摘要:

Transport in Aqueous Solutions of Group IIB Metal Salts (298.15 K) Part 4.4nterpretation and Prediction of Isotopic Diffusion Coefficients for Cadmium in Dilute Solutions of Cadmium Iodide BY RUSSELL PATERSON" AND LUTFULLAHJf Department of Chemistry, University of Glasgow, Glasgow G12 SQQ Receiued 6th May, 1977 A method for predicting isotopic diffusion coefficients for component ions in a self-complexed electrolyte has been developed. In an irreversible thermodynamic analysis the diffusion coefficient has been shown to be a function of the mobility coefficients of the free ion and its complexes together with the interionic coupling coefficients between labelled and unlabelled species. It is shown that these coefficients may be evaluated by an extension of the methods developed earlier to predict isothermal transport in such systems.Cadmium ion diffusion coefficients in aqueous cadmium iodide, Das, have been predicted (0.003-0.10 mol dm-3). These describe the maximum anticipated experimentally. Observed and calculated D, agree to within 3 % at 0.1 mol dm-3. In Paper 3,' measurements of the isotopic diffusion coefficients of cadmium-1 15 in aqueous cadmium iodide were reported. The variation of the diffusion coefficient of cadmium, D,,, with increasing concentration of salt was anomalous and had no parallel in the literature on isotopic diffusion of ions in aqueous media, fig. 2. and in earlier papers the isothermal transport properties of the unlabelled solution were reported and interpreted in terms of mobility and coupling coefficients between free and complexed ions in the solutions, using irreversible thermodynamic analyses.From the data already available it is obvious that the concentration dependence of D,, is one more consequence of self-complexing in this system. the diffusion coefficient of an isotopically labelled ion, a*, in an aqueous solution of salt (a, b) was represented by eqn (1). Aqueous cadmium iodide is extensively complexed In earlier discussions,4* D,, x loy3 = RT(L,JC,T- C,TL,*,/C,OCX) where C,' is the total molar concentration of ion a and C:, Cz are the concentrations of unlabelled and labelled species, respectively, (C,' = C:+C$). L, is the direct mobility coefficient of the ion (a) in the unlabelled binary solution (a, b/water). In eqn (1) D,, is expressed as cm2 s-l when RT, La, and Lza and concentrations are expressed in their usual units, J mol-l, mo12 J-1 c111-l s-l and mol dm-3, respectively.I,,, and the other mobility and coupling coefficients of dilute solutions of cadmium iodide have been predicted.2 It is observed that in solutions, below 0.1 mol dm-3, D,, passes through a maximum while L,,IC,T shows a minimum over this same concentration range. It is, therefore, obvious that the concentration dependence of D,, is largely dependent upon that of the isotope-isotope coupling term in eqn (1). It is the aim of this paper to obtain an explicit expression for this isotopic term as a function of coupling coefficients between labelled and unlabelled cadmium species t Present address : Institute of Physical Chemistry, University of Peshawar, N.W.F.P.Pakistan. 103104 TRANSPORT IN AQUEOUS SOLUTIONS in solution. Thereafter these interionic coupling coefficients may be evaluated, using the classical theories of Onsager, expressed in macroscopic irreversible thermodynamic form by Pikal.6 THEORETICAL Before developing an expression for isotopic diffusion in terms of the complex species present it is necessary to define nomenclature and symbols. As before,2 the complexes CdI,"-x are identified by the value of x, thus, for example, CdIZ- will be species 4. The free aqueous ions Cd2+ and I- will be denoted by a and b respectively. In any solution of cadmium iodide the total concentration of cadmium, CT, and iodide, Cz will be given by eqn (2) and (3). and (from the stoichiometry o f the salt Cz = 2C3. When a proportion of the normal unlabelled cadmium is removed from solution and replaced by an equal amount of isotopically labelled cadmium, the total concentra- tions of all species will remain unaltered, but each will now contain both labelled and unlabelled components.For any species i (i = a, 1,2,3,4), the total concentra- tion, Ci, will be the sum of concentrations cPand cT ; those of unlabelled and labelled respectively, eqn (4) The specific concentration of the labelled and unlabelled cadmium in each solution species are defined by p: and pf respectively, eqn (5) C,' = Ca+C,+Cz+C,+C4 (2) ct = C1+2Czf3C3+4C4+Cb (3) Ci = cP+cT i = a, 1,2,3,4. (4) p; = c;/C, and po = colCi i = a, I, 2 , 3 , 4 (5) (p"ppp = 1). It will be assumed that the labelled isotope is chemically identical to the normal bulk isotopic mixture of cadmium in the unlabelled solution.In consequence, there can be no isotopic enrichment of labelled species in any one of the complexes. This is a widely held assumption in diffusion studies and essentially equates isotopic diffusion as being equivalent to self-diffusion. In the present context it requires that the specific concentrations, p:, for all species involving cadmium are equal, and so the subscript i will be dropped in all further discussion. Similarly pp becomes po. The total concentration of labelled cadmium, Cz is therefore related to total cadmium concentration, C: by eqn (6) A The total contribution of unlabelled cadmium, Ct, is given by an expression analogous to eqn (6).A similar scheme for defining fluxes is required. Jay J b are the sums of the molar fluxes of all cadmium and all iodide species in solution respectively, eqn (7) and (8). Ja = ja +j, +J2 +j3 +j4 Jb = jb+jl f2j2 +3j3 +4j4. (7) (8) andR . PATERSON A N D LUTFULLAH 105 When a proportion of the cadmium is labelled, the total flux of any species will remain unaltered (by the assumption of the chemical identity of isotopes) but will now be the sum of labelled j : and unlabelled jp fluxes, eqn (9) ji = jP+j:, (9) The total flux of labelled cadmium species, J:, is the only flux to be measured experi- mentally thus : 4 J,* = jT (10) i = a and in any experiment J , = J,* + J," where J: is the net fiux of the unlabelled component. For a species i, the thermodynamic driving force which causes a direct flow is (-grad f i r ) ; the negative gradient of the electro-chemical potential of that species at any point in the non-equilibrium system, represented as xf, eqn (12) x f = -grad /it = -RTdln Cf/dx-RTd lny,/dx+Z,F(-d$/dx) (12) where Zi is the valency, including sign of the species i, and yr the molar activity coefficient.When the species is neutral, the valency is zero and the force is simply the negative gradient of chemical potential. When labelled and unlabelled species i are introduced with forces xp and $, respectively then :' (13) For unlabelled species the expression for xp is formally identical, except that the concentration and activity coefficient cp and yp replace those of the labelled species in eqn (13).Since the labelled and unlabelled species may be assumed to have activity coefficients which are equal to one another and to those of the total species in a local volume element and since dCi = dc? + dcf, from eqn (4), then from eqn (1 2) and (13), Under the special condition of isotopic diffusion, in which no bulk gradients of chemical potential exist and the sole source of non equilibrium is due to isotopic gradients; xi = 0 and so, from eqn (14), Equally there are no gradients of activity coefficients or of electrical potential ( ~ ) and so the forces xp and xT are simple functions of concentration gradients, defined (16) x? = -grad ,!iT = - RT d In cT/dx - RT d In yf/dx + ZiF( - dt,b/dx). cixi = cPxP+c'xf or x i = poxP+pox*. (14) pox; = -p*$. (15) by eqn (16) xr = - (RT/cT)( - dcf/dx) superscript F = O or *.IRREVERSIBLE THERMODYNAMIC ANALYSIS In Paper 2 the binary coefficients of cadmium iodide were expanded in terms of mobility and coupling coefficients of the uncomplexed and complexed species present in solution. For any species k, linear phenomenological equations were written, eqn (17). A106 TRANSPORT I N AQUEOUS SOLUTIONS Eqn (17) deals with an experimental conditions likely to be obtained in normal studies and it is assumed here and in all subsequent discussion that the Onsager Reciprocal Relations (O.R.R.) are obtained and so Iki = Zkf. If it is assumed that the experimental conditions to which eqn (17) refers are again obtained with the same solution in which now labelled cadmium isotope is present, then, from eqn (9), the sum of the flows of labelled and unlabelled isotopic species, k, are equal to the net flow, jk ; ( j , = jkq+jz).The phenomenological equations for isotopic flows will define a symmetrical 11 x 11 matrix of mobility coefficients, since five of the six flows (and six forces) of eqn (17) are now subdivided into two components. The phenomenological equations of the isotopic system are given by eqn (1 8)-(20) 4 j," = (jkiXp+?ki*$)+~k,~, ( k = a, 1, 2, 3, 4) (18) i = a and for b, which was unlabelled, From eqn (9), (18) and (19), A or, using eqn (14), Comparing terms in eqn (17) and (22), having noted that lkfxi = &@oxp+~~fp*x~, from eqn (14) then, and lki = (jki+?k*i)/po = ( j k i + + j k * i * ) / P * i and k = a, 1, 2, 3, 4 (23) (24) Ibb = 'bb (25) lbi = &i/Po = ?bi*/P* i = a, 1,2,3,4.(26) (27) zkb = (jkb+jk*b) k = a, 1, 2, 3, 4. Since the force on iodide, x b , is unaltered, under conditions when only cadmium species are labelled, then from eqn (14), (17) and (20), and From eqn (24) and (26), using the Onsager Reciprocal Relations, lkb = j,b/po = jk*b/p*. Eqn (23), (25) and (27) establish specific relationships between the mobility coefficients of the unlabelled system, eqn (17) and those of the isotopic matrix, eqn (18)-(20). We must now formulate expressions for the total flow of labelled cadmium .I,*, using eqn (lo), under the more restricted conditions of isotopic diffusion, in which no bulk chemical potential gradient exists and so all xi and j , are zero in eqn (17).R. PATERSON AND LUTFULLAH 107 Under these conditions xb, the force on iodide ion, is also zero and hence from eqn (lo), (18) and (19), Since, from eqn (14), p*xT = -pox; when xi = 0, and from eqn (23) From the analysis of the concentrations of complexes to the binary coefficients, Paper 2,2 A d where La, is the direct mobility coefficient of cadmium in the binary solution for which the phenomenological equations are : and Recalling the isotope-isotope terms in eqn (28), expansion of the summation shows that, when the Onsager Reciprocal Relations are assumed, eqn (30) is obtained.(30) Ja = Laaxa $- la bXb J b = LbaXa LbbXb. 4 4 4 4 (jki*ip*+jk*i/po) = x C (Iki*/Po P*)* k = a i = a k=a i = a From eqn (28)-(30), 4 4 J,* = RT[L,,- k=a i=a 1 (lki*/pop*)](-dp*/d.). When Cz, total cadmium concentration, is constant at all points in the solution (as it must be for isotopic diffusion) then from eqn (6), and eqn (31) becomes, dp*/dx = (l/Cz)dC:/dx (32) Eqn (33), therefore, has the form of Fick's first law of diffusion, eqn (34), The flow JZ is that measured by experiment and thus Da, is the isotopic diffusion coefficient obtained from such experiments.Comparison of eqn (33) and (34) gives, J,* = Daa( - dC,*/dx). (34) r 4 4 1 Eqn (35) and (1) reveal that expressions for isotopic diffusion coefficients in a complex or a simple dissociated electrolyte have identical forms. It may be noted that, from eqn (5), the term C,TLa,*/(C,OC,*) in eqn (1) is equal to (l/C,T)(Laa*/P*p*) making the108 TRANSPORT IN AQUEOUS SOLUTIONS identity more obvious.The isotope-isotope coupling term, which is a function of a single coupling coefficient Laas in eqn (l), is now replaced by a summation of coupling coefficients which include all possible interactions between labelled and unlabelled foriix of the cadmium containing species in the complexed solution. PREDICTION OF ISOTOPE-ISOTOPE COUPLING AND DIFFUSION COEFFICIENTS FOR CADMIUM As discussed earlier in relation to eqn (l), it is the complex variations in the function C,’L,,*/C,”C,* which determines the major trends in the concentration dependence of Daa. The direct mobility Laa/C: shows quite the inverse concentration dependence of D,, [fig. 1 of ref. (3)J In the theoretical section it was shown that the isotope-isotope coupling contribu- tion to Daa may be expressed as a summation, by comparing eqn (1) and (35), c:L,,*/C:c: = (L,,,lpop*)/c: 4 1 FikalY6 in his analysis of coupling coefficients, has used the Limiting Laws of Fuoss and Onsager to obtain a general expression for such coefficients, eqn (37), 10 2, /\/Cx = A,/p ppt [ B]l’.