(a) in contact with pyridine vapour (2.3 kN m-2), B(b)-(e) after evacuation (1 h) at (b) CQ. 298 K, (c) 373 K, ( d ) 423 K and (e) 473 K. Fig. 3A shows spectra of pyridine adsorbed on barium sulphate. Bands in the range 2900-3100 cm-' are due to CH stretching vibrations of adsorbed molecules. The bands due to ring vibrations which are commonly used to characterise surface sites were at 1612,1590,1575,1490, 1442 and, at high coverages, 1485 cm-'.The appearance of more than four bands' in this spectral region showed that more than one mode of adsorption of pyridine occurred on barium sulphate. This is borne out when the relative intensities of bands are considered as a function of increasing surface coverage. The ratio of intensities of the maxima at 1612 and 1575 cm-' changes with increasing amount of pyridine adsorbed and the band at 1485 cm-' only appears in the presence of pyridine vapour. Spectra of the pyridine vapour recorded with the disc lifted out of the infrared beam showed that the maximum at 1485 cm-l must have been due to an adsorbed form of pyridine. Desorption studies supported the suggestion that pyridine was adsorbed on a dry barium sulphate surface in two ways (fig.3B). Evacuation at 298 K largely removed adsorbed pyridine giving the band at 1575 cm-l but failed to desorb pyridine responsible for the band at 1612 cm-'. Partial reductions in band intensities at 1442, 1590 and 1485 cm-' also occurred, the latter band shifting to 1490 cm-'. Evacuation at 423 K left residual infrared bands at 1612, 1590, 1490 and 1442 cm-' which are attributed to pyridine which was ligated to exposed Lewis-acidic Ba2+ surface sites. The growth of bands at 1590, 1575, 1485 and 1442 cm-' in the presence of pyridine vapour is ascribed to a second mode of adsorption involving weak interactions between pyridine molecules and the barium sulphate surface.The absence of surface hydroxyl groups' obviated the possibility of hydrogen bonding between the surface and pyridine molecules.W. Neagle and C. H . Rochester 3619 I I 3500 3000 1700 1500 w avenumber/m- Fig. 4. Spectra of barium sulphate after (a) evacuation at 773 K, exposure to water vapour (2.39 kN mP2, 298 K) and evacuation (298 K, 16 h), (bk(e) exposure to increasing pressures of pyridine and evacuation (298 K, 30 min) and (f) exposure to deuterium oxide vapour (2.39 kN m-2) and evacuation (298 K, 10 min). Spectra of hydroxylated/hydrated barium sulphate showed that pyridine displaced both surface hydroxyl groups and adsorbed water molecules from adsorption sites (fig. 4). Displacement of water was incomplete even in the presence of an excess of pyridine, suggesting that one mode of adsorption of water involved surface sites (possibly exposed sulphate anions)' not preferred by pyridine.The band due to the bending vibration of adsorbed water shifted to 1640 cm-' in the presence of adsorbed pyridine and disappeared from spectra when water was replaced by deuterium oxide [fig. 4 0 1 . Bands due to vibrations of adsorbed pyridine were the same for both hydrated (fig. 4) and dehydrated (fig. 3) barium sulphate. However, the addition of deuterium oxide apparently removed a very weak band at 1540 cm-l, which could be ascribed to pyridinium ions,6 hence suggesting the possible existence of a very low population of Brsnsted-acidic surface sites on barium sulphate. The infrared band is absent from spectra of deuterated pyridinium ions pyD+.' The band at 1540 cm-l was also just discernible in spectra of pyridine on barium sulphate which had been preheated at 773 K (fig. 3) and which, before the adsorption of pyridine, gave spectra containing no clearly detectable evidence for surface hydroxyl gr0ups.l The existence of the very weak band at 1540 cm-' was not therefore due to Brsnsted acidity of hydroxyl groups responsible for the infrared band at 3680cm-' [fig.2(b)]. The present evidence for surface Brsnsted-acidic sites is not strong and must be open to question. If they do exist, then their surface concentration must be small compared with the concentration of surface hydroxyl groups generated by exposure of barium sulphate to excess water vapour at ambient temperature.The ammonia adsorption experiments confirm that the latter hydroxyl groups do not readily exhibit Brsnsted-acidic behaviour. The adsorption of hydrogen chloride on barium sulphate initially devoid of surface hydroxyl groups led to the appearance in spectra of a broad maximum centred at3620 Adsorption of Bases on BaSO, I I 1300 Fig. 5. Spectra of barium sulphate after evacuation at 773 K and (a)-(e) exposure to increasing pressures of hydrogen chloride at 298 K. I I I ~ I I I I I 4000 3500 3000 2500 w avenumber/an-' I I I I - 1700 1600 Fig. 6. Spectra of barium sulphate after (a) evacuation at 773 K, exposure to water vapour (2.39 kN m-2, 298 K) and evacuation (298 K, 16 h), (b)-(e) subsequent addition of increasing pressures of hydrogen chloride and (f) evacuation (298 K, 30 min).2900-3000 cm-l which shifted to ca. 2600 cm-l with increasing surface coverage (fig. 9, and a second maximum at 1310 cm-l. These bands may be assigned to vibrations of HSO, anionss. generated by the protonation of exposed sulphate ions. Similar bands appeared when hydrogen chloride was admitted to barium sulphate with surface hydroxyl groups and adsorbed water molecules (fig. 6). Surface hydroxyl groups (3680 cm-l) were displaced by hydrogen chloride and were not restored when HC1 was removed by evacuation. Bands at 3450 and 1630 cm-l [fig. 6(a)] due to vibrations of water molecules were depleted in intensity, broadened and shifted towards ca. 3400 and 1650 cm-l in the presence of adsorbed hydrogen chloride. A broad band also appeared at 3000 cm-l.Bands at 3400, 3000 and 1650 cm-' may be assigned to vibrations of hydroxonium ions H,O+.'9 lo The existence of Brarnsted-acidic surface species wasW. Neagle and C. H. Rochester 362 1 - 1700 1500 w avenumber/an- Fig. 7. Spectra of barium sulphate after (a) treatment similar to that for fig. 6 ( f ) and (b)-(d) exposure to increasing pressures of pyridine followed by evacuation (298 K, 30 min). confirmed by pyridine adsorption which led to spectra (fig. 7) containing bands at 1640 and 1540 cm-' characteristic6 of pyridinium ions. The band at 1310 cm-l disappeared from spectra when pyridine was adsorbed showing that HSO, anions behaved as Brlernsted acids in the presence of pyridine. A similar effect was observed for barium sulphate which had not been exposed to water vapour before being contacted with hydrogen chloride and subsequently evacuated at ca. 298 K prior to pyridine adsorption.No surface H,O+ ions were present before pyridine adsorption. However, infrared bands still appeared at 1640 and 1540 cm-l in parallel to the disappearance of the band at 1310 cm-', confirming the behaviour of HSO, anions as Brarnsted-acidic surface sites. The gravimetrically determined isotherm for the adsorption of ammonia on barium sulphate at 298 K was of type 11'' and gave an estimated monolayer capacity of ca. 0.2 x mol m-2, which is only slightly less than the monolayer capacity of barium sulphate for water in the presence of water vap0ur.l In accordance with the infrared results, evacuation at 298 K resulted in complete desorption of ammonia from the barium sulphate surface.Pyridine adsorption gave a type I isotherm with a monolayer capacity in the presence of pyridine of 0.12 x mol m-2. Subsequent evacuation at 298 K led to only partial desorption of pyridine and a residual surface coverage (after attainment of monolayer coverage in the presence of pyridine) of 0.09 x mol m-2. The gravimetric isotherms for ammonia and pyridine refer to a barium sulphate surface which had been dehydroxylated and dehydrated by heat treatment in vacuum at 773 K. Discussion There are two distinguishable modes of adsorption of pyridine on a dehydroxylated barium sulphate surface. Both the infrared and gravimetric results show that one mode of adsorption (bands at 1612, 1590, 1490 and 1442 cm-l) is somewhat stronger (not desorbed by evacuation at 298 K) than the other (bands at 1590, 1575, 1485 and 1442 cm-l; desorbed at 298 K).Band positions for the stronger mode of adsorption are3622 Adsorption of Bases on BaSO, consistent6 with the existence of exposed Lewis-acidic Ba2+ cationic sites in the barium sulphate surface. If the second mode of adsorption also involves Ba2+ sites then both the infrared band positions and the ease of desorption show that the ligating interactions at this second type of site must be extremely weak. One possibility would be that the approach of pyridine molecules to the Ba2+ sites is sterically hindered by adjacent sulphate anions. Multilayer adsorption of pyridine did not occur at 298 K. The infrared results for ammonia adsorbed on dehydroxylated barium sulphate could be ascribed to interactions involving two types of Lewis-acidic Ba2+ surface site.However, the gravimetric results show that the number of adsorbed ammonia molecules at monolayer coverage on barium sulphate was ca. 12 nm-2, which exceeds the estimated' number of ca. 4 nm-' for exposed Ba2+ ions in the surface. An explanation would be that some ammonia molecules are adsorbed at Lewis-acidic sites, but that a second mode of adsorption involves hydrogen- bonding interactions with exposed sulphate anions. A similar proposal has been made for water adsorption on barium sulphate.' A parallel between the behaviour of ammonia and water is shown by the similar monolayer capacity of barium sulphate for the two adsorbate molecules.In contrast, the monolayer capacity of barium sulphate for pyridine, which cannot donate hydrogen bonds, was only ca. half of that for water and ammonia. The ability of ammonia and water to form hydrogen-bonded aggregates favoured multilayer adsorption at high relative pressures of the gaseous adsorbates. The competitive adsorption of ammonia and water showed that adsorbed ammonia responsible for infrared bands at 3365 and 3240 cm-' was either displaced by water or its adsorption was impeded by preadsorption of water. These band positions are not far removed from those for solid a m m ~ n i a , ~ therefore they are tentatively ascribed to ammonia molecules involved in hydrogen-bond-donating interactions with surface sulphate anions. Ammonia adsorbed via ligands with Ba2+ cations was responsible for infrared bands at 3400 and 3300 cm-' and was not appreciably displaced or impeded by water. Providing there are no complications, such as multicentre surface-adsorbate bonding, the present proposals require that ammonia is a stronger electron donor and a weaker hydrogen-bond donor than water.Hydrogen chloride is dissociatively adsorbed on dehydroxylated barium sulphate in accordance with the reaction Ba2+ SO:- + HCl(g) --+ Ba2+Cl- HSO, to give C1- anions at vacant exposed Ba2+ ion ligand positions and HSO; anions which exhibit Brarnsted acidity". l3 in the presence of adsorbed pyridine. On a hydroxylated surface reactions (2) and (3) lead to the displacement of hydroxyl groups and water molecules, respectively : Ba2+0H- + HCl(g) - Ba2+C1- + H,O Ba2+ ...OH, + HCl(g) --+Ba2+C1- + H,O+(ads). (3) Protonation of sulphate ions by HCl or H,O+ also occurs and generates HSO; anions in the surface. Infrared spectra show that there must be an excess of H,O+ ions generated above the number of exposed sulphate anions available for protonation. The experiments involving HCl and pyridine adsorption help to confirm that neither of the surface reactions (4) or (5) occur when water is adsorbed on a dehydroxylated barium sulphate surface : Ba2+SOi- + H20(g) Ba2+0H- HSO, (4) Ba2+ + 2H20(g) Ba2+OH- + H,O'(ads). ( 5 ) The question therefore arises that if water is dissociatively adsorbed to form a surface hydroxyl group then what happens to the residual proton from the water molecule? The level of Brmsted-acidic activity after water adsorption was scarely detectable andW.Neagle and C. H. Rochester 3623 therefore neither HSO, not H,O+ ions must constitute significant adsorption products. One possibility would be that the surface of dehydroxylated barium sulphate contains a small concentration of 02- anions which are able to accept protons and hence generate a second hydroxyl group for each water molecule adsorbed. The low population of surface hydroxyl groups on fully hydroxylated barium sulphate' would be governed by the availability of exposed 02- anions which may be preferentially located at edge of defect sites. Hydroxyl groups formed at Ba2+ sites by water adsorption do not exhibit Brrzrnsted acidity but can act as hydrogen-bond donors to adsorbed acceptor molecules. We thank the S.E.R.C. for a research studentship, Unilever Research for collaboration through the CASE scheme and Dr I. D. Robb for helpful discussions. References I W. Neagle and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1988, 84, 3625. 2 H. Y. Sheng, E. F. Barker and D. M. Dennison, Phys. Rev., 1941, 60, 786. 3 C. M. Lewis and W. V. Houston, Phys. Rev., 1933, 44, 903. 4 F. P. Reding and D. F. Horning, J. Chem. Phys., 1951, 19, 594. 5 K. Nakamoto, Infrared Spectra of Inorganic Coordination Compounds (John Wiley, New York, 1963), 6 E. P. Parry, J. Catal., 1963, 2, 371. 7 S. J. Puttock and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1986, 82, 2773. 8 G. E. Walrafen and D. M. Dodd, Trans. Faraday SOC., 1961, 57, 1286. 9 R. Savoie and P. A. Giguere, J. Chem. Phys., 1964, 41, 2698. 10 D. E. Bethel1 and N. Sheppard, J. Chem. Phys., 1953, 21, 1421. 1 1 K. S. W. Sing, in Characterisation of Powder Surfaces, ed. G. D. Parfitt and K. S. W. Sing (Academic Press, London, 1976), p. 1. 12 0. Redlich, Chem. Rev., 1946, 39, 333. 13 W. L. Marshall and E. V. Jones, J. Phys. Chem., 1966, 70, 4028. V O ~ . 111, p. 143-151. Paper 8/00334C; Received 22nd January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403615

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1988, 84(10), 3625-3632 Infrared and Gravimetric Studies of the Adsorption of Water on Barium Sulphate William Neagle and Colin H. Rochester* Chemistry Department, The University, Dundee DDl 4HN Barium sulphate precipitated from aqueous solution contains bulk water which can be removed by prolonged evacuation at 773 K. Infrared spectra of barium sulphate exposed to water vapour show that both associative and dissociative adsorption of water occurs. Water is desorbed from multilayers by evacuation at ca. 298 K and from the first monolayer by evacuation at 383 K. Water in the first monolayer is adsorbed partly at weak Lewis acidic Ba2+ cations and partly via hydrogen-bonding interactions with exposed sulphate anions. Hydroxyl groups exist either in low-area planes or at defect or edge sites and are also removed by heat treatment at ca.383 K. Studies of the adsorption of gases on barium sulphate have demonstrated the surface heterogeneity of the solid.ll Comparison of the adsorption of aliphatic and aromatic hydrocarbons on barium sulphate showed that the adsorption of benzene involved interaction between the electrostatic field at the surface and ;n-electrons in benzene nuclei3 Adsorption from solution on to barium sulphate has primarily involved polymers adsorbed from water, and infrared spectroscopy has proved helpful in the characterisation of surface However, the advantage of infrared spectroscopy for the surface characterisation of solid particles has not been fully exploited for barium sulphate. The present paper reports an infrared and gravimetric study of water vapour adsorption on barium sulphate which had been characterised by X.r.d., t.g.a., s.e.m., i.r.spectroscopy and surface area and porosity measurements. Infrared studies of probe molecules on barium sulphate will be published elsewhere. Experimental AnalaR barium chloride and ammonium sulphate were recrystallised from water (deionised and triply distilled, once from alkaline potassium permanganate and twice from itself, all under nitrogen) and used to prepare barium sulphate by the method of Cafe and Robb.’ The product was washed three times in mixtures of water- methanol-ammonia and five times in water-methanol, contact times between solid and liquid in the successive washings ranging from 30 min to 1 week.Finally the solid was removed by centrifugation, dried under vacuum (333 K, 6 h and 293 K, 16 h), finely ground, dried under vacuum (453 K, 6 h and 293 K, 16 hj and reground before storage in a desiccator. X.r.d. traces were recorded using a Phillips instrument with a PWlOlO generator and Cu K, radiation. T.g.a. involved study of barium sulphate at 290-1073 K in a flow (25 cm3 min-l) of nitrogen using a Stanton Redcroft TG750 thermobalance. Surface areas were determined by the standard B.E.T. method with nitrogen as the adsorbate gas. Porosity was monitored with a Philips mercury porosimeter and electron micrographs were recorded with a Jeol JSM-35 scanning electron microscope. Infrared spectra were recorded with a Perkin-Elmer 580 spectrometer in conjunction with a 3600 infrared data station.36253626 15 w loo 10 “E !i H \ ‘t: 2 5 Water on Barium Sulphate I I 1 I I 0.0 n E v) ,o 0.5 3 1 .o 373 573 TIK 773 Fig. 1. (a) Surface area and (b) weight loss of barium sulphate as a function of evacuation temperature. Infrared study of barium sulphate involved pressed (50-60 MN mP2 on the die) self- supporting discs of 25 mm diameter and weights in the range 100-200 mg. The infrared cell had an external furnace and fluorite optical windows and was glassblown