(37) The units of are those used in this study (moI2 J-l cm-l s-l) and concentrations are expressed in molar units. A is a constant (0.107 40), [B] is a combination of relaxation and electrophoretic terms, discussed below and I is the true ionic strength of the solution, eqn (38), (38) 11: is defined by Pikal as the ionic strength fraction, eqn (39), where the superscript f- may be either O or * for unlabelled and labelled components respectively : (39) Thus in eqn (37), I = 3 2 (cj”+cri;)Zj” ( j = a, 1,2, 3,4, b).pJ = cjzj/c (cg+cj*)zj = c;zj/zr. dppp: = JcpcE 1Zizk[/21. (40) From eqn (37), (40) and ( 5 ) , eqn (41) is obtained, which has the same mathematical form as eqn (35) The term [B] used in these equations is defined by Pikal, eqn (42) where terms in Aik* and Bo involve the relaxation and electrophoretic effects respect- ively. In eqn(42), x = &‘&‘/lZiZkl. It is assumed from the postulatedidentical chemical characteristics of labelled and unlabelled species that Aj’ = A;* where A; is the equivalent conductance of the ion at infinite dilution, and y = E/ijA;/lZjl, where the summation is over all species, labelled and unlabelled, in the solution. From eqn (39) and (4), [B] = [ x / v Aik*-(BO/2) z,&] (42)R. PATERSON AND LUTFULLAH 109 Pikal has shown that Aik* = alZiZkl where a at 298.15 K is a constant, 0.229 62.Similarly Bo is a constant, 60.495. Thus eqn (42) becomes, Combination of eqn (37) and (44) gives : (relaxation term) (- lzicil I z k c k l (ZiZk)AB,/4)I-'. (45) (electrophoretic term) It is to be noted that all terms on the right hand side of eqn (45) may be evaluated without a knowledge of the absolute concentration of isotopically labelled cadmium. Only total concentrations and equivalent conductances at infinite dilution of free and complexed ions need be known. These have been evaluated in a similar but independent analysis of the transport properties of the unlabelled salt.2 Pikal's (S.L.L.) analysis is based upon the Limiting Laws and so is precise only in very dilute solutions.Equally, since it deals with interionic coupling coefficients, there can be no satisfactory method for estimating coupling between the neutral complex (2 or 2*) and its environment of ions. This means that all Z2j* and Z2+j are indeter- minate and must be set equal to zero. This limitation was also encountered in the analysis of the unlabelled salt.2 Such omissions may cause little error however, since in dilute solutions, neutral complex amounts to no more than 5 % of the total cadmium present .z In Paper 2,2 concentrations of free and complexed ions have been obtained as a function of total salt concentration. In that paper also an optimisation procedure for estimating the 1; for complexed ions was reported. Using these data, the isotope-isotope coupling terms of eqn (36) were calculated in the concentration range C:, (0.003-0.1 mol dm-3).For clarity eqn (36) has been expanded [eqn (46)] and the individual coefficients displayed in this format in table 1. In eqn (46) the parameters lik*/pop" are represented as (jk::) 4 4 1 (~ki*/Pop*>/c~ k=a i = a = 1ICi [(aa*) + (a1 *) + (a2*) + (a3*) + (a4*) + (1 a*) + (1 1 *) + (129 - (1 3*) + (14*) + (2a*) + (21 *) + (22*) + (23") + (24*) + (3a*) +(31*) +(32*) +(33*) +(34*) + (4a*) + (41 *) + (42*) - (43*) + (44*)]. (46) Since lik*/pop* equals lki*/pop*, by the Pikal analysis, eqn (45), the matrix of coefficients is symmetrical. Electrophoretic terms in eqn (45) dominate, and so for ions of like charge lik*/pop* will be negative, since (Z,Z,) is positive. These negative terms appear at the upper left and lower right hand corners of the matrix, eqn (46).For the same reasons coefficients between ions with opposite charges are positive and these appear110 TRANSPORT IN AQUEOUS SOLUTIONS on the upper right and lower left corners, eqn (46). Representative points are shown in table 1 and a complete set of predictions over the full concentration range are given in table 2. From table 1, it is observed that in the most dilute solutions, where higher (negatively charged) complexes are not present in significant proportions, the negative contributions to (Laa*/pop*)/C~ are largely [(aa*) + 2(al*) + (1 1 *)I. Only at the highest concentrations do the negative-to-negative ion interactions, [(33 *) + 2(34*) + (44*)] contribute significantly.Positive contributions to the total isotope-isotope coupling term are solely 2[(3a*) + (31 *) + (4a*) + (41 *)] and these are important only when the higher negatively charged complexes (3 and 4) are significant. ( -La,*/pop*)/C,T increases rapidly from its value of zero at infinite dilution largely because the only significant cadmium species are Cdlf (1) and cadmium ion itself (a). As complexing increases with further TABLE 1 .-COMPONENT ISOTOPIC MOBILITY COEFFICIENTS (ik*) AS DISPLAYED AND DEFINED BY EQN (46) AT REPRESENTATIVE TOTAL CONCENTRATIONS OF CADMIUM IODIDE, C:, (mol dm-3) [as in eqn (45) each coefficient must bemultiplied by the factor x to obtain the individual -7.405 97 (-4) -1.539 19 (-4) 1.560 66 (-6) 2.273 28 (-7) - - 1.133 92 (-3) - 3.387 95 (-4) 7.823 27 (-6) 1.81425 (-6) - - 1,873 03 (- 3) - 8.736 12 (-4) 5.762 32 (- 5) 2.453 72 (- 5) - - 4.057 79 (- 3) - 3.095 45 (- 3) 7.723 56 (-4) 7.470 19 (-4) - - 6.399 45 (- 3) - 5.407 69 (- 3) 2.060 94 (- 3) 2.687 05 (-3) - zik* in units mo12 J-' cm-l s-'1.C,'= 0.003 -1.539 19 (-4) - 1.560 66 (-6) -3.165 54 (-5) - 3.484 24 (- 7) 3.484 24 (-7) - -1.551 38 (-9) 5.053 51 (-8) - -2.416 24 (- 10) - - - C,' = 0.005 -3.387 95 (-4) - 7.823 27 (-6) -1.001 97 (-4) - 2.507 70 (-6) 2.507 70 (-6) - -2.582 72 (-8) 5.79094(-7) - - 6.394 98 (-9) - - - c,T = 0.01 -8.736 12 (-4) - 5.762 32 (- 5) -4.034 15 (-4) - 2.880 00 (- 5) 2.880 00 (-5) - - 8.593 36 (- 7) 1.221 26 (-5) - - 3.901 97 (- 7) - - - C z = 0.03 - 3.095 45 (- 3) - 7.723 56 (-4) -2.337 11 (-3) - 6.323 72 (-4) 6.323 72 (-4) - - 7.000 18 (- 5) 6.090 37 (-4) - - 7.232 52 (- 5) - - - C: = 0.05 -5.407 69 (-3) - 2.060 94 (- 3) -4.519 86 (-3) - 1.875 16 (-3) 1.875 16 (-3) - - 3.050 85 (-4) 2.434 15 (-3) - -4.264 28 (-4) - - - 2.273 28 (- 7) 5.053 51 (-8) -2.416 24 (- 10) -3.733 35 (- 11) - 1.814 25 (-6) 5.790 94 (-7) - 6.394 98 (- 9) -1.571 23 (-9) - 2.453 72 (- 5) 1.221 26 (- 5) - 3.901 97 (- 7) - 1.758 45 (- 7) - 7.470 19 (- 4) 6.090 37 (-4) - 7.232 52 (- 5) -7.414 38 (- 5) - 2.687 05 (- 3) 2.434 15 (- 3) -4.26428 (-4) -5.910 51 (-4) -R .PATERSON AND LUTFULLAH 111 TABLE 1 .-corztd. C z = 0.07 -9.177 09 (-3) -7.930 51 (-3) - 3.744 54 (-3) -7.930 50 (-3) -6.774 58 (-3) - 3.494 12 (- 3) - - - - 3.744 54 (-3) 3.494 12 (-3) - -6.802 83 (-4) 5.811 05 (-3) 5.398 14 (-3) - - 1.135 48 (-3) c: = 0.09 -1.243 73 (-2) -1.069 78 (-2) - 5.755 54 (-3) - 1.069 78 (-2) -9.091 26 (-3) - 5.358 42 (-3) 5.755 54 (-3) 5.358 42 (-3) - - 1.154 53 (-3) 1.007 95 (-2) 9.341 12 (-3) - -2.180 00 (- 3) - - - - c: = 0.10 - 1.425 31 (-2) -1.217 49 (-2) - 6.871 61 (-3) - 1.217 49 (-2) -1.027 27 (-2) - 6.359 90 (-3) 6.871 61 (-3) 6.359 90 (-3) - - 1.419 32 (-3) 1.263 72 (- 2) 1.164 22 (-2) - -2.818 65 (-3) - - - - 5.811 05 (-3) 5.398 14 (-3) - 1.135 48 (-3) - 1.878 39 (- 3) - 1.007 95 (-2) 9.341 12 (-3) -2.180 70 (-3) - 4.080 39 (- 3) - 1.263 72 (- 2) 1.164 22 (-2) -2.818 65 (-3) -5.544 04 (-3) The individual isotopic mobility coefficients, (ik*), are represented with the power of ten bracketed, thus - 6.399 45 (-3) equals -6.399 45 x loe3.The isotope-isotope term, (l/Cz)(Laa*/pop*),.of eqn (46) is obtained by summation of the coefficients of each matrix divided by the correspondmg value of C,T. increase in concentration of total salt, the positive-to-negative ion coupling interactions become significant. This causes (-Laa*/pop*)/Cz to pass through a maximum and subsequently decrease, see fig. 1, tables 1 and 2. TABLE 2.-PREDICTED ISOTOPIC DIFFUSION COEFFICIENTS FOR CADMIUM IN AQUEOUS CADMIUM IODIDE mol dm'3 0.003 0.005 0.01 0.02 0.03 0.04 0.05 0.06 0.07 0.09 0.10 -3.585 7 - 3.772 6 -3.779 1 -3.202 7 -2.451 0 - 1.727 0 - 1.073 8 -0.489 8 0.036 19 0.949 86 1.354 55 2.738 2 2.757 9 2.790 3 2.858 7 2.932 8 2.999 3 3.055 1 3.101 5 3.140 4 3.200 6 3.224 3 7.677 7.772 7.854 7.882 7.878 7.863 7.840 7.810 7.776 7.698 7.657 Values of (1 /Ca Laa*/pop* were obtained from eqn (45) and (46), using concentrations of individual ions in solution and their corresponding equivalent conductances at infinite dilution, A:, as obtained in Paper 2.' (The optimised values of A; for the complex ions were A;, 33.51 ; A;, 39.70 and A:, 68.67 cm2 s1-' equiv-I). The coefficients LaaIC,T given in the third column of this table are those predicted previously where LaalC,' = 2LaallV in the notation of ref.(2). The dimensions of Laa and Laa* are moI2 J-' cm-I S-I.I12 TRANSPORT IN AQUEOUS SOLUTIONS In terms of these Limiting Law predictions the additional positive contributions [(33*) + 2(34*) + (44*)] never overcome this downward trend. Above 0.2 mol dm-3 the experimental data increase once more and pass through a second maximum, fig.1. Such concentrations are far beyond the range of classical theory and it is not surprising that this analysis cannot reproduce them. I 9 - 2.0 ! c. 2 0.4 0.6 dC FIG. 1.-Isotope-isotope coupling terms in eqn (1) and (35) as a function of the square root of rnolarity of cadmium iodide, C,T. Experimental values obtained in Paper 3 and calculated points, from table 2, are represented as and 0 respectively. In Paper 2 the binary coefficients for cadmium iodide were predicted using a similar analysis and calculated values of L,,/C,T obtained. (These were represented as Laa/Ca in that paper.2) This intrinsic mobility coefficient passes through a minimum in dilute solutions, while D,, shows an initial rise to a maximum in the same range.Since both terms in eqn (1) have now been predicted, using the same optimised values for A:, it was of interest to predict the cadmium ion diffusion coefficient, D,,, purely by these theories. The results are given in fig. 2 and table 2. Calculated diffusion coefficients are shown to increase from infinite dilution and pass through a maximum as inferred from experimental data.l Although the method is limited to dilute solutions, the predicted diffusion coefficient is only 2.5 % lower than observed at 0.1 mol dm-3. This method of calculation does not provide an insight into the basic causes of the initial maximum in D,,, shown in fig. 2. From eqn (45) of this text and eqn (20) of Paper 2, it is easily shown that Zik = Iik/p0p* i # k. (47) From eqn (47), (29) and (35) D,, may be expressed as a summationR.PATERSON AND LUTFULLAH 113 where ai equals C,/C,', the proportion of total cadmium present as complex i. dil have the form of diffusion coefficients and are defined by eqn (49) The functions dir are the isotopic diffusion coefficients of the complexes (i). These would be obtained by hypothetical experiments in which complex i alone was labelled and was no longer in dynamic equilibrium with the other cadmium species in solution. 8.2 i 'm N 7.8 ---. \o 2 X ce < 7: 4 7.0 I 1 I I I I 0,2 0; 4 0:6 FIG. 2.--Isotopic diffusion coefficients Daa x lo6 (cm2 s-l) for lr5Cd2+ in aqueous cadmium iodide as a function of the square root of molarity of cadmium iodide, C,T. Measured coefficients obtained previously and those calculated here (table 2) are represented as 0 and 0 respectively.D,, is, therefore, the weighted sum of these diffusion coefficients. At infinite dilution & = d i = RTIIP/IZ,IH2 and from eqn (45) and eqn (20)2 it is easily shown that I I1 (relaxation term) Eqn (50) contains no electrophoretic terms, showing that for a complexed electrolyte, as for dissociated ones, isotopic diffusion is affected only by relaxation contributions. Using limiting equivalent conductances estimated previously,2 d& dTl, d3\ and d& are 7.12 x 8.92 x 10.57 x and 9.14 x respectively in units of cm2 s-I. Each of the complexed ions has a larger diffusion coefficient at infinite dilution than free cadmium ion itself. As concentration is increased and complexation becomes significant (table 2)2 the first summation in eqn (50) increases progressively.Relaxation terms, although zero at infinite dilution, make an increasingly negative contribution, causing D,, to describe a maximum. The numerical output of this calculation is of course identical to that presented in fig. 2 and table 2 and so tabulation of the terms in eqn (50) is not presented.114 TRANSPORT I N AQUEOUS SOLUTIONS What may be concluded from these limiting law predictions is that although all possible coupling interactions between ions have been considered, coupliiig coefficients between chemically different species cancel in eqn (49, reducing the thirty basic mobility coefficients to only eight. These in turn may be grouped to define the four isotopic diffusion coefficients of cadmium and its complexes, each making a propor- tional contribution to D,,. This calculation, taken with those earlier predictions of binary electrolyte transport in cadmium iodide solutions,2 shows that the classical concepts of transport may be applied with considerable success to the predictions of both isotopic diffusion coefficients, transport numbers, conductance and salt diffusion in this complexed electrolyte. We are grateful to the Ministry of Education of the Government of Pakistan for a Research Scholarship to Lutfullah during the period of this work. R. Paterson and Lutfullah, Paper 3, J.C.S. Furaday I, 1978, 74, 93. R. Paterson, J. Anderson, S . S. Anderson and Lutfullah, J.C.S. Faruduy I, 1977, 73, 1773, Paper 2. R. Paterson, J. Anderson and S. S. Anderson, J.C.S. Faraday I, 1977, 73, 1763, Paper 1. J. Anderson and R. Paterson, J.C.S. Furaduy I, 1975, 71, 1335. S. Liukkonen, Acra Polytechnica Scund. (Chemistry including Metallurgy Series), No. 11 3, Helsinki, 1973. M. J. Pikal, J. Phys. Chem., 1971, 75, 3124. 0. Kedem and A. Essig, J. Gen. Physiol., 1965, 48, 1047. R. Paterson, Furaday Disc. Chem. SOC., 1978, 64, in press. (PAPER 7/770)