to a conventional vacuum apparatus capable of maintaining a dynamic vacuum of ca. lo-* N mP2. For comparative purposes gravimetric study of water vapour adsorption involved barium sulphate which had been pressed into self-supporting discs under identical conditions to those employed in the infrared experiments.A C.I. Instruments mark IIA vacuum microbalance was used. Results The X.r.d. pattern for the present material was consistent with previous data for synthetic barium s ~ l p h a t e . ~ The X.r.d. lines were sharp, indicating a high degree of crystallinity. The results were unaffected by heat treatment of the barium sulphate in vacuum at 523, 673 or 773 K. A typical t.g.a. trace for barium sulphate is shown in fig. 1 (b). Three regions of weight loss occurred in the temperature ranges 29&373, 450-550 and 620-1000K. A final steady weight occurred above ca. 1073 K and corresponded to a total weight loss of ca. 1.8%. A sample heated at 773 K for 24 h gave a weight loss of 2.9%. After heat treatment at 773 K (30 min) one sample was exposed to the laboratory atmosphere for 30 min before being subjected to t.g.a. The resulting trace only exhibited the low- temperature weight loss (ca.0.1 %) which is ascribed to the removal of weakly adsorbed water from the barium sulphate surface. A series of 18 samples of barium sulphate were subjected to heat treatment under vacuum and their surface areas subsequently determined. The results are shown in fig. 1 (a). The area was 10.4 m2 8-l after evacuation at 298 K for 16 h. This decreased to 9.7m2 g-l after evacuation at 373 K (48 h). A maximum area of 13.8 m2 g-’ was observed after evacuation at 573 K (16 h). Evacuation at higher temperatures resulted in lower areas, falling to 6.3 m2 g-’ after heating to 773 K (24 h) and further to 2.7 m2 g-l after 48 h at 773 K.W.Neagle and C. H . Rochester 3627 4000 3500 3000 1800 1600 1400 1200 w avenumber/cm-' Fig. 2. Infrared spectra of barium sulphate (a) in air, and (b)-(r) evacuation at temperatures/K of (b) 298 (1 h), (c) 298 (20 h), ( d ) 483 (2 h), (e) 523 (1 h), (f) 523 (7.5 h), ( g ) 523 (18.5 h), ( h ) 523 (72 h), (i) 583 (1 h), (j) 598 (12 h), ( k ) 598 (24 h), (I) 623 (2 h), (m) 623 (17.5 h), (n) 683 ( 5 h), (0) 688 (2.5 h), ( p ) 723 (12 h), (4) 773 (6 h) and ( r ) 773 (24 h). Mercury porosimetry showed that the barium sulphate contained no pores of diameter > 5 nm. The absence of extensive microporous structure was confirmed by the consistency between surface areas (fig. 1) and the average dimensions of barium sulphate particles estimated by s.e.m.Low magnification showed roughly spherical particles with ca. (1-7) x m diameters, although high-resolution micrographs showed that these were aggregates of much smaller particles with average diameter (1.2 1 .O) x lW7 m. This size corresponds to an area of ca. 13 m2 g-l. Heating samples in vacuum at 773 K for short times showed no clear-cut changes in the electron micrographs. However, evacuation at 773 K for 168 h led to increases in particle size corresponding to a reduction in estimated surface area to 5 m2 g-l. Nitrogen adsorption and B.E.T. analysis gave 0.5m2 g-l as the area of this sample. The discrepancy here may be due to interparticulate fusion at points of particle contact for which there was evidence in the electron micrographs.Infrared spectra of barium sulphate in a KBr disc exhibited bands at 1188, 1120 and 1075 cm-' [v,(SOq-)], 980 cm-l [v,(SOq-)], 637 and 606 cm-l [v,(SO:-)] in good agreement with literature data.'', l1 Very weak bands at ca. 3420 and 1640 cm-l showed that some water was present in the sample. The presence of bands due to molecular water was more clearly revealed in spectra of self-supporting discs of barium sulphate (fig. 2). The latter spectra also contained, in accordance with previous results and assignments of Schroeder et a1.,12 several overtone and combination bands ascribable to vibrations of the sulphate anion and bands at 1437, 1400, 925, 890 and 876cm-' due to small concentrations of carbonate impurities. Additional sharp bands at 1355 and 1670 cm-' (fig.2) may also be due to carbonate species.13 Fig. 2 shows the effects of heat treatment of self-supporting discs of barium sulphate under vacuum on infrared bands due to hydroxyl and carbonate species. Evacuation at 298 K resulted in the disappearance of a weak shoulder at 3690 cm-l and significant decreases in the intensities of maxima at 1633 and 3600-2800 cm-'. Similar decreases in band intensities could also be achieved by exchange with deuterium oxide vapour. Depletion of the maxima was therefore due to loss of adsorbed water and surface hydroxyl species. The residual maxima at 360&2800 and 1633 cm-l after evacuation at3628 Water on Barium Sulphate 1 I I I I I 10% !L 4000 3500 1700 2800 2500 wavenumber/m-' Fig. 3. Spectra of barium sulphate after (a) evacuation at 773 K, ( b t ( f ) contact with increasing vapour pressures of water and evacuation (30 min, ca.300 K). Spectra ( b ' ) - ( f ) are corresponding difference spectra obtained by subtraction of the background (a). Spectra (g)-(l) are difference spectra for the adsorption of deuterium oxide on barium sulphate. ca. 383 K were unaffected by prolonged exposure to deuterium oxide vapour at ambient temperature and are therefore ascribed to molecular water and hydroxyl groups occluded within the barium sulphate bulk structure. Heat treatment at increasing temperature led to the gradual disappearance of water from the bulk lattice. At the same time the bands at 1670 and 1355 cm-l, due to carbonate species, disappeared at 688 K. The more intense band maxima at 1437 and 1400 cm-l, due to cabonate ions, diminished in intensity with increasing temperature but still remained as sharp bands after treatment at 773 K for 24 h.These bands were unaffected by the adsorption of probe molecules on the barium sulphate surface, therefore the carbonate ions probably existed at sites in the bulk lattice. The bands could be enhanced in intensity by the deliberate addition of carbonate ions to the solutions used in the preparation of barium sulphate. Even the most rigorous attempts to remove carbon dioxide from preparative solutions failed to produce barium sulphate which was completely free from carbonate ions. A self-supporting disc of barium sulphate was freed from adsorbed and bulk water by heat treatment at 773 K. Subsequent adsorption of water vapour gave spectra (fig.3) containing infrared bands at 3700 (sh) and 3680 cm-' due to isolated hydroxyl groups, at 3580cm-' ascribed to hydrogen-bonded hydroxyl groups and at 3450, 3200 and 1630 cm-l due to vibrations of water molecules. A plot of intensities at the two main adsorption maxima (3680 and 3450 cm-l) against each other was curved in the sense that the extent of non-dissociative adsorption of water became increasingly greater with increasing contact vapour pressure of water than the extent of dissociative adsorption to form surface hydroxyl groups. Adsorption of deuterium oxide vapour on barium sulphate gave infrared bands at 2730 (sh, vw), 2700 and 2630 cm-l due to OD groups and at 2520 and 2380 cm-' assigned to vibrations of D,O molecules.The expected band at ca. 1200 cm-l due to the deformation vibration of D,O molecules was obscured by the intense maximum at 1188 cm-l resulting from the v, vibration of sulphate anions.W. Neagle and C . H. Rochester 3500 3( 3500 3000 w avenumkr /an- 3629 Fig. 4. (a) through to (b) : Spectra of barium sulphate after exposure to water vapour at saturated vapour pressure at ca. 293 K and evacuation at 373 K for increasing times. Spectra of barium sulphate after exposure to water vapour (2.4 kN m-2) and evacuation at (c) ca. 298 K (16 h), ( d ) 328 K (24 h), (e) 373 K (1 h), and (f) 383 K (1 h). Confirmation that the sharp bands at 3680 and 2720 cm-l could be ascribed to surface OH and OD groups, respectively, rather than to vibrations of adsorbed water or deuterium oxide molecules was obtained14 by recording spectra of HDO adsorbed on barium sulphate.The only two sharp bands in the spectra were those at 3680 and 2720 cm-l. If these were due to vibrations of H,O and D,O, respectively, then HDO would have given two new bands in easily distinguishable spectral positions. Fig. 4 exemplifies the results of water desorption experiments involving self- supporting discs of barium sulphate. Spectra 4(a) to (b) were recorded using the repeat-scan mode of the spectrometer, each scan taking ca. 5 min. Comparisons of the rates of change of absorbances at 3680,3580 and 3450 cm-l showed that the relative ease of removal of surface species was in the sequence: associatively adsorbed water molecules > hydrogen-bonded OH groups > isolated OH groups.The large change in the overall intensity of the broad absorption envelope centred at 3450cm-l which occurred between the first [fig. 4(a)] and second scans may be ascribed to the removal of water from multilayers which exist in equilibrium with pressures of water vapour approaching the saturated vapour pressure. Evacuation at 383 K was sufficient to remove all traces of water or surface hydroxyl groups from barium sulphate [fig. 4 0 1 . The isotherm for the adsorption of water vapour was of type 11, indicating unrestricted monolayer-multilayer adsorption on a heterogeneous surface. l5 Estimation of monolayer capacity from the isotherm,lG combined with a value 0.125 nm2 for the cross-sectional area of an adsorbed water molecule,17 gave 10.6 m2 g-l for the barium sulphate surface area.The sample had been preheated in vacuum at 773 K (24 h) to remove bulk water and reduce the level of carbonate impurity before water adsorption. Nitrogen adsorption at 77 K gave 6.3 m2 g-l for a similarly pretreated sample. This represents good agreement in view of possible uncertainties concerning the use of water adsorption for surface area determinati~nl~ and also the presently observed time- 119 FAR I3630 Water on Barium Sulphate 0 1 2 water vapour pressure/kN m-2 Fig. 5. Adsorption of water on barium sulphate (a) isotherm at 298 K, (b) and (c) residual water after subsequent evacuation at 298 K for 30 min and 1 h, respectively. dependent sintering effects of heat treatment at 773 K on barium sulphate.The occurrence of dissociative adsorption may contribute to the higher apparent monolayer capacity for water than for nitrogen. After evacuation at 298 K amounts of residual water remaining on the barium sulphate surface corresponded to ca. 0.1 x lo-* mol m-2 [fig. 5(b) and (c)]. The infrared spectra show that after evacuation at 298 K the surface contained both hydroxyl groups and residual adsorbed water molecules. The concentration of surface hydroxyl groups was estimated from the gravimetric and spectroscopic data. The calculation involved an iterative procedure in which it was assumed that the extinction coefficient at 3450 cm-l was the same for water in the first adsorbed layer and in multilayers. The surface population of hydroxyl groups was 0.30nm-2 after barium sulphate had been exposed to the saturated vapour pressure of water and subsequently evacuated for 24 h at 298 K.The residual population of adsorbed water molecules was ca. 5.3 nm-2. Discussion Both t.g.a. and infrared results for barium sulphate suggest the presence of bulk water, in accordance with previous observations. Walton and Walden18 proposed that three water molecules replaced one barium sulphate unit in the bulk structure, and that loss of water on heat treatment resulted in the formation of voids in the crystal lattice. The suggestion that one water molecule is present for every 5-10 barium sulphate 23 would be compatible, taking into account the losses of surface water and bulk carbonate, with the present weight changes on heating.The increase in surface area of barium sulphate after heat treatment at ca. 373-573 K could arise because loss of water from the lattice may induce cleavage of crystallites to give smaller particles or may lead to the creation of microporous holes near the solid surface which are accessible to adsorbate molcules. Electron micrographs of theW, Neagle and C. H. Rochester 363 1 barium sulphate before and after heat treatment at 573 K (16 h) showed no detectable differences, suggesting that particle cleavage was not an important factor. Voids in the surface region induced by loss of water18 generate micropores which were accessible to nitrogen molecules but were too narrow to be detectable by mercury porosimetry. Decreases in surface area on high-temperature treatment of barium sulphate have been reported before.20*2' One studyz0 gave similar weight losses to the present results but considerably greater and more rapid losses in surface area.This difference in behaviour is attributed to the much longer times for which the barium sulphate was aged in the present preparative procedure. It has been shown that rapidly prepared precipitates give significant decreases in surface area on being heated at high temperatures, but that slowly prepared precipitates are less susceptible to change in surface area.2' Thermal mobility of ions in barium sulphate at elevated temperatures probably accounts for the losses in surface This may lead to the annealing of accessible voids generated by loss of water or fusion of particles in closely packed aggregates.Electron micrographs suggest that the latter effect was important here. A conclusion that water is not chemisorbed on barium s ~ l p h a t e ~ * + ~ ~ is at variance with the infrared spectroscopic evidence that water dissociates on barium sulphate to form surface hydroxyl groups. The chemisorption reaction was completely reversed by evacuation at ca. 383 K (fig. 4). The estimated populations of hydroxyl groups (0.30 nm-2) and water molecules (5.3 nm-2) on barium sulphate after saturation with water and evacuation at 298 K compares with a figure of ca. 3.94.2 nmP2 for the surface concentration of barium ions estimated on the assumption that the (001) and (210) faces constitute the predominant exposed crystal planes. The excess of water molecules over barium ions probably arises because some water molecules may be adsorbed via hydrogen-bonding interactions with oxygen atoms in exposed sulphate anions.This mode of adsorption and the proposal that some water molecules would be ligated to weak Lewis-acidic Ba2+ surface sites would be compatible with the ease of desorption (ca. 383 K) of water from the barium sulphate surface. The relatively low population of surface hydroxyl groups suggests that these exist either on an exposed plane which only contributes a small amount to the total area of the solid particles or at defect or edge sites. We thank the S.E.R.C. for a research studentship, Unilever Research for collaboration through the CASE scheme and Dr I. D. Robb for helpful discussions. References 1 N.N. Avgul, G. I. Berezin, A. V. Kiselev, I. A. Lygina and G. G. Muttik, Zh. Fiz. Khim., 1957, 31, 2 N. N. Avgul, G. I. Berezin, A. V. Kiselev and I. A. Lygina, Bull. Acad. Sci. USSR Div. C h h . Sci., 3 L. D. Belyako, A. V. Kiselev and M. V. Lomosov, Bull. Acad. Sci. USSR Dit:. Chem. Sci. 1962, 969. 4 S. S. Khamraev, A. I. Yusupov and K. S . Akhmedov, Uzt. Khim. Zh., 1974, 18, 29. 5 Y. V. Erman, A. V. Uvarov, S. N. Tolstaya and N. A. Aleksandrova, Makromol. Granitse Radzela 6 B. Dobias, Colloid Polym. Sci., 1977, 255, 682. 7 M. C. Cafe and I. D. Robb, J . Colloid Interface Sci., 1982, 86, 41 I . 8 D. R. Bain, M. C. Cafe, I. D. Robb and P. A. Williams, J . Colloid Interface Sci., 1982, 88, 467. 9 A. A. Colville and K. Staudhammer, Am. Mineral., 1967, 52, 1877. I l l . 1960, 1948. Faz., 1971, 100. 10 S. D. Ross, Spectrochim. Acta, 1962, 18, 1575. 11 P. Tarte and G. Nizet, Spectrochim. Acta, 1964, 20, 503. 12 R. A. Schroeder, E. R. Lippencott and C. E. Weir, J. Inorg. Nucl. Chem., 1966, 28, 1397. 13 G. Busca and V. Lorenzelli, Muter. Chem., 1982, 7 , 89. 14 D. M. Griffiths and C. H. Rochester, J . Chem. Soc., Faraday Trans. I , 1977, 73, 1510. 15 K. S. W. Sing, in Characterisation of Powder Surfaces, ed. G. D. Parfitt and K. S. W. Sing (Academic Press, London, 1976), p. 1. 119-23632 Water on Barium Sulphate 16 S. Brunauer, P. H. Emmett and E. Teller, J. Am. Chem. Soc., 1938, 60, 309. 17 A. L. McClellan and H. F. Harnsberger, J. Colloid Interface Sci., 1967, 23, 577. 18 G. Walton and G. A. Walden, J. Am. Chem. Soc., 1946, 68, 1750. 19 V. B. G. Varma, Chem. Age India, 1963, 14, 329. 20 I. M. Kolthoff and W. M. McNevin, J. Phys. Chern., 1940, 44, 921. 21 E. Buzagh-Gere, F. Paulik and L. Erdey, Talania, 1966, 13, 731. 22 F. Paulik, E. Buzagh, L. Polos and L. Erdey, Acta Chim. Acad. Sci. Hung., 1963, 38, 31 1. 23 L. Erdy, F. Paulik, E. Buzagh and L. Polos, Acta Chim. Acad. Sci. Hung., 1964, 41, 109. 24 T. Morimoto, J. Takeshita and M. Nagao, Rep. Res. Lab. Surf Sci., Okayama Unio., 1973, 4, 45. 25 T. Morimoto, J. Kishi, 0. Osamu and T. Kadota, Bull. Chem. Soc. Jpn, 1980, 53, 1918. Paper 8/00335A; Received 22nd January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403625