ISSN:0300-9599    

DOI:10.1039/F19787400103

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Dielectric Properties of N-Methyl Acetamide in Carbon Tetrachloride Solution BY MUSA M. OMAR* Department of Chemistry, Faculty of Science, University of Tripoli, Tripoli, Libya Received 10th June, 1976 The dielectric relaxation parameters have been measured for N-methyl acetamide in carbon tetrachloride solution (0.0016 to 1.5 mol dm-9, at 30.1"C and a frequency range of 450 to 1000 MHz. The static permittivities have been also measured at 100 kHz for twelve different concentrations. The apparent dipole moment and the relaxation times of the amide as functions of concentration have been evaluated. N-methyl acetamide is probably the simplest molecule which can be taken as an adequate model of the units in polypeptide and protein structures. This, in addition to its large permittivity values has caused extensive study of its molecular properties. The studies of this compound have been principally concerned with the dielectric properties of the pure 1iquid.l Little attention has been given to the dielectric study in solution for smaller amide molecules.The dielectric properties of pure liquids or concentrated solutions are attributed to the properties of molecular aggregates. In this work we tried to trace the molecular dielectric parameters of a solution at various concentrations hoping to obtain a complete spectrum of properties, starting from a highly associated medium to a very low or negligible associated system. This was done by utilising the available infrared spectroscopic studies of N-methyl acetamide in carbon tetrachloride solution. - EXPERIMENTAL N-methyl acetamide was purified by fractional distillation under reduced pressure.The solvent carbon tetrachloride was of spectroscopic grade. The amide and the solvent were checked for purity, using infrared spectroscopy and refractive index measurement. Static permittivity measurements were made using a W.T.W. Diplometer (Model 06) with MFL 1 cell for concentrations up to 0.4 mol dm-3, and MFL 2 for high concentrations. The cells were thermostatted through the outer water jacket using a water thermostat providing a temperature control to +O.l"C. The static readings were taken at 100 KHz, which is considered to be below the dielectric absorption frequency range. The other dielectric properties were studied in the frequency range 450 to 1000 MHz, using a General Radio slotted line which was designed by the General Radio Company, Massachussetts, U.S.A.This apparatus has been described fully by William~.~ The material studied was contained in a coaxial line. This was achieved by dipping the line into a vessel containing the liquid. The liquid enters through small holes at the end of the line. The liquid depth was measured using a cathatometer. The vessel including the line was then immersed in a thermostat at a fixed temperature of 30.1"C. RESULTS The dielectric absorption of five different concentrations of the N-methyl acetamide in carbon tetrachloride solution (1.5 to 0.0016 mol dm-3) was studied at a constant 115116 DIELECTRIC STUDY OF METHYL ACETAMIDE TABLE DI DIELECTRIC PARAMETERS OF N-METHYL ACETAMIDE IN CCI, AT 30.1"C onc./mol dm-3 1.5 0.8 0.1 0.02 0.0016 fIMHz 450 650 850 1000 450 650 850 1000 450 650 850 1000 450 650 850 1000 450 650 850 1000 d 3.1 8+ 0.01 2.98+ 0.01 2.91 kO.01 2.79& 0.08 2.87k0.04 2.67+0.00 2.642 0.02 2.55k0.07 2.48+ 0.05 2.45k0.01 2.47+ 0.02 2.44+ 0.07 2.29k0.01 2.28 & 0.03 2.26k0.01 2.23 + 0.01 2.24k0.04 2.23k0.03 2.20+ 0.01 2.1 8 _+ 0.02 E *' 1.700k0.020 1.11 5+ 0.031 0.857 & 0.025 0.774k 0.050 1.018+0.007 0.570+0.001 0.5 I2+ 0.007 0.366k0.045 12.1 & 0.5 x 11.8+0.4x 9.8k0.7 x 8.4k0.5 x 6.0k0.3 x 6 .9 k 0 . 6 ~ 7.2+ 0.9 x 6.92 0.4 x 1.5k0.2~ 2.8+ 0.3 x 6.1k0.5~ 5.6k0.5 x temperature of 30.1"C and at four different frequencies (450,650,850 and 1000 MHz). The measured permittivity of the carbon tetrachloride of 2.22 + 0.02 is in good agree- ment with the literature value of 2.218 at the same temperature. The losses observed for the solvent are too small to have much significance in the losses for the concentrations above 0.02 mol dm-3.The relaxation parameters as deduced are given in table 1. The reported dielectric parameters are the mean values of two or three determina- tions using different sample depths. It is found that the agreement between different TABLE 2.-cOMPARISON OF THE ESTIMATED DIPOLE MOMENTS USING DIFFERENT VALUES FOR Em concl mol dm-3 1.5 1 .o 0.8 0.6 0.4 0.2 0.1 0.08 0.04 0.02 0.01 0.001 6 EO 22.15 12.66 11.86 9.56 5.97 3.78 2.78 2.65 2.38 2.29 2.25 2.23 PD 2 ea = ND 10.71 9.51 10.21 10.31 9.09 8.44 7.52 7.54 7.47 8.54 10.96 8.02 PD Em = E0* 10.67 9.46 10.61 10.24 8.97 8.18 6.93 6.79 5.86 5.43 5.86 7.32 PD C + O Em = EO 10.67 9.46 10.15 10.24 8.96 8.15 6.89 6.74 5.75 5.24 5.40 4.63 * PD 10.68 10.01 - - - 1 5.63 - - 5.07 - -M. M.OMAR 117 depths varied from 1 to 3 % in permittivity (8') and from about 2 to 10 % in the loss ( E " ) values. The static permittivity of twelve different concentrations (including the one mentioned above) was also measured at the same temperature (30.1"C) and the results are reported in table 2. ANALYSIS OF THE DATA The experimental results for the four higher concentrations (1.5, 0.8, 0.1 and 0.02 mol dm-3) were analysed in terms of a distribution of relaxation times, using the Fouss-Kirkwood equation (see fig. 1) cosh-I ~ ~ , ~ / s " [ 2 +(l/skax)2/2 +(I/s')~] - - pin f!i where B is a measure of deviation from a single relaxation time, 0 < p < 1, and E k a x is the permittivity obtained at maximum absorption.This equation was used 1.4 \ FIG. 1 .-Fuoss-Kirkwood plots for (c) 1.5 (O), (b) 0.1 (9) and (a) 0.02 mol dm-3 (9) NMA solution at,30.1 "C. x = 2+('/&rnax)*/2+('/~~ax)* E" 22 - ---__ -- El FIG. 2.-Cole-Cole plots for (a) 1.5 (0) and (&) 0.8 mol dm-3 (0) NMA solution at 30.1OC.118 DIELECTRIC STUDY OF METHYL ACBTAMIDE for the 1.5 and 0.8 mol dm-3 solutions, which have large losses. For 0.1 and 0.02 mol dm-3 solutions where E” 4 E’ and E’ = &&ax, the following reduced equation was used The results were checked against a Cole-Cole plot (see fig. 2) to ensure that the permittivities were consistent with the relaxation parameters obtained.In all cases a distribution of relaxation times was observed, and the Fuoss-Kirkwood distribution parameter (p) was converted to the Cole-Cole distribution parameter (a) (p = I for single relaxation times) and shown in table 3. TABLE 3 .-COLE-COLE AND FUOSS-KIRKWOOD DISTRIBUTION PARAMETERS B B &Ax Emax conc/mol dm-3 Cole-Cole Fuoss-Kirkwood Cole-Cole Fuoss-Kirkwood 1.5 0.87 0.90 7.10 7.10 0.8 0.90 0.92 4.00 4.00 0.1 0.89 0.98 0,155 0.155 0.02 - 0.97 - 7.21 x 10-3 The results for the 0.0016 mol ~ l m - ~ solution indicate two overlapping absorptions, but it was not possible to analyse them according to the Fuoss-Kirkwood distribution of relaxation times. These results were analysed, assuming two overlapping simple Debye absorptions and using a suitable IBM 1620 computer program (see fig.3). log f FIG. 3.-Analysis of dielectric absorption of the 0.0016 mol dm-3 NMA solution, 0 experimental points. The relaxation parameters are listed in table 4. The distributed relaxation times The dipole moments associated with these relaxation processes were calculated are estimated to be correct to between 8 and 12 %. from the absorption intensities using the Onsager equation :M. M. OMAR 119 where S is the number of the polar molecules per cm3 of solution, p2 is the dipole moment of the solute, gal and em2 are the static and the optical permittivities of the solvent and solute respectively, and c0 and em are those for the solutions. The calculated dipole moments are listed in table 2.TABLE 4.-RELAXATION PARAMETERS FOR N-METHYL ACETAMIDE SOLUTION IN CARBON TETRA- CHLORIDE AT 30.1"C , conc/mol dm-3 t x 1012/s Emsx B EO Em PD 1.5 3 690 7.10 0.90 22.15 2.65 10.68 0.8 3650 4.00 0.92 11.85 2.50 10.01 0.1 546 0.155 0.98 2.78 2.41 5.63 0.02 187 7 . 2 1 ~ 0.97 2.29 2.23 5.07 0.001 6 62.3* - - 2.23 * 71 I - For the 0.0016 mol dm-3 solution, the dipole moment associated with each process is calculated from the absorption intensities using the Cole treatment : P: = A&;I(&o -&m)P2 (80 -Em) where ,ul and p2 are the dipole moments associated with the low and high frequency absorptions respectively. 2 A 4 I 2 P2 = - P DISCUSSION As has been shown earlier, the results for the 1.5, 0.8, 0.1 and 0.02moldm-3 solutions of N-methyl acetamide (see table 4) are adequately described in terms of a distribution of relaxation times.In all these cases the distribution parameter (fl) increases, as expected, with decreasing concentration and tends to unity (corresponding to a Debye single relaxation time). Bass, Meighan and Cole have measured the dielectric parameters for pure liquid N-methyl acetamide in the frequency range 1-250 MHz. Their results have been described by a simple Debye function with a single relaxation time. This does not follow the trend which we expected from our solution results and for such strong hydrogen bonded molecules. The values obtained hare for the relaxation time of N-methyl acetamide solution are much larger than expected for such small molecules. The estimated value for the relaxation time of the 1.5 mol dm-3 solution is 3690x 10-l2 s.This is much higher than the value (7,) obtained for the pure liquid (N-methyl acetamide) of 74 x 10-l2 s at 31.4"C. At the same time the relaxation time obtained for 1.45 rnol dm-3 solution of N-methyl palmitamide in par& wax at 91°C is 6 8 0 2 ~ s. Taking into consideration the molecular size of the latter, our value for N-methyl acetamide is in fair agreement with that for N-methyl palmitamide solution. The large values for N-methyl acetamide could be attributed to the molecular aggregation formed by chain association. Decreasing concentration decreases the molecular association of such agglomerates which leads to a decrease in the relaxation times (see table 4 and fig. 4). This does not seem to be true at higher concentrations (0.8 and 1.5 mol dm-3) where the relaxation times kept almost constant.Taking into consideration the small values obtained for the pure liquid, one could suggest that at much higher concentration the apparent relaxation times may start to decline120 DIELECTRIC STUDY OF METHYL ACETAMIDE with increasing concentration. This behaviour may be explained by the formation of some agglomerates (cyclic association) in pure liquids and highly concentrated solutions which are different from those formed at low concentrations (chain association) and rotate faster than the latter. The dielectric behaviour of the 0.0016 mol dm-3 solution which is described by two overlapping Debye processes should not be taken as giving unique values, because the values obtained carried a large uncertainty due to the small values of the dielectric losses.However, the values obtained for low and high frequency relaxation times are z1 = 62.3 x 10-l2 s and z2 = 21.5 x 10-l2 s respectively. For pure liquid n-ethyl alcohol, which is associated in a linear form similar to mono-substituted amides, Saxton l1 resolved the overall absorption into two separate Debye relaxation times, z1 = 170 x 10-l2 s and z2 = 1.6 x 10-l2 s at 20°C. Garg and Smyth l2 have observed for n-propyl alcohol three Debye relaxation processes with single relaxat ion times. I 1 . - - -- as 0.5 1.0 1.5 concentration/mol dm-3 FIG. 4.-Variation of NMA static permittivity (0) and relaxation time (9) with its molar concmtra- tion at 30.1 "C. Similar results have been obtained lo for a dilute solution of N-methyl palmitamide (0.097 and 0.049 mol dm-3) in paraffin wax, where the dielectric absorption has been resolved into two overlapping Debye relaxation processes.The values obtained for the 0.049 mol dm-3 solution are z1 = 208 x 10-l2 s and z2 = 30 x 10-l2 s, at 91°C. Our relaxation time values (2,) for 0.0016 mol dm-3 solution of N-methyl acetamide seem to be in good agreement with the N-methyl palmitamide solution. Since the two relaxation times were not followed up at lower concentrations or at different temperatures, because of the experimental difficulties due to the small values of the observed losses, it is not possible to establish any general pattern or conclusion, but some idea could still be formed from these results.Davies and Evans have studied the infrared spectroscopy of N-methyl acetamide solution in carbon tetrachloride. They have shown that the (N-H) non-bonding stretching vibration absorption dominates at the 0.001 6 mol dm-3 concentration. Therefore the assignment of the two relaxation times could be suggested as follows : the low frequency relaxation time (zl) could be attributed to the rotation of the free moleculesM. M. OMAR 121 (monomeric), and the high frequency relaxation time ( T ~ ) could be associated with the rotation of the whole polar head group around the R-C1 bond.1° 0 // \ / \ R-C1 CH3 N H The rotation round the C-N bond is much faster l 3 and cannot be expected to be observed at this frequency range. The other possibility is that the z1 is attributed to the rotation of the associated molecules ; this is supported by the large estimated value of the dipole moments; while the z2 is associated with the non-associated molecules.I 12 I a2 0.6 1.0 L.4 1.6 concentrationlmol dm-3 FIG. 5.Variation of NMA dipole moment with its molar concentration at 30.1"C. The variation of static permittivity with concentration is shown in fig. 4; it is found, as expected, that eo is a linear function of the molar concentration except at the lower concentration (below 0.1 mol dm-3) where eo values hardly varied. The choice of a value for 8, presents a problem in the calculation of the apparent dipole moment, since if we select E , = n& for each solution we find that the calcula- tions are quite meaningless.Using E , = E' (solvent) improves the apparent dipole moment. However it seems that using the extrapolated value of E~ (solution) at zero concentration as gives better results. In all the above cases the ga2 was taken to be l4 1.05 &. In table 2, four values of the apparent dipole moment (p) are listed. It is observed [even using E , (observed), obtained at high frequency or from Cole-Cole plots] that the dipole moments are much larger than those obtained for N-methyl acetamide l5 gases (3.71 D). This confirms that mono-substituted amides do associate in linear chain and that the expected dipole moments of the associated molecules must be larger than the monomer. Similar results have been obtained by Cole l * lS and122 DIELECTRIC STUDY OF METHYL ACETAMIDE coworkers for pure liquids and a benzene solution of N-methyl acetamide.The large, strongly concentration dependent (see fig. 5 ) dipole moment shows large association effects, a phenomenon which is, as also explained by Cole,I a result of chainwise association by hydrogen bonding with the single amino-hydrogen such that adjacent molecular dipole moments are nearly parallel. This phenomenon seems to exist even in highly diluted solution (0.0016moldm-3) where a value of - 1.3 times that of the dipole moment of a monomer was obtained. However, the reported low concentration values of the dipole moment are very sensitive to (E,, - E,) values, i.e. any small error in go or 8, leads to a large difference in the calculated dipole moment. Thus the figures reported for 0.01 and 0.0016 mol dm-3 solutions are exposed to a large uncertainty.The dipole moments p l , associated with the low frequency relaxation mechanism and p2, which is associated with high frequency one, were calculated from the absorption intensities and values of 2.00 and 4.17 D were obtained respectively. I thank the Faculties of Science, Engineering and Education at the University of Tripoli for use of facilities and equipment during this research. I am also indebted to Dr.-Eng. M. Wanas and Eng. E. El-Ghazzawi for their valuable assistance in installing the equipment and apparatus. S. J. Bass, W. I. Nathan, R. M. Meighan and R. H. Cole, J. Phys. Chem., 1964, 68, 509. M. Davies, J. C. Evans and (in part) R. L. Jones, Truns. Furahy SOC., 1955,51,761. (a) G. Williams, Ph.D. Thesis (University College, Wales, 1966) ; (6) A. Von Hippel, Dielectric Materials and Applications (NIT Press, Boston, 1961). Table of Dielectric Constants of Pure Liquids, ed. A. A. Maryott and E. R. Smith (National Bureau of Standards Circular 514, 1951). R. M. Fuoss and J. G. Kirkwood, J. Amer. Chem. SOC., 1941,63,385. K. S. Cole and R. H. Cole, J. Chem. Phys., 1941, 9, 341. S. R. Gough, Ph.D. Thesis (University College, Wales, 1964). S. L. Onsager, J. Amer. Chem. SOC., 1936,58, 1486. R. H. Cole, J. Chem. Phys., 1965,42,637. lo M. M. Omar, Ph.D. Thesis (University College, Wales, 1967). l1 Dielectric Properties and Molecular Behaviour, ed. N. Hill, W. E. Vaughan, A. H. Price and l2 S. K. Garg and C. P. Smyth, J. Phys. Chem., 1965, 69, 1294. l3 T. Drakenberg, K. Dahlqvist and S. Forskn, J. Phys. Chem., 1972, 76, 2178. l4 C. Campbell, G. Brink and L. Glasser, J. Phys. Chem., 1975, 79, 660. l5 R. M. Meighan and R. H. Cole, J. Phys. Chem., 1964, 68,503. M. Davies (Van Nostrand, 1969), p. 352. (PAPER 6/1099)