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1988, 84(10), 3633-3639 Temperature-programmed Desorption of p-Xylene from ZSM-5, ZSM-11 and Theta-1 Li-feng Chen and Lovat V. C. Rees* Physical Chemistry Laboratories, Imperial College of Science and Technology, London S W7 2A Y The temperature-programmed desorption of p-xylene from three high- silica-type zeolites, ZSM-5, ZSM- 1 1 and Theta- 1, in their mixed H+/Na+ forms has been determined. Using two recently developed analytical procedures the activation energies and entropies of desorption as a function of coverage have been obtained. The saturation capacities of these three samples for p-xylene have been measured. The saturation capacities and desorption energies have also been determined in ZSM-5 and ZSM-11 samples with varying Si/Al ratios.Two new methods of analysing temperature-programmed desorption (t.p.d.) profiles have recently been reported from our laboratory. In the first method the profiles from ten different heating rates are analysed as a function of coverage. This method will be referred to as the variable heating rate (v.h.r.) meth0d.l In the second method the profile from only one heating rate is analysed, and will be referred to as the single heating rate (s.h.r.) method.2 Both methods will be used to obtain activation energies and entropies of desorption of p-xylene from three high-silica-type zeolites. The effect of change of the Si/Al ratio on these energies will also be reported. Theory By assuming first-order desorption kinetics, a linear heating rate and the Arrhenius equation to be applicable to t.p.d.kinetics eqn (1) was obtained:l exp (- E/RT) dT where 8 is the fractional coverage and p is the heating rate. 8, and T, are the initial coverage and temperature, respectively, and Oi and the corresponding values at some time, t . A is the pre-exponential function and E the energy of activation for desorption in the Arrhenius equation. The values of A and E are obtained at specific values of Bi using a minimization procedure and an approximate solution of the temperature integral in eqn (1) due to Balarin in the v.h.r. method.' In the s.h.r. method an exact solution of the temperature integral is obtained by the introduction of a new variable u = E/RT. A and E are now obtained over a small range of Bi using a minimization procedure which is fully described in ref.(2). Experimental A Stanton-Redcroft TG 762 thermogravimetric balance with an on-line Commodore CBM 8032 microcomputer was used for both collection and analysis of the t.p.d. data. Ca. 4.5 mg of zeolite was dehydrated in situ in the t.g. balance by heating to 673 K at 10 K min-l in a dry argon stream (30 cm3 min-l). After cooling to room temperature a 36333634 T.P.D. of p-Xylene from Zeolites secondary argon stream was passed through the liquid p-xylene (East Anglia Chemicals) and then over the sample in the balance until sorption equilibrium attained. The relative pressure of adsorbate was p/po z 0.5. Ten heating rates from 2 to 20 K min-’ were used in the v.h.r. method, while 6 and 10 K min-l were used for the s.h.r.method. In the first part of this study the three zeolites used had the following unit-cell formulae after calcination in air at 823 K for 60 h to burn off the template: ZSM-5 (Si/Al = 15) : H4~48Nal~,2A16SigoOlg2 ZSM-11 (Si/Al = 15) : H4~36Nal~66A16Sigo0,g2 Theta- 1 (Si/Al = 32) : Ho~45Na2~46A12~glSig3~ogOlg2. In the second part the unit-cell compositions of the samples after calcination were: ZSM-5 (45) ZSM-5 (78) H0.54Na0.68A11.22si94.780192 Silicalite-1 (8200) Si/Al = 8200 The numbers in parentheses indicate the Si/A1 ratio, and this notation will be used in the second part when we refer to these samples. Results and Discussion The saturation capacities for p-xylene of the low Si/Al zeolites obtained in the t.g. balance at p/po z 0.5, room temperature and a sorption time of 45 min were 5.96, 6.21 and 1.97 molecules per unit cell in ZSM-5, ZSM-11 and Theta-1, respectively.ZSM-11 contains two sets of straight channel^,^ while ZSM-5 contains one set of similar channels intersected orthogonally by another set of sinusoidal channel^.^ In ZSM-5 the channel intersections are all equivalent, whereas in ZSM-11 there are two types of intersections; one has the same volume as those in ZSM-5 while the other has a volume which is ca. 30 % greater.5, This accounts for the greater saturation capacity of ZSM- 1 1. Theta- 1 has only linear channels, parallel to the c axis, with no intersections, and has a much smaller theoretical sorption capacity compared with ZSM-5 or The saturation capacities found with these low-Si/Al zeolites are much smaller than those found with the ‘pure’ silica analogues of these framew0rks.l’ These low-Si/Al samples have H+ and Na+ ions present in the channels and intersections which may block part of the channel system or, more likely, reduce the packing efficiency of the p-xylene molecules in the channels and intersections.ZSM- 1 1. Temperature-programmed Desorption Studies From the differential of the t.p.d., profiles of p-xylene in ZSM-5, ZSM-11 and Theta-1 the peak temperatures and peak widths shown in table 1 were obtained. The similarity of these temperatures and widths in table 1 indicate that the desorption energies of p-xylene from these three zeolites should be similar. The variation of the desorption energy, Ed, with coverage calculated by the v.h.r.method is shown in fig. 1 for the three zeolite samples, while in fig. 2 the corresponding variation of the entropy of desorption, -ASS, is demonstrated. The positive entropies of desorption observed for ZSM-11 and Theta-1 over part of the coverage range is difficult to explain. The value of E d averaged over the whole coverage range, E d , is given in table 2. The significant increases in Ed for ZSM- 1 1 and Theta- 1 over the correspondingL. F. Chen and L. V. C. Rees 120- - 100. I 1 2 80- 3635 60- Table 1. T.p.d. peak temperatures and widths for p-xylene: p = 10 K min-' peak temperature/K peak width/K higher zeolite peak I peak I1 peak peak I peak I1 ZSM-5 320 405 peak I 300-370 370-470 ZSM- 1 1 320 405 peak I 300-370 370-470 Theta- 1 315 395 peak I 300-370 370-460 1 0 1.0 2.0 3.0 4.0 5.0 coverage/molecules per unit cell Fig.1. Activation energy of desorption of p-xylene as a function of coverage (v.h.r. method of analysis) : 0, ZSM-5 ; 0, ZSM- 1 1 and 0, Theta- 1 . value for ZSM-5 is not consistent with the similar peak temperatures in table 1 for these three zeolites. The corresponding variations of the desorption energy, E d , and entropy, - ASt, with coverage calculated by the s.h.r. method are shown in fig. 3 and 4, respectively. Although the variations of Ed with coverage for the ZSM-5 and ZSM-11 zeolites in fig. 1 and 3 are different, the values of Ed are not greatly different. By both methods Ed is ca. 105 kJ mol-' at low coverages, decreasing to 60-70 kJ mol-' at high coverages for ZSM-5. The v.h.r. method shows an increase in Ed to a maximum value of 120 kJ mol-1 at 2 molecules per unit cell, whereas Ed steadily decreases with increasing coverage by the s.h.r.method. Richards and Rees' found that E d varied from ca. 75 to 78 kJ mol-' over the same coverage range for a H-ZSM-5 sample with Si/Al= 132. The large difference of ca. 20 kJ mol-' in Ed at low coverage between the high- and low-Si/Al ZSM-5 samples undoubtedly arises from the specific interaction of the n-bonds in the benzene ring and the electric fields around the A1 sites in the low-Si/Al sample. With ZSM-11 the v.h.r. method tends to give a constant value of E d of ca. 100 kJ mo1-l at lower coverages, decreasing slightly to ca. 85 kJ mol-' at high coverages, whereas the s.h.r. method shows a steady decrease from ca.105 to ca. 70 kJ mol-'.3636 T.P.D. of p-Xylene from Zeolites 12C 80 - I 3 40 I 1 +I 1 r/l Q I 0 -40 - 80 0 1.0 2.0 3.0 4.0 5.0 coverage/molecules per unit cell Fig. 2. Activation entropy of desorption of p-xylene as a function of coverage (v.h.r. method of analysis; /? = 6 K min-l): a, ZSM-5; 0, ZSM-11 and 0, Theta-1. Table 2. Ed values for p-xylene desorption (in kJ mol-') ZSM-5 ZSM- 1 1 Theta-1 v.h.r. method 89 98 110 s.h.r. method 85 86 80 The variation in E d with coverage for Theta-1 by the two methods of analysis shown in fig. 1 and 3 are quite different. With the v.h.r. method Ed tends to increase with increasing coverage, whereas E d decreases with the s.h.r. method, although both methods tend to give similar Ed values at low coverages.The values of Ed averaged over the whole coverage range from the s.h.r. method of analysis are given in table 2, where they may be compared with the corresponding E d values from the v.h.r. method. The s.h.r. E d values are similar for the three zeolites and consistent with the peak temperatures in table 1 . -A$ values by the s.h.r. method in fig. 4 all increase with increasing coverage, consistent with the increase in ordering of the sorbate at higher loadings.L. F. Chen and L. V. C. Rees 3637 40 0 1.0 2.0 3-0 4 . 0 5.0 coverage/molecules per unit cell Fig. 3. Activation energy of desorption of p-xylene as a function of coverage (s.h.r. method of analysis; = 6 K min-I): 0, ZSM-5; 0, ZSM-11 and 0, Theta-1. 1 0 1.0 2-0 3-0 4.0 5-0 coverige/molemles per unit cell Fig.4. Activation entropy of desorption of p-xylene as a function of coverage (s.h.r. method of analysis; B = 6 K min-I): 0, ZSM-5; 0, ZSM- 1 1 and 0, Theta-1. Desorption of p-Xylene from ZSM-5 and ZSM-11 Zeolites with Different Si/Al Ratios The t.p.d. of p-xylene from ZSM-5 (15), ZSM-5 (49, ZSM-5 (78), Silicalite-1 (8200), ZSM-11 (15) and ZSM-11 (72.9) was studied. (N.b. the numbers in parentheses are the Si/Al ratios.) The peak temperatures, widths and maximum rates of desorption derived from these t.p.d. profiles are given in table 3. Table 3 clearly shows that as the Si/A1 ratio increases the temperature of peak I increases, while the temperature of peak I1 decreases. Secondly, as the Si/Al ratio increases the widths of peaks I and I1 decrease and the maximum rate of desorption increases.These results suggest that as the Si/Al ratio increases Ed should decrease, in good agreement with the Ed values derived from these profiles by the s.h.r. method3638 T.P.D. of p-Xylene from Zeolites Table 3. T.p.d. peak temperatures, peak widths and maximum rates of desorption of p-xylene @ = 10 K min-') peak temperature peak width maximum rate of desorption /K higher /K peak I peak I1 peak peak I peak I1 /mg s-' ZSM-5 (1 5) 320 405 I 300-370 370-470 4.202 E-03 ZSM-5 (45) 335 380 I 300-350 350-400 6.557 E-03 ZSM-5 (78) 340 380 I 310-350 350-400 8.621 E-03 Silicalite- 1 (8200) 342 370 I 315-350 350-390 9.453 E-03 ZSM-11 (15) 320 405 I 300-370 370-470 3.361 E-03 ZSM- 1 1 (72.9) 325 400 I 300-360 360440 6.942E-03 100 80 i P 60 q" 40 1:O 2:O 3-'0 4:O 5-'0 6:O coverage/molecules per unit cell Fig.5. Activation energy of desorption of p-xylene as a function of coverage for ZSM-5 zeolites with various Si/AI ratios (s.h.r. method of analysis; p = 10 K min-I): ., ZSM-5 (15); 0, ZSM-5 (45); 0, ZSM-5 (78) and 0, Silicalite-1 (8200). shown in fig. 5. In this figure E d is seen to decrease in all cases with increasing coverage, and also decreases significantly at low coverages with increasing Si/A1 ratio. However, the value of E d at high coverages tends to be the same for all samples at ca. 75 kJ mol-1 down to ca. 50 kJ mol-l. The value of Ed averaged over the whole coverage range from the s.h.r. method of analysis for all these zeolites is given in table 4, which also includes the saturation capacities of these samples for p-xylene.Table 4 shows that with both ZSM-5 and ZSM-11 as Si/Al increases the sorption capacity increases. As Si/Al increases the steric hindrance of the decreasing number of H+ and Na+ ions in the channels decreases and the micropore volume increases. This volume increase, however, is larger than the volume of the H+ and Na+ ions removed, suggesting that the cations in the channel network of ZSM-5 and ZSM- 1 1 reduce the packing efficiency of p-xylene molecules. The directing influence of the charge centres in the channel network does seem to beL. F. Chen and L. V. C. Rees 3639 Table 4. p-Xylene saturation capacities and Ed values obtained by s.h.r. method (B= 10 K min-l) amount adsorbed /molecules per Ed zeolite unit cell /kJ mol-' ZSM-5 (1 5) 5.95 86 ZSM-5 (45) 6.80 78 ZSM-5 (78) 6.89 76 Silicalite- 1 7.05 70 ZSM-11 (15) 6.28 87 ZSM-11 (72.9) 6.9 80 the reason for the decrease or disappearance of the hysteresis loop found in sorption/desorption studies of p-xylene in ZSM-5 as Si/Al decreases.Conclusion The analysis of t.p.d. profiles by the s.h.r. method seems to give more reasonable Ed and -AS$ values as a function of coverage. The use of the more exact solution of the temperature integral in eqn (1) in the s.h.r. method is probably the reason for the improved quality of the data. The results reported in this paper clearly show the decreased sorption capacity and higher activation energies for desorption of p-xylene in the channels of these three pentasil type zeolites as Si/Al decreases. The presence of charge centres in the channels reduce the packing efficiency of p-xylene sorbate molecules. References 1 R. E. Richards and L. V. C. Rees, Zeolites, 1986, 6, 17. 2 E. Dima and L. V. C. Rees, Zeolites, 1987, 7, 219. 3 G. T. Kokotailo, P. Chu, S. L. Lawton and W. M. Meier, Nature (London), 1978, 275, 119. 4 G. T. Kokotailo, S. L. Lawton, D. H. Olson and W. M. Meier, Nature (London), 1978, 272, 437 5 I. D. Harrison, H. G. Leach and D. A. Whan, Zeolites, 1987, 7 , 21. 6 J. M. Thomas and G. R. Millward, J. Chem. Soc., Chem. Commun., 1982, 1380. 7 S. A. I. Barri, G. W. Smith, D. White and D. Young, Nature (London), 1984, 312, 533. Paper 8/00473K; Received 29th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403633

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1988, 84(10), 3641-3647 Infrared Study of the Adsorption of C,,EJ on Silica immersed in Carbon Tetrachloride Shona Keith and Colin H. Rochester* Chemistry Department, The University, Dundee DDl 4HN Infrared spectra of C,,E, adsorbed on silica in carbon tetrachloride show that adsorption primarily involves hydrogen bonding between surface silanol groups and oxygen atoms in ethyleneoxy segments of C,,E, molecules. At low surface coverages every ethyleneoxy segment in each adsorbed molecule is involved in hydrogen bonding with a silanol group. The average number of segments in each adsorbed molecule which are bonded to silanol groups decreases as the coverage of available silanol groups increases. At full coverage ca. 3.5 oxygen atoms in each adsorbed molecule remain hydrogen bonded to the surface.The use of infrared spectroscopy to monitor the number points of attachment of adsorbed polymer molecules to an oxide surface was first reported by Fontana and Thomas.' Killmann et a1.2, developed methods involving infrared spectroscopy for the in situ study of a variety of polymers adsorbed on silica at the solid/liquid interface. The numbers of attached segments for poly(ethy1ene glycol) molecules were deduced from a combination of adsorption isotherm data and the changes in absorption intensity of infrared bands due to silanol groups on the silica surface. Using experimental techniques developed in our Lijour et al.7 have studied the adsorption of four octylphenolethoxylate non-ionic surfactants on silica immersed in carbon tetrachloride.The infrared spectra led to valuable information about the mode of adsorption and orientation of adsorbed surfactants containing ethyleneoxy segments in the hydrophilic component of the molecules. In 1986 we embarked on a project involving infrared study of surfactants adsorbed from aqueous solution on to silica. As a part of this work spectra have been recorded of 0-n-dodecylpentaethylene glycol (C,,E,) adsorbed from carbon tetrachloride on to silica. The appearance of the paper by Lijour et aZ.? prompts this report of our results for Cl,E5 which lead to complementary conclusions to those for the octylphenolethoxylates. Experimental Spectra of pressed discs of aerosil silica (176 m2 g-') immersed in solutions of C12E5 (Nikko Chemical Co) in carbon tetrachloride (Fisons spectrograde) were recorded using a Perkin-Elmer 1750 F.t.i.r.spectrometer (50 scans at resolution 1 cm-l) in conjunction with an infrared cell with silica optical windows.' The general experimental procedure was as before.' Discs were heated at 873 K for 1 h in vacuum before being cooled prior to immersion in carbon tetrachloride. All spectra are reported with respect to a background spectrum of the empty infrared cell. All spectra refer to systems at equilibrium for which there were no further changes in spectrum with time. Spectra of C12E, in carbon tetrachloride were recorded using a fixed-pathlength infrared cell with fluorite windows. For comparative purposes spectra were also recorded of dodecan-1-01 and poly(ethy1ene glycol) 200 (both B.D.H.) in carbon tetrachloride and adsorbed on silica immersed in carbon tetrachloride.t 0-n-dodecylpentaethylene glycol. 364 13642 C12E5 on Silica Results Spectra of C,,E, in solution contained bands at 2927 and 2856cm-l due to CH stretching vibrations. The maximum absorbance values of both bands were linear functions of C,,E, concentration up to 0.15 mol dm-, with therefore a constant value (1.44) for A,,,,/A,,,,. The intensities of two much weaker bands at 3604 and 3495 cm-' were also linear functions of concentration up to 0.1 1 mol drn-,. These bands are ascribed to OH stretching vibrations of hydroxy groups which are unperturbed and perturbed by hydrogen bonding, respectively. The constant intensity ratio with increasing concentration implies that the terminal OH groups in some C,,E, molecules in solution are involved in intramolecular hydrogen- bonding interactions with adjacent oxygen atoms in the same molecule^.