ISSN:0300-9599    

DOI:10.1039/F19787400115

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Photo-oxidation of Floating Hydrocarbon Oils in the Presence of some Naphthalene Derivatives B Y WILLIAM H. K. SANNIEZ AND NEITON PILPEL" Department of Pharmacy, Chelsea College, University of London, Manresa Road, London, SW3 6LX Received 6th October, 1976 Samples of pure isopropylbenzene and methylbenzene and of a commercial mineral oil, Dobane, containing various concentrations of 1-hydroxynaphthalene, naphthalene-a-aldehyde and l-nitro- naphthalene were floated on water as layers 0.6 cm thick and irradiated with ultraviolet light over periods of 240 min. By combining measurements of the changes occurring in interfacial tension with chemical and physical analysis of the oil and water phases, it has been possible to identify some of the products of photo-oxidation and to propose mechanisms for the reactions.The oxidation of hydrocarbons is a complex process which involves a large number of concurrent and consecutive, independent and connected stages and this frequently makes it difficult to analyse the re~ults.l-~ Oxidation by photochemical methods can sometimes be useful for providing mild conditions which can be controlled to prevent the onset of complex secondary reacti0ns.l Recent work has shown that oxidation of certain hydrocarbons can also be effected by irradiating them with U.V. light in the presence of suitable photo sensitizer^.^-^ Thus anthracene and certain polynuclear hydrocarbons act as photosensitizers for the homolytic decomposition of acetyl peroxide, yielding the same products as in the thermal reaction. 1 -hydroxynaphthalene acts as a photosensitizer for the oxidation of certain alkanes and alkylbenzene~.~* Chemical reactions in organic compounds can sometimes be investigated by floating them on the surface of an aqueous substrate and observing the changes that occur in their phase potential and surface pressure?* lo* l1 In the present work, hydrocarbon oils containing selected additives were floated on triply distilled water and irradiated with U.V.light for various periods of time. By combining measurements of the changes in their interfacial tensions with physicochemical analysis of the oil and water phases, it was hoped to establish the mechanisms and kinetics of the photochemical reactions. EXPERIMENTAL MATERIALS Puriss isopropylbenzene and AnalaR methylbenzene were further purified by passing them three times under slight pressure through a tightly packed bed of Fuller's earth 2.5 cm thick, which was renewed after each passage. l2 The hydrocarbon oil, Dobane JN from Shell Chemicals, was purified by heating three times with 5 % w/v Fuller's earth at - 100°C for 30 min and then filtered [m.w.232 ; q20 1 I .4 CP ; d$O 0.878 ; interfacial tension (IFT) 31.4 mN m-'1. 1 -hydroxynaphthalene and 1-nitronaphthalene were good grades from BDH and were 123I24 FLOATING HYDROCARBON OILS further purified by crystallisation from water or ethanol + water. Puriss naphthalene-a- aldehyde was obtained from Koch-Light laboratories and was used without further purification. The physicochemical properties of all the materials-density, surface and interfacial tension, refractive index and melting point-agreed closely with reliable literature values.Triply distilled water, surface tension 72.1 mN m-l, conductivity 1.2 x 10-‘ Q-’, was used throughout. FIG. 1.-Irradiation chamber. A, mercury lamp (125 W) ; By cool air blower ; C, silica filter (2 mm thick) ; D, 7.5 cm diameter fan ; E, thermometer; F, 5.0 cm Du Nouy Pot ; G, glass thermostat (water temp. 25 k 2°C) ; H, blocks. TECHNIQUE A N D APPARATUS Dilute solutions of the naphthalene derivatives were prepared in the three hydrocarbon solvents using an electromicrobalance for the weighings. 12cm3 of each sample was floated on 20 cm3 of distilled water in a 5 cm diameter glass pot, producing a layer 0.6 cm thick and irradiated with U.V. light for various periods of time in the apparatus shown in fig.1. The U.V. source was a 125 W Phillips MB/U medium pressure arc equipped with a silica filter which transmitted in the range 290-325 nm.” * During irradiation the temperature of the samples was maintained at 2552°C by circulating a 1 % w/v aqueous solution of chrome alum+cupric sulphate (1 : 1) which was also transparent to light of this wavelength. After irradiation, the interfacial tensions, y, of the samples were measured in sitri with a du Nouy tensiometer using the equation (1 1 where y is the interfacial tension, P the tensiometer reading, C is the circumference of the platinum ring, D is density of water at 20°C and d is the density of hydrocarbon at 2OoC, and the results were corrected to 20°C by the factor -0.04 m Nm-l “el. This technique required the floating hydrocarbon layer to be at least 0.5 cm thick.The peroxides produced in the reactions were determined by iodometric titration 1 3 9 and the other products in the oil and water phases, after 4 h irradiation, were analysed by g.1.c. and thin layer chromatography combined with U.V. spectroscopy. = P(0.725+ r0.0145 P/C2(D-d)]*) RESULTS The effects of irradiation on the interfacial tensions of representative samples are shown in fig. 2(u)-(c). All the systems exhibited decreases in IFT with increasing periods of irradiation; the decreases depended also on the nature and amount of naphthalene derivatives that had been added to the hydrocarbons.W. H. K . SANNIEZ AND N . PILPEL 125 It was found that for all the unirradiated systems there was a linear relationship between the change in interfacial tension and the log concentration c, of the additive, showing that the Gibbs' equation -dy/d log c r = 2.303 RT where r is the surface excess, R is the gas constant and T is the temperature in K ap~1ied.l~ Values for the surface excess are given in table 1.1 0 ,, , u L_ I I IC 1 I I u- 0 I00 2 0 0 300 irradiation time/min FIG. 2.-(u) Plot of interfacial tension against irradiation time for methylbenzene system. A, methylbenzene ; A, methylbenzene+2.2 x mol dm-3, 1-nitronaphthalene ; 0, methylbenzene+ 2.2 x mol dm-3 hydroxynaphthalene ; 0, methylbenzene+2.2 x lod3 rnol dm-3 naphthalene- a-aldehyde. (b) Isopropylbenzene system. mol dm-3 1-nitronaphthalene ; 0, isopropylbenzene+ 2.2 x 1O-j mol dm-3 hydroxynaphthalene ; 0, cumene+2.2 x rnol dm-j naphthalene-a-aldehyde.(c) Dobane system. A, Dobane ; A, Dobane+2.0x rnol dm-3 1-nitronaphthalene ; 0, Dobane+2.0 x mol dm-3 l-hydro- A, isopropylbenzene ; A, isopropylbenzene+ 2.2 x xynaphthalene ; 0, Dobane+ 2.0 x mol dm-3 1-naphthalene-a-aldehyde. For the isopropylbenzene and methylbenzene, though less accurately for the Dobane systems, the slopes -(dy/dt) of the graphs in fig. 2(a) and (b) also remained reasonably constant over the first 2 h of irradiation and when log (-dyldt) was plotted against the log concentration of the additives in these three solvents, straight lines were obtained as shown in fig. 3(a)-(c). TABLE 1 .-SURFACE EXCESS VALUES FOR THE VARIOUS SYSTEMS surface excess I'/pmol m-2 1 -naphthalene solvent I-nitronaphthalene 1-hydroxynaphthalene -a-aldehyde isopropyl- benzene 0.14 0.22 0.24 Dobane 0.001 0.004 0.007 methyl- benzene 0.14 0.13 0.18126 FLOATING HYDROCARBON OILS Fig.4(a)-(c) are representative plots to show how the peroxide concentrations in the oils varied with the period of irradiation, with the nature of the oil and with the amount of additive employed. Initially the Dobane appeared to be more staPle than either isopropylbenzene or methylbenzene, but the occurrence of maxima and points of inflexion in the curves suggests that in all three solvents several different mechanisms could be contributing to the initial formation and, at later stages, to the decomposition of the peroxides formed. -‘r -21----- - I - -_25 -4 -3 -2 0 log concentration FIG.3 . 4 2 ) Plot of -log dy/dt against -log concentration. A, methylbenzene+ 1-nitronaphthalene 0, methylbenzene + 1-hydroxynaphthalene ; 0, methylbenzene+ 1-naphthalene-a-aldehyde. (b) Isopropylbenzene. A, isopropylbenzene + 1-nitronaphthalene ; 0, isopropylbenzene + 1 -hydroxy- naphthalene ; 0, isopropylbenzene + 1-naphthalene-a-aldehyde. (c) Dobane. A, Dobane + 1- nitronaphthalene ; 0, Dobane + 1-hydroxynaphthalene ; 0, Dobane+ 1-naphthalene-a-aldehyde. T.1.c. analysis of the isopropylbenzene and methylbenzene systems after 4 h and of the Dobane system after 3 h of irradiation revealed several fluorescent spots when viewed in light of 350 nm wavelength. Their R, values are shown in table 2. The individual spots were extracted with diethyl ether, redissolved in methanol and their U.V. absorption spectra were recorded.Typical spectra are shown in fig. 5 and 6 . By means of g.1.c. analysis and the use of standard spectra, the main reaction products from the isopropylbenzene and methylbenzene systems were identified and their amounts determined (table 2). The amounts formed in the Dobane systems, however, were mostly too small for them to be analysed.0 I Q E m U 0 151 G 10 5 0 0 I00 200 300 irradiation timelmin FIG. 4.-(a) Graph of peroxide concentration against irradiation time. A, methylbenzene ; A, methylbenzene+2.2 x mol dm-3 1-nitronaphthalene ; 0, methylbenzene+2.2 x lo-' rnol dm-3 1 -hydroxynaphthalene ; 0, methylbenzenef 2.2 x mol dm-3 1-naphthalene-a-aldehyde. (6) A, Isopropylbenzene ; A, isopropylbenzene+2.2 x mol 1-nitronaphthalene ; 0, iso- propylbenzene+2.2 x rnol dm-3 1-hydroxynaphthalene ; 0, isopropylbenzenef2.2 x mol 1-naphthalene-a-aldehyde.(c) A, Dobane ; A, Dobane+ 1.0 x lo-' rnol dmV3 l-nitro- naphthalene ; 0, Dobane+ 1.0 x lo-' mol dm-3 1 hydroxynaphthalene ; 0, Dobane+ 1.0 x 10-1 mol dm-3 1-naphthalene-a-aldehyde. I . "9 8 1.2 - $I I.0- e % 0.8- 0.6 - 0 . 4 1 0 . 2 - 0.0 ' 2 0 0 2 5 0 3 0 0 350 400 wavelength/nm acid ; B, naphthalene. FIG. 5.-U.v. adsorption spectra of oxidation products from the methylbenzene system. A, benzoic128 FLOATING HYDROCARBON OILS -1 1 2 0 0 2 5 0 3 0 0 350 400 waveIength/nm met hox ynapht halene ; D , met hylphenylket one. FIG. 6.-U.V. absorption spectra of oxidation products from the isopropylbenzene system.C DISCUSSION It is apparent from the results that the three hydrocarbon oils on their own are relatively stable to U.V. irradiation in the range 290-325 nm over periods of 240 min. They all absorb U.V. light within this range ’* * but the changes produced in their interfacial tensions and peroxide values are very small. Of the three oils, isopropyl- benzene appears to be marginally the most affected. This is presumably because its benzylic hydrogen is more reactive than that of methylbenzene (due to the proximity of the extra methyl groups) and that of Dobane, which being a commercial oil has been specially blended to be resistant to oxidation.lg In the presence of the additives, however, all three oils are quite rapidly oxidised, as shown by the build up of peroxide/hydroperoxide [fig.4(a)-(c)] and the corres- ponding decreases in their interfacial tensions [fig. 2(a)-(c)]. It is well known 4 9 2o that the oxidation of hydrocarbon oils proceeds via the formation of peroxides and/or hydroperoxides. These subsequently start to de- compose [as demonstrated by the maxima and points of inflexion in fig. 4(a)- (c)] to yield surface active products-alcohols, ketones, etc., some of which have been identified by g.l.c., t.1.c. and spectroscopy in table 2 and fig. 5 and 6 . On more prolonged oxidation, a variety of other products may be formed, including acids, polymers and waxes,lS 2o but the routes followed are then very complex and not readily amenable to the present type of analysis. For isopropyl benzene on its own, the initial slow oxidation could proceed as follows: hv 202 2*PhCH(CH3)2 -+ PhC(CH3)200H+ Phd(CH3)2 + HG2.(3) (hydroperoxide) The corresponding reaction for methylbenzene on its own could be 2o as in reaction (4) hv 202 2PhCH3 -+ PhCH200H + PhdHz + HG2. (4) (hydroperoxide) * Ph = phenylW. H. K. SANNIEZ AND N. PILPEL 129 TABLE PR PRODUCTS OF OXIDATION AND THEIR CONCENTRATION AFTER 240 min IRRADIATION isopropylbenzene a 1 -ni tronapht halene 1 -hydroxynaphthalene naphthalene- a-aldeh yde 1 -nitronaphthalene 1 -hydroxynaphthalene naphthalene- a-aldehyde 1 -nitronaphthalene 1 -hydroxynaphthalene naphthalene- a-aldeh yde spot 1 2 1 2 1 1 2 1 1 1 1 1 2 retention Rf value timelmin 0.75 6.4 0.92 1.2 0.75 6.3 0.91 4.2 0.76 6.4 methylbenzene a 0.72 0.8 0.92 1.2 0.71 0.8 0.72 0.9 Dobane 0.83 1.3 0.84 0.8 0.39 0.85 conc.x identity of spot lO4lmol d m - 3 met hylphenyl 2.0 naphthalene 1.5 methylphen yl 2.4 met hoxy- 1.5 met hylphenyl 3.1 ketone ketone naphthalene ketone benzoic acid 1.7 naphthalene 1.6 benzoic acid 2.1 benzoic acid 2.5 naphthalene 0.6 benzoic acid 1.1 * Products from 12 cm3 of hydrocarbon solvent containing 2.5 x lo-' mol dm--3 additive ; b products from 12 cm3 of hydrocarbon solvent containing 1.5 x 10-1 mol dm-3 additive. The additives 1 -nitronaphthalene, 1 -hydroxynaphthalene and 1 -naphthalene-a- aldehyde act as sensitizers by causing the isopropylbenzene and the methyl benzene to produce more radicals, thus accelerating their oxidation. It is known that aromatic nitro compounds, such as 1-nitronaphthalene, are thermally oxidised to their parent hydrocarbons 21 and if they are similarly decom- posed by U.V.light, this would account for the detection of naphthalene (see table 2) in the systems to which 1-nitronaphthalene had been added. Thus the reaction involving the 1 -nitronaphthalene could be h v 0 2 PhCH3 + ArN02 PhCOOH + *ArH + HN02 (5) (benzoic (naphthalene) acid) and the reaction involving the 1-hydroxynaphthalene could be reaction (6) 0 hr 11 PhCH(CH3)2 + ArOH 3 PhCCH3 + ArOCH, + H20 0 2 (methyl- (met hoxy- phenyl naphthalene) ketone) * Ar = naphthyl130 FLOATING HYDROCARBON OILS to form the methylphenyl ketone and the methoxynaphthalene which were detected in the isopropylbenzene system (table 2). With naphthalene-a-aldehyde, however, the only detectable products were methylphenylketone (from the isopropylbenzene) and benzoic acid (from the methyl- benzene), eqn (6) and (5).Thus, although naphthalene-a-aldehyde accelerates the photo-oxidation of isopropylbenzene and of methyl benzene it does not itself appear to be decomposed.6 This was confirmed by analysis of the test samples over the period of the irradiation. As has been explained in greater detail elsewhere? one can employ the results in fig. 3(u)-(c) to determine the overall orders of reaction (n) for the three oil systems with respect to the initial concentrations of the additives in each. The equation of these graphs can be written as eqn (7) log dy/dt = log K+n log c (additive) (7) where K is a constant for each graph and n is the slope.It was found that the slopes of all the graphs fell between zero and 0.5 showing that virtually zero order kinetics was being followed during this first stage of photo-oxidation, irrespective of the nature of the oil concerned. The similarity in the results obtained for Dobane and for the pure hydrocarbon oils suggests that the first stage in the oxidation of Dobane proceeds by a similar mechanism to that for isopropylbenzene and methylbenzene though for this oil it was not always found possible to analyse the products of the reaction. The techniques employed in this work should prove useful for assessing the oxidation stability of other oils and the sensitizing properties of other additives. N. M. E. Emmanuel, The Oxidation of Hydrocarbons in the Liquid Phase (Pergamon, London, 1965). R. G. W. Norrish and M. H. Searby, Proc. Roy. Soc. A, 1956,237,464. N. Pilpel and B. F. J. Hunter, J. Colloid Interface Sci., 1970, 33, 615. W. A. Noyers Jr., G. S. Hammond and J. N. Pitts Jr., Advances in Photochemistry (Interscience, New York, 1963), vol. 1, p. 209. C. Luner and M. Szwarc, J. Chem. Phys., 1955,23,1978. C. Walling and M. J. Gibian, J. Amer. Chem. Soc., 1965,87, 3413. ' A. E. Klein and N. Pilpel, J.C.S. Faraduy I, 1973,69, 1729. A. E. Klein and N. Pilpel, J.C.S. Faraday I, 1974,70, 1250. M. Ottolenghi, J. Amer. Chem. Soc., 1963, 85, 3557. lo D. R. Augood and G. H. Williams, Chem. Rev., 1957,57, 123. l1 N. Pilpel, J. Colloid Sci., 1956, 11, 51. l2 D. R. P. Murray, Bull. Entomol. Res., 1939, 30, 211. l3 D. K. Banjee and C. C. Budke, Analyt. Chem., 1964,36,2367. l4 C. H. Lea, J. Soc. Chem. I d , , 1939,65,286. l5 E. Hutchison, J. Colloid Sci., 1948,3,219. l6 Friedel and Orchin, Ultraviolet Spectra of Aromatic Compounds (John Wiley, N.Y., 1951). la J. C. W. Chien, J. Phys. Chem., 1965,69,4317. l9 Shell Chemicals, London, data sheet on Dobane. *O V. F. Fedorova, in The Oxidation of Hydrocarbons in the Liquid Phase (Pergamon, London, J. G. Calvert and J. N. Pitts Jr., Photochemistry (John Wiley, N.Y., 1966), chap. 5. 1965). E. K. Fields and S. Meyerson, J. Amer. Chem. Soc., 1967,89,724. (PAPER 6/1872)