^ However, for a C,,E, concentration of 0.15 mol drn-,, the constant intensity ratio was no longer maintained in the sense that the band at 3495 cm-' was more intense and that at 3604 cm-' was less intense than expected from the linear relationships between intensities and concentration up to 0.1 1 mol drn-,.The deviations would be consistent with the onset of dimerisation caused by intermolecular hydrogen-bonding interactions. Spectra of silica in carbon tetrachloride exhibited a single maximum at 3689 cm-' due to isolated surface silanol groups which were perturbed by contact with liquid carbon tetrachloride. Spectra of silica immersed in solutions of C,,E, in carbon tetrachloride are shown in fig.1. The diminution in band intensity at 3689 cm-' with increasing C,,E, concentration and the concomitant appearance of a broad maximum at 3360cm-' (shifting to 3290 cm-' with increasing surface coverage) is consistent with the formation of hydrogen bonds between surface silanol groups and oxygen atoms in C,,E, molecules. The overall shape of the band envelope in the spectral range 2800-3000 cm-' was very similar to the shape of the corresponding band envelope for solutions of C,,E, in carbon tetrachloride. However, a small difference was that a distinct maximum at 2874 cm-l (fig. 1) replaced a shoulder of similar intensity at ca.2873 cm-' in the solution spectra. The spectral features in fig. 1 ascribed to vibrations of C,,E, molecules were entirely due to C,,E, adsorbed on silica and not to C,,E, in solution. Replacement by flushing of the solution phase by pure carbon tetrachloride produced no change in the intensities of the bands due to CH stretching vibrations. Bands due to C,,E, in solution made a negligible contribution to the recorded spectra (fig. I), suggesting that the equilibrium concentration of C,,E, in solution was low and that C,,E, was fairly strongly adsorbed on silica in carbon tetrachloride. The latter conclusions were substantiated by the lack of desorption of C,,E, on flushing the surface with pure carbon tetrachloride. The overall band envelope arising from vibrations of OH groups must contain contributions from bands due to OH stretching vibrations of C,,E, in the adsorbed state.These would be expected at 3604 cm-' for unperturbed OH (cf. C,,E, in solution) and ca. 3460 cm-l for a hydroxy group perturbed by hydrogen-bonding interactions with surface silanol groups. lo Under the conditions of the present experiments these bands are comparatively weak and the spectra are dominated by the much more intense maxima due to unperturbed and perturbed silanol groups. Fig. 2 demonstrates the inter-relationships between the intensities of the four dominant maxima in the infrared spectra of C,,E, on silica. The linear relationship between the absorbance values A,,,, and A,,,, for the maxima due to unperturbed and perturbed silanol groups, respectively, is typical of results for adsorption processes involving the formation of hydrogen bonds between silanol groups and adsorbed molecules.11 The non-linear plots of band intensities at 2927 and 2856cm-' against absorbance at 3290 cm-l show that with increasing adsorption a decreasing number of silanol groups were perturbed by each adsorbed C,,E, molecule. Thus the numbers of ethyleneoxy segments in each adsorbed C,,E, molecule which were perturbed by hydrogen-bonding interactions with surface silanol groups decreased with increasing2.2 8 -E 1.1 4 -2 c S.Keith and C. H. Rochester I I I 1 I w avenurnber/cm-' 3643 Fig. 1. Spectra of silica immersed in solutions of increasing concentrations of C,,E, in carbon tetrachloride. I I I 2.C 1.5 8 B -e 1.c 2 % 0.5 C absorbance at 3290 err? Fig.2. Absorbance values at (a) 3689, (b) 2927 and (c) 2856 cm-' as a function of absorbance at 3290 cm-' for silica in solutions of C,,E, in carbon tetrachloride.3644 C12E5 on Silica fractional coverage of the available silanol group adsorption sites. The average numbers of ethyleneoxy segments bonded to the surface have been estimated from the spectroscopic data as a function of the coverage of adsorption sites. The total number of surface silanol groups in a 2.54cm diameter silica disc was calculated from the weight of the disc, the surface area of the silica and the surface population of OH groups. The latter was taken to be 1.7 nmP2 for aerosil silica which had been heated at 873 K.12 Combination with the fractions foH of silanol groups perturbed by C12E5 deduced13 from absorbance values at 3689 cm-' gave the numbers no, of silanol groups in the disc perturbed by the adsorption of C12E5.The number n, of molecules of surfactant adsorbed in the disc was calculated from integrated absorption intensities for the band envelope at 2600-3050 cm-l due to CH stretching vibrations of C12E5. It was assumed that the integrated extinction coefficient was the same for C,,E5 in solution in carbon tetrachloride as for C12E5 adsorbed on silica immersed in carbon tetrachloride. For C12E5 in solution at concentration [C12E5] mol in a cell of pathlength 1 cm (calibrated using interference fringes) the integrated extinction coefficient is related to the integrated absorbance (in absorbance-wavenumber units) by the equation 3050 3050 A,dv = [C12E5] E , ~ V J2600 in which A , and E, are the absorbance and extinction coefficient, respectively, at wavenumber V.The corresponding equation for a silica disc of radius rcm and containing n, molecules of surfactant is given by 3050 3050 12600 I 6 0 0 A,dv = (n,/N, nr2) E,dv in which NA is Avogadro's number. Integrated absorption intensities were deduced using the dedicated Perkin-Elmer F.t.i.r. computer against a baseline which for the disc involved, taking into account the low-wavenumber tail of the broad band due to perturbed silanol groups. The resulting values of n, were combined with noH to give perturbation numbers p = (n,,/n,), which are the average numbers of surface silanol groups which were perturbed by each adsorbed surfactant molecule.The variation of p with fractional coverage of available silanol group sites is shown in fig. 3 for four separate series of adsorption experiments. The error in measuring p was greatest at low surface coverages. However, the results clearly established that p was ca. 6 at low coverages, indicating that all the six oxygen atoms in each C12E, molecule were involved in hydrogen-bonding interactions with surface silanol groups. The p value fell to ca. 3.5 asf,, -+ 1. The development of a maximum at 2874cm-l for C12E, adsorbed on silica was accompanied by a second detectable effect in which a pronounced shoulder at 2953 cm-' for C12E5 in solution became a less distinct shoulder at slightly lower wavenumbers (Av< 1 cm-l) for adsorbed C12E5.Attempts were made to analyse quantitatively the differences between the spectra of C12E5 in solution and in the adsorbed state. However, the spectral changes were too small to allow reliable meaningful difference spectra to be obtained. Despite this, comparison of the spectra of C12E5, dodecan-1-01 and PEG 200 in solution and adsorbed on silica immersed in carbon tetrachloride suggested that the spectral changes observed for C12E5 could be attributed to perturbation of ethyleneoxy segments in the hydrophilic head groups of adsorbed molecules. Spectra of C,,E, in solution could be replicated by addition, with appropriate weighting factors, of the separate solution spectra of dodecan- 1-01 and PEG 200. This enabled the band envelope for C,,E, to be resolved into the contributions due to the C,, alkyl chain and the E, head group.Spectra of dodecan- 1-01 on silica at a coverage foH = 0.4 of silanol group sitesS. Keith and C. H. Rochester 3645 a 6 P 4 2 C I I I I I 0 O O 0 0 0 0 0 00 I 1 I I 1 0.2 0.4 0.6 0.8 1.; f O H Fig. 3. Numbers, p , of silanol groups perturbed by each adsorbed C,,E, molecule as a function of the fraction, foH, of surface silanol groups perturbed by hydrogen bonding to C,,E,. contained bands at similar positions to those in spectra of dodecan-1-01 in solution. Spectra of PEG 200 in solution exhibited a band envelope with a main maximum at 2873 cm-' and subsidiary maxima which were scarcely more than shoulders at 2910 and 2943 cm-'. The adsorption of PEG 200 on silica resulted in shifts in the positions of the bands at 2873 and 2943 cm-' to higher and lower wavenumbers, respectively. Similar shifts in the component of the Cl,E5 spectra due to ethyleneoxy segments would be consistent with the overall observed changes in the spectra accompanying the transfer of C1ZE5 from solution to the adsorbed state.Discussion The present results for Cl,E5 support previous conclusions for poly(ethy1ene glycols)14 and octylphenolethoxylates7 that ethylenoxy segments in molecules adsorbed on silica immersed in carbon tetrachloride form hydrogen bonds with isolated surface silanol groups, Average numbers of silanol groups which were perturbed by each adsorbed octylphenolethoxylate molecule were deduced' using spectroscopic data alone, as opposed to an alternative method which requires knowledge of the adsorption i~otherm.~ The extinction coefficients of bands at 2960 and 2880 cm-l were assumed to be unchanged by transfer of octylphenolethoxylates from solution to the adsorbed state.This assumption was deemed valid for the band at 2960 cm-l but was thought to be not strictly true at 2880cm-l because the latter band exhibited a decrease in relative absorption intensity when adsorption took place. This change was attributed to perturbation of ethyleneoxy groups.' Similar effects were observed in the present study. Values of the p factor deduced using absorbance values for the absorbance maximum at 2927 cm-l were in good agreement with the corresponding values (fig. 3) deduced using integrated absorption intensities of the overall band envelope between 2600 and 3050 cm-l.However, p values deduced from absorbances at 2856 cm-l were slightly higher than those from the other two methods. This discrepancy arises because of the shifting position on adsorption of C,,E, of an underlying band due to vibrations of ethyleneoxy groups. The integrated absorption intensity procedure has the advantage that theC12E5 on Silica intensities of shifting bands are retained in the total intensity of the overall band envelope. An assumption inherent in the procedure is that shifts in the positions of individual bands which contribute to the overall band envelope are not accompanied by significant changes in integrated absorption intensities of the individual bands. This assumption is probably valid in the present context since the CH band shifts induced by adsorption of C12E5 were very small.Infrared spectra of ethyl ethanoate15 and cyclohexanone16 on silica immersed in aprotic solvents established that adsorption involved either pairs of surface silanol groups simultaneously forming two hydrogen bonds to a single carbonyl group in each adsorbed molecule or single silanol groups each forming a hydrogen bond with one carbonyl group. Isolated silanol groups on silica which had been heated at high temperatures were too far apart from each other to interact in pairs with single carbonyl groups.15 By analogy it is highly unlikely that pairs of silanol groups on silica which had been preheated at 873 K could form pairs of hydrogen bonds with individual oxygen atoms in adsorbed C1zE5.A p value of 6 for C12E, at low foH values therefore implies that the six oxygen atoms in an adsorbed C12E5 molecule are each bonded to the silica surface via formation of a hydrogen bond with one isolated silanol group. The conclusion that the ethyleneoxy E, chain lies flat on the silica surface at low surface coverages is identical to the corresponding conclusion for octylphenolethoxylates adsorbed on silica in carbon tetrachloride. Decreases in the number of points of attachment of ethyleneoxy chains to silica surfaces with increasing surface coverage have previously been observed by infrared spectroscopy for adsorbed polyethylene glycols,2 nonylphen~lethoxylates~~. l8 and octylphenolethoxylates. In the latter case surfactants with average numbers of oxygen atoms per molecule of 2.03, 10.7 and 68.3 gave average numbers of adsorbed segments at saturation surface coverage (foH -+ 1) of 2, 5 and 32, respectively, which compared with corresponding figures of 2.2, 10 and 66 at low coverages (f,, -+ O).’ The results for C,,E5 with six oxygen atoms and p falling from 6 at foH + 0 to 3.5 at foH -+ 1 therefore fits into exactly the same pattern of behaviour.Previous spectra of polyethylene glycol^'^ and octylphenolethoxylates7 on silica showed no obvious shifts with increasing surface coverage in the position of the infrared band due to silanol groups which were perturbed by the formation of hydrogen bonds with adsorbed molecules. Furthermore, for poly(ethy1ene glycols) the intensities of the band were linear functions of the adsorption enthalpies suggesting that the average interaction energy per perturbed silanol group was independent of surface coverage.In contrast, several series of spectra for C12E, at increasing coverages on silica consistently showed a trend in the infrared band maximum from 3360 to 3290cm-l. A plausible conclusion would be that the average strength of the hydrogen bonds between silanol groups and adsorbed Cl,E, increases with increasing surface coverage. Two explanations of this effect can be envisaged. First, C12E, adsorbed in the flat configuration ( p = 6) is not displaced or reorientated ( p < 6) with increasing coverage. Residual adsorption sites do not allow incoming molecules to adopt the flat configuration, and the number of points of attachment of incoming molecules will decrease, eventually to one as f,, -+ 1.Molecular models suggest that the geometry of the flat configuration is not ideal for the maximum possible strength of interaction between every silanol group and the oxygen atom to which it is bonded. An incoming molecule at high coverages which, say, only interacts with a single silanol group will be able to adopt a configuration which gives the strongest attainable hydrogen bond, Diethyl ether, for example, adsorbed on silica in carbon tetrachloride” gives an SiOH band shift to 3230 cm-’, which corresponds to a stronger hydrogen bond than the average bond strength for Cl2E, as fOH -+ 1 . This explanation implies that the p value of 3.5 at high coverage represents the average of a wide range of p values (from 6 possibly to 1) corresponding to molecules cohabiting the surface in quite different configurations.The second explanation would involveS. Keith and C . H. Rochester 3647 reorientation of adsorbed C,,E5 molecules in order to accommodate incoming C,,E, molecules in the same configuration on the surface. Desorption of some ethyleneoxy segments in each adsorbed molecule would again favour the attainment of stronger individual hydrogen-bonding interactions and an enhanced shift in the position of the infrared band assigned to perturbed silanol groups. The average p value on this model would result from a narrow range of individual p values. On balance this second explanation would probably be more consistent with the average p values asfoH + 1 for C,,E, and octylphenolethoxylates all being equal to about half the total number of oxygen atoms in the surfactant no matter whether the latter was 6 (for C12E5), 10.7 or 68.3 (for octylphenolethoxylates7). It is most likely that this is due to desorption of alternate ethyleneoxy segments asfoH --+ 1 rather than desorption of a tail of half the ethyleneoxy chain leaving a train of equal length flat on the surface.We thank the S.E.R.C. for a studentship, Unilever Research for collaboration through the CASE scheme and Dr J. W. Mactaggart for helpful discussions. References 1 B. J. Fontana and J. R. Thomas, J . Phys. Chem., 1961, 65, 480. 2 E. Killmann, H. J. Strasser and K. Winter, 6th Znt. Congress on Surface Active Agents, Zurich, 1972 (Carl Hanser Verlag, Munich, 1973), p. 221. 3 E. Killmann, M. Korn and M. Bergmann, in Adsorption from Solution, ed. R. H. Ottewill, C. H. Rochester and A. L. Smith (Academic Press, London, 1983), p. 259. 4 K. Marshall and C. H. Rochester, J . Chem. SOC., Faraday Trans. 1, 1975, 71, 2478. 5 C. H. Rochester, Progr. Colloid Polym. Sci., 1980, 67, 7. 6 C. H. Rochester, J . Oil Colour Chem. Assoc., 1985, 285. 7 Y. Lijour, J-Y. Calves and P. Saumagne, J . Chem. SOC., Faraday Trans. 1, 1987, 83, 3283. 8 C. H. Rochester and G. H. Yong, J . Chem. SOC., Faraday Trans. 1, 1980, 76, 1466. 9 C. H. Rochester, in The Chemistry of the Hydroxyl Group, ed. S . Patai (Interscience, London, 1971), p. 327. 10 K. Marshall and C. H. Rochester, Faraday Discuss. Chem. SOC., 1975, 59, 117. 11 C. H. Rochester, Adu. Colloid Interface Sci., 1980, 12, 43. 12 V. Y. Davydov, A. V. Kiselev and L. T. Zhuravlev, Trans. Faraday SOC., 1964, 60, 2254. 13 C. H. Rochester and D-A. Trebilco, J . Chem. Soc., Faraday Trans. 1 , 1977, 73, 883. 14 E. Killmann and H-J. Strasser, Angew. Makromol. Chem., 1973, 31, 169. 15 S. N. W. Cross and C. H. Rochester, J . Chem. SOC., Faraday Trans. 1 , 1979, 75, 2865. 16 C. H. Rochester and D-A. Trebilco, J . Chem. Soc., Faraday Trans. 1 , 1979, 75, 221 1. 17 H. Rupprecht and H. Liebl, Kolloid Z . , 1970, 239, 685. 18 H. Rupprecht, Archiu. Pharm. (Weinheim), 1972, 305, 149. 19 D. M. Griffiths, K. Marshall and C. H. Rochester, J . Chem. SOC., Faraday Trans. 1, 1974, 70, 400. Paper 8/00554K; Received 1 lth February, 1988

ISSN:0300-9599    

DOI:10.1039/F19888403641

出版商:RSC    

年代:1988

数据来源: RSC