ISSN:0300-9599    

DOI:10.1039/F19787400123

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Site Group Interaction Effects in Zeolite-Y Part 1.-Structural Examination of the First Stages of the Ag Ion Exchange BY MARTIN COSTENOBLE AND ANDRE MAES* Katholieke Universiteit Leuven, Centrum voor Oppervlaktescheikunde en Colloidale Scheikunde, De Croylaan 42, B-3030 Heverlee, Belgium Received 13th October, 1976 The localization of Na+ and Ag+ ions in hydrated zeolites NaAg-Y containing 2,7.25 and 14 Ag+ ions/u.c. is studied by X-ray diffraction. A high preference of Ag for site I is observed. On increasing the silver content the Ag occupancy of site I reaches a maximum value of %4 Ag/u.c. at the expense of Na+ ions in site I’. The Ag to Na preference pattern derived is; site I > site I1 > site I’ W unlocalized. The majority of ion exchange data on zeolites-X and -Y have been explained in terms of a static model of neutralization in small and large cavities, combined with hydration energy effects and the crystallographic radius of the exchanging ions.The number of possible neutralization points being in excess of the number of exchangeable cations, repositioning can occur during ion exchange.2 9 For example, the distribution in hydrated Na-Y is different from that of the K analog~e,~ pointing to the flexibility in network neutralization. However, two main constraints seem to limit the cation positioning. First, Mortier et aL4 state that each framework oxygen should within small limits bear t of the charge. Secondly there are exclusion rules which must be obeyed. Both the characteristics of the ions and the ion exchanger determine the selectivity of an exchanger for a certain cation.Recently, Maes and Cremers ti argued that cation positioning in small cages influences the selectivity in the large cages and vice versa. In Part 2 7 more evidence is presented showing the influence of the different nature of the ions in the large cavities on the Na+-Ag+ selectivity in the small cages. Studies of cation location in hydrated zeolites are still relatively scarce and are mostly restricted to the homoionic forms. It is the purpose of the present study to investigate cation positioning in hydrated zeolite-Y containing two different cations. In particular, three compositions, corresponding to the initial stages of the Ag+-Na+ equilibrium are investigated by powder X-ray diffraction techniques.The distribu- tion of the exchangeable cations over the different crystallographic sites is discussed in terms of the crystallographic radius, hydration energies and electronegativities of the two competing ions. EXPERIMENTAL SAMPLE PREPARATION A Linde Y-zeolite saturated with Na+ ions by repeated exchange with 1 mol dm-3 NaCI solution, subsequently washed until Cl- free (liquid/solid ratio on a weight basis ~ 3 5 ) ~ and air dried is called Na-Y and contains 55 Na+ ions/u.c. From the ion exchange data in the Part 2,’ zeolite-Y is shown to be initially very selective for Ag+. Three samples were 131132 SITE GROUP INTERACTION IN ZEOLITES prepared by adding a specific amount of a AgN03 solution to air-dried zeolite (liquid/ solid ratio -4). Overnight equilibration was followed by centrifugation and decantation of the supernatant liquid.The samples were air dried without previous washing, as the Ag+ content in the equilibrium solution was respectively 0.05,l and 5 % of the total solution ion content, resulting in negligible Ag in the interstitial solution of the sediment. The air-dried samples were analysed for Ag by atomic absorption methods after acid breakdown and contain respectively 2, 7.25 and 14 Ag+ ions/u.c. The sample table specifies the synthetic faujasite (Y), the ions (Na+/Ag+) and the amount of Ag+ ions/u.c. (e.g. Na/7.25 Ag-Y). The site notation of Smith is used. X-RAY DIFFRACTION AND CALCULATIONS The recording of the X-ray diffraction spectra, the data sampling and the calculation procedures are given elsewhere.lO The scattering due to unlocalised cations and water molecules was accounted for by applying a liquid scattering function.The electron densities of the ions and water molecules were summed and compared with those expected for water molecules. The unlocalised contribution was smeared out in spheres of radii 0.25 and 0.57 nm, corresponding to the sodalite units and large cavities respectively. An additional temperature factor of 15 x nm2 was applied. The residual of the intensities was defined as RI = z ~ ( k Io- IJ/& k Io. A least-squares refinement was carried out with the full-matrix program POWOW,ll based on intensities. The following weighting scheme was introduced The program ORFFE l2 was used to calculate the interatomic distances and angles.RESULTS AND DISCUSSION Table 1 lists the extra framework located material and other relevant information, such as unit cell dimensions, RI values and interatomic distances within the structures Na-Y,I0 Na/2Ag-Y, Na/7.25Ag-Y and Na/14Ag-Y. The structure parameters, framework interatomic distances and bond angles and the values of I. and Ic are not essential to the discussion and are not included in the results. These data are available upon request. Since three kinds of electron density (Ag+, Na+, HzO) can overlap on the same site the following guidelines were used to interpret the observed electron densities at a specific distance from the framework oxygen. The theoretical ionic distances are Na+-02- = 0.235 nm, Ag+-02- = 0.2266 nm and H202--02- = 0.280 nm.According to the short distances observed (see table 1) on site I (site 1-03) and site I1 (site 11-02) these electron densities are attributed solely to Ag+ and Na+ ions. The long site 11'--02 distance suggests that only water occupies site 11', as is generally accepted.13-l The framework of the relaxed hydrated faujasite structures is neutralized according to a specific pattern.4* lo Nevertheless the number of charges coordinated around the O3 oxygens (site I+site 1') is approximately constant: 18.5 in Na-Fj,lG 16 in Na,Ca-Fj l4 and 17.3 in Na-Y.I0 Since even in the case of the Na/l4Ag-Y sample the overall exchange of Na+ by Ag+ ions had only proceeded to -25 %, it would be expected that no dramatic change in the neutralization pattern of the studied hydrated structures will occur.The unit cell dimension is constant throughout the sequence of samples, which indicates the absence of any distortive influence of the Agf ions.M . COSTENOBLE AND A . MAES 133 The interpretation of the electron densities around the O3 oxygens is based on 17.5 charges, as found in the Na-Y parent material. The differentiation between Ag+ and Na+ is obtained according to the constancy of the number of charges around each oxygen4 and taking into account that the electron density of 1 Ag+ ion is similar to 4.5 Na+ ions. The Ag+ to Na+ conversion factor arises from the interpretation of the respective atomic scattering factors in the sin O/h range scanned in the experiments. The electron density observed at site I in the Na/2Ag-Y structure can be assigned to 1.9 Ag+ ions or 8.6 Na+ ions.However it is improbable that the introduction of 2 Ag+ ions results in a migration of 8.6 Na+ ions to site I, since in hydrated Na-Y lo no electron density was observed in the hexagonal prism (see table 1). TABLE AS ASSIGNMENT OF ELECTRON DENSITIES AT THE DIFFERENT SITES AND INTERATOMIC DISTANCES IN NM a site I site I’ site 11‘ site I1 total localized Ag chem. determined Ag unlocalized Ag total charge on sites I +I’ + I1 Tnteratomic distances site I’-Os - 0 2 -site I’ -site 11’ site 11’-O2 - 0 4 -site I‘ -site 11’ site 11-02 4 3 4 -site 11’ site I-O3 Rr a0 Na-Y - 1 7.3( 1 .3)Na+ 13.4( 1.1)HzO 10.2( 1.6)Na+ - - I 27.5 0.25 3 (9) 0.307(10) 0.3 8 8( 1 8) 0.259( 16) 0.283(13) 0.372(13) 0.259( 16) 0.343(25) 0.279( 13) 0.3 19( 13) 0.346( 18) 0.278(3) 0.1658 2.4700(3) Na /2Ag-Y 1.9(2)Agf 1 5.7( 1 .O)Na+ 15.7(1.3)HzO 9.7(1.9)Na+ 1.9 1.9 27.3 - 0.255(6) 0.309(6) 0.3 89( 12) 0.247( 12) 0.296( 11) 0.388(11) 0.247( 12) 0.298 (22) 0.253(20) 0.307(20) 0.342(23) 0.279(2) 0.1535 2.4690(4) Na17.25Ag-Y 3.9(4)Ag+ 13.5( 1.7)Na+ 0.3(4)Agf 8,0(2.7)Na+ 1.8(6)Ag+ 6.0 7.25 1.25 27.50 23.5(3.2)HzO 0.250(5) 0.311(5) 0.392(8) 0.249(6) 0.249(6) 0.384(6) 0.249(6) 0.253( 8) 0.302( 8) 0.352(9) 0.277(3) 0.1747 2.4690(2) - Na/ 1 4Ag-Y 4.4(6)Ag+ 1 1.7(9)Na+ 1.8(4)Ag+ 4.0(3)Na+ 6.0(7)Ag+ 12.2 14.0 1.8 27.90 18.3(3)H20 0.247(4) 0.312(4) 0.3 8 8 (6) 0.249(6) 0.296(5) 0.3 8 1 (6) 0.249( 6) 0.307(9) 0.255(5) 0.303(5) 0.340( 6) 0.272(2) 0.1376 2.4690(4) a Estimated standard deviations on the last figures in parenthesis.Upon increasing the Ag content to 7.25 ions/u.c. the observed electron density on site I is equivalent to 3.9 Ag+ ions. Using the same argument as for Na/2Ag-Y it is physically impossible to attribute the electron density in site I to 17.6 Na+ ions, since the exclusion rule (I+$ I’ < 16) cannot be obeyed in this case. According to the premise that w 17.5 charges are located around the O3 oxygens in all the hydrated samples, and assigning the entire electron density in site I to 3.9 Ag+ ions, the distribution in site I’ must be 13.5 Na+ and 0.3 Ag+ ions. The observed electron density on site 11 being equivalent to 16.1 Na+ ions, exceeds that for the Na/2Ag-Y and Na-Y and is thought to be due to the presence of silver.I34 SITE GROUP INTERACTION I N ZEOLITES By analogy with Na-Y lo and Na/2Ag-Y where 10 charges (Na+ ions) (see table 1) were located on site 11, 8 Na+ ions and 1.8 Ag+ ions are assigned to site 11.On the basis of the preceding arguments the entire electron density on site I is attributed to 4.4Ag+ ions in the Na/l4Ag-Y structure. Taking account of the standard deviation it is seen in table 1 that the observed distances in all the samples are very similar to the Na-Y. The 1’-03 distance (0.247+_0.004nm) in the Nal 14Ag-Y sample has a tendency to be even smaller than in the other structures, suggesting that minor amounts of Ag+ are present on that site. Since, moreover, it was assumed that the charge distribution/oxygen ratio will only slightly vary, 1.8 Ag+ and 11.7 Na+ ions are assigned to site 1’.The sum of the charges around 0, then amounts to 17.9 units ofcharge/u.c. (see table 1). The high electron density observed on site I1 suggests a major Ag+ occupancy. Keeping the number of charges on site I1 equivalent to 10 the observed electron density is satisfied by 4 Na+ and 6 Ag+ ions, In Part 2 the selectivity of Ag+ for Na-Y will be studied.’ The high preference for Ag+ in the early stages of the Na+-Ag+ exchange in the presence (ternary system) or absence (binary system) of a large excess of Cs or NHZ, will be considered to originate from a high preference for a few Ag+ ions in the small cages, the ideal coordination in the hexagonal prism being the most probable position. From the foregoing X-ray assignments in the sequence of samples studied it emerges that site I is preferentially occupied by Ag+ ions at low loadings of silver in the zeolite.Beyond a limit of ~4 Ag+ ions, additional introduction of Ag+ into the zeolite fails to increase site I occupancy, but results in a distribution over different sites 1’, I1 and the unlocalized part according to the preference pattern : site I1 > site I’ 21 unlocalized. The preference of Ag+ for site I may originate from its large radius. The ideal distance Ag+-W- = 0.266 nm, which is close to the observed distance site I-O3 of 0.275 nm. However, K+ which is theoretically still better suited (Kf-02- = 0.273 nm; 1’-03 = 0.280 K-Y)4 to fit the hexagonal prism has an insignificant occupancy on site I.4 The higher hydration energy of Agt- and its ability in general to form complex ions must therefore favour the formation of a type of coordination compound with the 6 oxygens of the hexagonal prism.Sodium forms a stable coordination complex with H20 in the sodalite units lo and does not occupy site I. Silver with its still higher than Naf hydration energy is reluctant to remain in the sodalite cages. The diameter of the Ag+ ion being 0.06 nm larger than Na+ ion, it is sterically hindered in its coordination with site 1‘, while site 11’ is occupied by the H20, and therefore Ag+- prefers site 11. The occupancy in the small cage sites seems to be a subtle inter-play between hydration and coulombic forces governed by steric factors. As a consequence, the distribution of exchangeable cations over small and large cavities fails to follow a fixed pattern. The specific behaviour of any kind of ions or mixture of ions along with the pore system characteristics of zeolite-Y result in a neutralization which cannot solely be discussed in terms of small and large cavities.The authors thank the Belgian Government (Programmatie van het Wetenschaps- beleid) for financial support. H. S. Sherry, Adu. Chem. Ser., 1971, 101, 350. B. K. G. Theng, E. F. Vansant and J. B. Uytterhoeven, Trans. Firday Soc., 1968, 64, 3370. E. F. Vansant and J. B. Uytterhoeven, Adu. Chern. Ser., 1971, 101,426. W. J. Mortier and H. J. Bosmans, J. Phys. Chem., 1971, 75, 3327. W. J. Mortier, H. J. Bosmans and J. B. Uytterhoeven, J. Phys. Chem., 1972, 76, 650. A. Maes and A. Cremers, J.C.S. Faraday I, 1975, 71, 265.M. COSTENOBLE AND A . MAES 135 ’ A. Maes and A. Cremers, J.C.S. Faraday I, 1978, 74, 136. A. Maes and A. Cremers, Adu. Chem. Ser., 1973, 121,230. J. V. Smith, Ado. Chem. Ser., 1971, 101, 171. lo M. L. Costenoble, W. Mortier and J. B. Uytterhoeven, J.C.S. Faraday I, 1976, 72, 1877. l1 W. C. Hamilton, POWOW (Brookhaven National Laboratory, Brookhaven, New York, 1962). l2 W. R. Buning, K. 0. Martin and H. A. Levy, ORFFE (Oak Ridge, National Laboratory, l3 D. H. Olson, J. Phys. Chem., 1970, 74, 2758. l4 W. M. Baur, Amer. Mineral., 1964, 49, 697. l5 J. J. Pluth and J. V. Smith, Mar. Res. Bull., 1973, 8,459. l6 T. H. Hseu, Ph.D. Thesis (University of Washington, 1972). Oak Ridge, Tenn., 1964). (PAPER 6/1919)

ISSN:0300-9599    

DOI:10.1039/F19787400131

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Site Group Interaction Effects in Zeolite-Y Part 2.-Na-Ag Selectivity in Different Site Groups BY ANDRE MAES* AND ADRIEN CREMERS Katholieke Universiteit Leuven, Centrum voor Oppervlaktescheikunde en Colloi'dale Scheikunde, De Croylaan 42, B-3030 Heverlee, Belgium Received 13th October, 1976 The influence of the nature of the ion neutralizing a certain site group on the selectivity in another group is studied in zeolite-Y. The high Ag to Na preference in the small cages in the binary system is reduced by filling the large cavities with Cs+ or NH: ions. The selectivity data are fitted in terms of the occupancy and selectivity coefficient of 4 (binary system) and 2 (ternary system) types of sites. The calculated occupancies in each site agree with the observed ones, which are obtained from X-ray studies.The thermodynamic ion exchange formalism for the case of several groups of sites is discussed. One of the characteristic features of most zeolitic ion exchangers is the association of the framework charge with various types of cationic sites. Very often, the number of available crystallographic sites exceeds the number of negative charges to be neutralized and changes in the occupancy factors are expected, depending on the nature of the neutralizing cati0ns.l. Such site heterogeneity, combined with the fact that ion exchange reactions are rarely thermodynamically ideal and that certain ions may only be partially exchanged, for steric or other reasons, leads to ion exchange isotherms of great complexity which are often difficult to rationalize in terms of the properties and extents of the various site groups.In a few cases, ion exchange isotherms have been sucessfully analysed in such terms 3 9 and these analyses indicated that a given group of sites may be assigned a characteristic set of thermodynamic state functions for some arbitrary pair of cations. In fact, this is the point taken by Barrer and Klinowski in their theoretical approach to the case of several groups of homogeneous sites. Admittedly, these authors recognize that each group of sites may fail to behave ideally, and this is formally taken care of by assigning to each group a characteristic non-ideality pattern of its own. By an appropriate choice of combinations, they were able to produce some complex Kielland plots which agreed closely with experimental data.More recently, Barrer, Klinowski and Sherry extended this approach to cases in which part of the zeolitic ions are not exchangeable. One of the most crucial points in these papers is the assumption of the existence of a given equilibrium constant for some arbitrary exchange reaction in a given set of sites. The purpose of the present paper is twofold: first, to elaborate on some conceptual difficulties involved in formalizing thermodynamically the exchange of ions in ion exchangers with interacting site groups, and secondly to show experi- mentally that such site group interactions do in fact exist. Synthetic zeolite-" is an excellent system for comparing a given exchange reaction in the small cages in the absence and presence of a cation which is unable to penetrate these areas of the crystal.In particular, we present a study of the Naf-Agf equilibrium, in the presence of a large excess of cesium ions, which confine the 136A . MAES AND A . CREMERS 137 Na+-Ag+ equilibrium to the small cages, and show that such an equilibrium is significantly different from what is observed in the binary system. THEORETICAL Confining, for simplicity, our attention to monovalent ion exchange reactions, the overall thermodynamic constant for the ion exchange reaction between a solid (z) and a solution (s) is defined as &+A, f A,+B, (1) Azf K , = - Bzf BaA' A,, B, represent the equivalent fractions of the ions in the exchanger and fA/fB the activity coefficient ratio in the solid; aB/aA represents the activity ratio in solution which, at sufficiently low ionic strength, may be identified with the molality ratio.The overall selectivity coefficient K , is defined by B A Kc = Ktf If (3) As usual, the standard and reference states are the homoionic solids, in equilibrium with an infinitely dilute solution of the corresponding ions. In the case of an exchanger containing n types of site, Barrer and Klinowski and Barrer, Klinowski and Sherry define an equilibrium constant for the ith set by in which At, BL represent the equivalent fraction of ions in set i andf:lfy the activity coefficient ratio, which formally takes care of the non-ideality effects within the corresponding site group. Evidently, if Xi represents the equivalent fraction of the ith set (with respect to the ion exchanger capacity), then n n A, = XiAi and B, = XiBi.1 1 The individual thermodynamic constants are related to the overall thermodynamic constant by n n K , = n K f i or AGO = zXiAGP ( 5 ) 1 1 in which AGO stands for the overall standard free energy change and AG; for the standard free energy change within the ith set. Implicit in the foregoing approach is the definition of the chemical potential pq in the ith set: ,up = ',u;+RT 1dA'f A. ' (6) Of course,fq -+ 1 when At -+ 1. The activity coefficient f:, or more generally, the activity coefficient ratio ftfl? within the ith set is some complex function of the ion exchanger composition and carries the burden of all interaction effects resulting from changes in composition between ions positioned in sites belonging to all different groups.From a rigorous thermodynamic point of view, this approach may be criticized138 SITE GROUP INTERACTION I N ZEOLITES on the following grounds. Eqn (5) is based upon the view that the standard free energy content of one equivalent of zeolite may be expressed as a sum of terms, GP, referring to the standard free energy contents of the various site groups (including the interaction terms between the ions positioned in the site group concerned). However, such a sum should contain a number of cross terms referring to interaction energies between different site groups, and these terms are not necessarily identical for two different ionic forms. Therefore, the meaning of the concept of a standard free energy content of a given site group is not clearly delineated and is ambiguous.The difficulties may be analysed in a more detailed fashion. In studying the thermodynamic properties of a given set of sites, one is assigning to it a thermodynamic status of its own, and one is therefore faced with the logical consequence that the activity coefficients are interrelated through the Gibbs-Duhem equation. This cannot be expected to be the case for the following reason. In attempting to express the composition dependence of the activity coefficients, one has a choice between two alternatives : the first is to express it as a polynomial in terms of composition changes of all different sets ; the second alternative, apparently taken by Barrer and Klinowski, is to rely on a ‘‘ simple concentration dependence ”, for example in terms of composi- tion changes in the site group concerned.The second alternative may be justified in view of the fact that all site groups are in equilibrium with each other, (being in equilibrium with the liquid phase) so that a given composition in one site group automatically fixes the composition of all other groups. Whichever choice is made, the coefficients in such a polynomial must carry the interdependence of composition changes between the different site groups. This interdependence, however, is primarily governed by differences in affinities of the competing ions for the different site groups and not by cation-cation interactions. It is therefore to be expected that the activity coefficients are not interrelated through the Gibbs-Duhem equation.These coefficients, and consequently the thermodynamic equilibrium constant for a particular site group, are not to be considered as thermodynamic quantities in the true sense. There may be no reasonable alternative to the one presented by Barrer and Klinowski and the case of interacting site groups may not be treated in a rigorous thermodynamic manner altogether. In any case, the foregoing criticism should not prevent an attempt to analyse ion exchange equilibria in terms of “equilibrium constants ” or preferably selectivity coefficients for the various site groups. One should keep in mind, however, that these quantities only yield quantitative measures of differences in ion-site affinities and are not to be identified with thermodynamic quantities in the formal sense.EXPERIMENTAL The zeolite used in this study is the Na-Y faujasite sample of unit cell formula Na54A154Si1380384 241 H20. Preparation methods and details of experimental procedure can be found el~ewhere.~ The Ag+-Na+ binary equilibria were studied at 25°C and 0.01 mol dm-3 total concentration, and the silver and sodium distributions were monitored using y-emitting isotopes llomAg and 22Na. The Ag+-Naf equilibria in the presence of cesium and ammonium ions were studied using the following procedure : first, a separate study was made of the ion exchange behaviour of Cs+ in Na-Y at high Cs+ loading in order to verify the maximum exchange limit and selectivity behaviour. The Ag+-Na+ equilibria were then studied in the presence of an excess of 3.55 x lo3 and 1.8 x lo3 ions/u.c.Csf ions and an excess of 7.1 x lo3 ions/u.c. NHZ ions. The corresponding molarities in the equilibrium solution are 0.3 and 0.15 (Cs+) and 0.6 (NH1;) mol dm-3. Preliminary tests showed this excess to be sufficient to exclude both Na+ and Ag+ from the supercages and confine these ions exclusively to the small cages.A. MAES AND A . CREMERS 139 RESULTS AND DISCUSSION The silver-sodium binary equilibria and those in the presence of a high excess of cesium or ammonium are summarized in fig. 1, which shows the variation of In K,(Ag+-Na+) as a function of the overall silver loading. It should be emphasized that, in calculating selectivity coefficients, the Na+/Ag+ molality ratios have been identified with the activity ratios in both binary and ternary systems.This simplification requires no comment in the binary study which was carried out at 0.01 mol dm-, total concentration, but requires some justification for the case in the presence of a high excess of CsN0,. The problem relates to obtaining activity coefficients of traces of AgN03 and NaN03 in a relatively concentrated CsNO, solution (0.15 or 0.3 mol dm-3). Harned and Robinson * showed that the activity coefficient of a trace of a given salt in an excess of another salt is intermediate 4 G 3 22 C - 2 \ I I I 025 050 0.75 ZA, FIG. 1.-Surface composition dependence of in $K, in the binary system (O), and in presence of an excess of 3.55 x lo3 (O), 1.8 x lo3 (A) Cs+ ions/u.c., and 7.1 x lo3 (A) NH: ions/u.c. between the activity coefficient of its own solution and the activity coefficient of the other salt at the same ionic strength, but generally closer to the latter.Taking the 0.3 mol dm-3 CsNO, case as an example, the activity coefficients of the three electro- lytes at 0.3 ionic strength are 0.602 (CsNO,), 0.606 (AgNO,) and 0.666 (NaNO,). It is apparent that the behaviour of AgN0, and CsNO, is entirely similar and that, therefore, the trace activity coefficient of AgN0, in a 0.3 mol dm-, CsNO, solution will be given by the numbers shown. For NaNO,, the activity coefficient of which differs by some 10 % from the CsNO, value, its trace activity coefficient is expected to differ by < 5 % from the value of CsNO,, in view of the argument presented above. Therefore, the Na+/Ag+ molality ratio may safely be identified with the activity ratio.The standard free energy for the Ag+-Na+ system, as obtained by graphical integration,1° is - 5.43 kJ mol-l. This figure shows that the silver-to-sodium]I 40 SITE GROUP INTERACTION IN ZEOLITES selectivity is significantly lowered in the presence of Cs+ or NH: ions, the effect being roughly the same in both cases. The ion exchange data for Cs+ in Na-Y are summarized in fig. 2, the upper part of which shows that the maximum exchange limit is 68 %, i.e. some 36.5 ions/u.c., a result which is in agreement with the one obtained by Sherry.ll The lower part of this figure shows the normalized selectivity coefficient log K: (Cs+-Na+) at high cesium loadings. These data are also in excellent agreement with those of Sherry, i.e.the limiting value of K,” (Cs+-Na+) at high Cs+ loading is ~ 3 . SCS 0.8 0.9 I I I 0.5 - k2 M 0.4- 0 - ZCS FIG. 2.-Na+-Cs+ exchange isotherm showing the limit of maximum Cs+ exchange (upper part) and the log E K c as a function of the normalized surface composition (lower part). Data of Sherry (solid line) are shown for comparison. Inspection of fig. 1 shows clearly that the Ag+ to Naf selectivity in the presence of Cs+ or NHZ ions is significantly lower than in the binary system. Before analysing these differences quantitatively, it is imperative to show that, in the presence of these other ions, the Ag+-Na+ exchange is restricted to the small cages. A first indication is provided by the data in fig. 3, which shows that the sum total of Ag++Na+ never exceeds the limit of 17.5 i0nsju.c.The small cage occupancy in the presence of NH: is lower and may be understood in terms of a slightly different neutralization pattern in small and large cavities.I2 A second indication is to be found in the Cs+ selectivity at high loading. As far as Na+ is concerned, the Cs+/Na+ ratio in the equilibrium solution is x 100 (in the case of a Cs+ excess of 3.55 lo3 ions/u.c.) irrespective of the amount of Ag+ added. Since the K, (Cs+-Na+) value is m3 at high Cs+ loading (see fig. 2) it follows that the Na+ occupancy in the supercages is 0.15 i0nslu.c. at most. When the Cs+ excess is reduced to 1.8 x lo3, this would amount to some 0.3 ions/u.c.A . MAES AND A . CREMERS 141 The Ag+ occupancy in the supercages can be estimated on a similar basis. On the basis of calculations to be presented below, the selectivity coefficient K, (Cs+-Ag+) in the supercages is near unity.Therefore, at low Ag+ loading, when the Cs+/Ag+ ratio in the equilibrium solution is very high (=lo4), the Ag+ occupancy in the supercages is entirely negligible. At higher Ag+ loadings, the Cs+/Ag+ ratio in the equilibrium solution is 200 to 300, which would correspond to 0.3 Ag+ ions/u.c. Similar calculations can be made in the case of NHZ; the selectivity coefficient KF(NH+,-Na+) at high loading is =4,l2 and the NHZ excess is even larger. Consequently the involvement of the supercage in the Ag+-Na+ exchange is restricted to roughly 1 % of its capacity and we may be confident in analysing the data in the presence of Cs+ or NHZ in terms of ion competition for the small cage sites.I I I I I 2 4 6 8 Aga& ions1u.c. FIG. 3.-Determination of the small cage capacity in an excess of 0, 3.55 x lo3 Cs ions1u.c. ; and A, 7.1 x lo3 NH; i0nslu.c. In order to analyse the selectivity data quantitatively, we used the X-ray data, presented in the Part 1 ; as a guideline, we assume that the total charge in the small cages is 17.5 and that the maximum number of silver ions which can be accommodated in site I is =4 (no Na+ ions are found in site I).2 In attempting to assign the remaining 30.5 ions to the supercages, we may choose from two alternatives : 10 ions/u.c. can be assigned to site 11, leaving the remaining 26.5 ions unlocalized. This choice could be justified on the basis of the finding of X-ray work, that no more than 10 Na+ ions could be localized in site IL2 An excellent fit of the Ag+-Naf binary data, as shown in fig.4, is then obtained by using the K , data listed in table 1. Another alternative would be to assign 20 ions/u.c. to site 11. The motivation for this choice is found in the fairly high selectivity of Ag+ ions in the supercages as indicated by the data of Theng, Vansant and Uytterhoeven l2 and the existing parallelism between selectivity and extent of ion 10calization.l~ This possibility is confirmed by the fact that in Y-zeolite, more K+ ions are localized, 20 in site IIY1 than Na+ ions.2 This1 42 SITE GROUP INTERACTION I N ZEOLITES choice necessitates a slight shift in the K, value assigned to I’ and I1 and gives an equally good description of the data. In any case, the shape of the Kielland plot at increasing Ag+ loading is not very sensitive to the choice of combinations for the supercage and it would be logical to expect shifts in the course of the exchange.Either choice leads to values for the overall AG values [using eqn (5)] which are in 1 I I I I 5 10 15 Ag i0nslu.c. FIG. 4.-Comparison of the experimental (same symbols as in fig. 1) and calculated selectivities as a function of the number of Ag+ ions/u.c. in the binary (solid line = fit a ; broken line = fit b) and the ternary system. excellent agreement with the result obtained by integration: -4.94 (fit 1) or -5.28 (fit 2) kJ mol-l. The values are slightly more negative than those obtained by Sherry,ll which were derived only on the basis of data at high loading. The data in table 1 are used to calculate the way in which the various site groups are occupied by Ag+ ions as the silver content is increased in the equilibrium solution.The results of these calculations are summarized in fig. 5 which shows the relative occupancy of the site groups as a function of the overall silver loading. Apparently, TABLE 1.-Two SETS (a AND b) OF PARAMETERS (SELECTMTY COEFFICIENT Kc AND MAXIMUM NUMBER OF IONS IN EACH SITE GROUP/U.C.) USED TO FIT THE BINARY Na+-Ag+ EXCHANGE EQUILIBRIUM. THE FRACTION OF TOTAL CAPACITY (Xi), In Kc AND THE FREE ENERGY ASSOCIATED WITH THE RESPECTIVE SITE GROUPS (- Xi RTln Kc) ARE ALSO TABULATED. ma. -Xi RT In Kc site KC ions1u.c. (Xc) In Kc jkJ mol-1 I a 2000 4 b 2000 4 I’ a 2 13.5 b 4 13.5 I1 a 20 10 b 10 20 unloc. a 4 26.5 b 4 16.5 0.074 0.074 0.25 0.25 0.185 0.37 0.49 0.25 7.601 7.60 1 0.693 1.386 2.996 2.303 1.386 1.386 1.41 1.41 0.44 0.88 1.39 2.14 1.70 0.88A .MAES AND A . CREMERS 143 the most important conclusion to be drawn from these data, is the very high affinity of silver ions for site I. An illustration of the nature of the agreement of X-ray data with those obtained by least squaring the ion exchange data is shown in table 2. The K, values listed in the tables represent only a quantitative measure of ion site affinities. However, it is worth emphasizing that the K, value, assigned to site I, 0 50 0.40 I00 x +.I ." 0.30 0.75 8 Q +a .C1 v) 4-r 0 .C1 Y 0 0.20 (d Ct: 0.10 425 5 10 15 Ag ions/u.c.FIG. 5.-Fraction of each site occupied by Agf- as a function of the total Ag+ content in ions/u.c. in site I (A, A), site I' (0, 0), site I1 (a, 0) and unlocalized (+, 0). Filled and open symbols respectively refer to the first (fit a) and second assumption (fit b). is ambiguous, as it really refers to a displacement reaction of Na ions from site I' by silver ions taking up positions in site I. The Ag+-Na+ equilibrium data for the small cages in the presence of Cs+ or NH: are impossible to fit in terms of the same occupancy numbers used for the binary data, and a significant shift of Ag+ ions has to be postulated from site I -+ 1'. The TABLE 2.-cOMPARISON BETWEEN THE POPULATION OF Ag+ IONS IN EACH SITE GROUP DETER- MINED BY X-RAY DIFFRACTION TECHNIQUES AND CALCULATED BY USE OF THE PARAMETERS OF FIT a AND b I I' I1 U total Ag+/ (ions/u.c.) tit a fit b X-ray fit a fit b X-ray fit CI fit b X-ray fit a fit b X-ray 2.0 2.00 2.00 1.9 0.01 0.03 - 0.1 0.1 - 0.05 0.03 - 7.25 3.84 3.84 3.9 0.3 0.61 0.3 1.90 2.05 1.8 1.19 0.74 1.25 14.00 3.92 3.92 4.4 1.16 1.89 1.8 4.8 5.9 6.0 4.08 2.31 1.8144 SITE GROUP INTERACTION I N ZEOLITES results are sumniarized in table 3.It is apparent from the nature of the fit, shown in fig. 4, that the parameters shown in table 3, give a close description of the ion exchange data. Two important observations can be made at this point : first, the presence of Cs+ (or NH;) ions in the supercages leads to a significant reduction in the silver affinity for site I ( ~ 4 kJ mol-I). The free energy loss associated with the exchange of Na+ for Ag+ ions in the small cages in the binary system can be calculated from the parameters in table I and is compared in table 3 with the data from the ternary system.The Ag+ to Na+ preference in the small cages is reduced from 5.72 kJ (fit a) or 7.05 kJ (fit b) to 4.6 kJ mol-I when the large cavities are respectively occupied by mostly Naf and Cs+ ions. Therefore, it appears that the free energy change for a given ion exchange reaction in a particular set of sites, (leaving aside the basic difficulties treated in the theoretical section) is very much dependent on the composi- tion of other site groups. Perhaps the nature of the effects described here can be qualitatively understood in terms of a strong interaction between the poorly hydrated TABLE 3.-cOMPARISON OF THERMODYNAMIC DATA FOR THE Na+-Ag+ EXCHANGE IN THE SMALL CAGES ONLY IN THE BINARY AND TERNARY SYSTEMS fraction of max.small cage -Xi R T h Kc site K, i0nslu.c. C.E.C. in Kc /kJ mol-1 ternary I 400 1.75 0.1 5.99 1.49 I‘ 4 15.75 0.9 1.356 3.11 binary I a 2000 4 0.229 7.601 4.37 b 2000 4 0.229 7.601 4.37 I’ a 2 13.5 0.774 0.693 1.34 b 4 13.5 0.774 1.386 2.70 and very polarisable Cs+ (and perhaps NH4) ions and the oxygen framework; this may lead to a weakening of the coordination bond of the silver ions in the hexagonal prisms. The second important aspect is the significant shift of Ag+ ions from the hexagonal prisms to the cubo-octahedra (I + 1’) upon introducing Cs+ ions into the supercages. Evidently, there is some logic in this since, as explained on theoretical grounds by Mortier,14 a change in energy difference between two sites, I and I’ in our case, should be reflected by a corresponding change in the ratio of occupancy numbers at these sites.In conclusion, we have presented experimental evidence that a change in the composition of the crystal may lead, as has sometimes been invoked in the past, to a relocation of some ions from one type of site to another. In addition, it is shown that the selectivity pattern for a particular site may be affected by the nature of the ions which are used to neutralize the electrical charge in other site groups, as was hypothesized in a recent paper.15 Acknowledgment is made to the Belgian Government (Programmatic van het Wetenschapsbeleid) for financial support. W. J. Mortier and H. J. Bosmans, J. Phys. Chem., 1971, 75, 3327. M. L. Costenoble, W. Mortier and J. B. Uytterhoeven, J.C.S. Paraday I, 1976, 72, 1877 E. Gallei, D. Eisenbach and A. Ahmed, J. Catalysis, 1974, 33, 62.A . MAES A N D A . CREMERS 145 * R. M. Barrer and W. M. Meier, J. Inoug. Nuclear Chem., 1966,28, 629. R. M. Barrer and J. KIinowski, J.C.S. Faraday I, 1972, 68, 71. R. M. Barrer, J. Winowski and H. S. Sherry, J.C.S. Faraday II, 1973, 69, 1669. A. Maes and A. Cremers, Adv. Chem. Ser., 1973,121,230. R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 1959). * H. S. Harned and R. A. Robinson, Multicomponent Electrolyte Solutions (Pergamon, 1968). lo G. L. Gaines and H. C. Thomas, J. Chem. Phys., 1953,21,714. l 1 H. S. Sherry, J. Phys. Chem., 1966,70,1158. l2 B. K. G. Theng, E. F. Vansant and J. B. Uytterhoeven, Trans. Faraday Soc., 1968, 64, 3370. l3 M. Costenoble and A. Maes, J.C.S. Faraduy I, 1978, 74, 131. l4 W. J. Mortier, J. Phys. Chem., 1975, 79, 1447. A. Maes and A. Cremers, J.C.S. Faraday I, 1975, 71,265. (PAPER 6/1920)

ISSN:0300-9599    

DOI:10.1039/F19787400136

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Active Centres on NaH-Y Zeolite in But4 -ene Transformations BY JAN GALUSZKA," ANDRZEJ BARA~SKI AND STANISLAW C~CKIEWICZ 30-060 Cracow, Poland Institute of Chemistry, Jagellonian University, Krupnicza 41, Received 17th November, 1976 The kinetics of but-1-ene transformations on zeolites NaH-Y with varying degrees of Na+ ion exchange (0-82 %) have been investigated at constant pressure and also by temperature-programmed desorption combined with gas liquid chromatography. A correlation between the degree of cation exchange and localization of the Bronsted acid centres was noted and an assignment of the active centres involved in the but-1-ene transformations is proposed. Hydrogen sodium zeolites obtained from Na-Y parent specimens by ion exchange are useful for various catalytic processes, such as transformations of but-1-ene as studied by the temperature programmed desorption TPD meth0d.l The experimental data from that study were used as a basis for the following scheme of consecutive transformations of but-1-ene on NaH-Y zeolite.but-1-ene(gas) -+ adsorption -+ isomerization -+ polymerization dismutation cracking " further reactions ". (1) - - called for brevity 4 \ dehydrogenation The aim of the present paper is mainly to identify the active sites involved in butene transformations. The distribution of potentially active Na+ and H+ ions among the possible accessible positions in the crystal lattice of zeolite NaH-Y is still an open question despite many studies.2 However, a comparison of reaction rates for catalysts con- taining different amounts of different accessible sites, combined with a temperature programmed desorption study, should throw some light on the location of the active centres in the zeolite crystal lattice.EXPERIMENTAL MATERIALS Na-Y zeolite of (Si02/A1203 molar ratio 5.22) was kindly supplied by the Institute of Industrial Chemistry, Warsaw. Samples of NH4Na-Y zeolite were obtained by Na+/NHz ion exchange followed by pelletizing without binder. For the various samples prepared, degrees of exchange 0 ; 9.5 ; 12 ; 23 ; 43 ; 59 ; 68 ; 77 and 82 % were determined by flame photometry. The pellets of NH4Na-Y were stored over saturated aqueous NaCl at room temperature. Water content in the zeolite estimated derivatographically amounts to - 21 %. But-1-ene of purity 99.8 mol % supplied by Fluka AG contained isobutane, isobutene and butadiene as commercially listed impurities.Commercial cylinder nitrogen was used as carrier gas after deoxidation over BTS catalyst (B.A.S.F., FRG) and passage through silica gel and liquid nitrogen-cooled traps. 146J. GALUSZKA, A. B A R A ~ S K I AND s. C~CKIEWICZ 147 APPARATUS The TPD apparatus and the manostat system were as described previously. PROCEDURE In order to avoid possible deactivation fresh zeolite pellets (5 mm in diameter, 3.5-4 mm thick, 0.1 g average weight) was used in each experiment. NH4Na-Y zeolite was decomposed in situ in a TPD reactor at 430°C for 4 h in vacuo ( p < 1.3 x N m-2 = Torr). NaH-Y zeolite samples, obtained after calcination, were kept in a reactor in vacuo for 24 h and then again outgassed for 1 h at 430°C.The reactor was then cooled down to room temperature, when sorption of but-1-ene from gas phase was carried out at constant pressure 693-746f 7 N m2 (5.2-5.6kO.05 Torr). The kinetic curves of butene consumption were automatically recorded. After the kinetic run, but-1-ene remaining in the gas phase was removed from the reactor by nitrogen flushing for 15 min. It was proved additionally that the removal of gaseous but-1-ene by outgassing for 1 h using a diffusion pump introduces no change in the TPD spectra. Accordingly, the simpler flushing procedure only was used. ,4fter removing the gas phase, temperature programmed desorption was followed up to 430°C using a heating rate of 25f 1 deg min-'. The stream of carrier gas containing the desorbed species was monitored by a Gow Mac thermal conductivity cell, after which condensable products were frozen out in a liquid nitro- gen-cooled trap.These products, again after evaporation, were transferred by a syringe into a gas chromatograph. The packing of the chromatograph column consisted of chromo- sorb W (80-100 mesh size) covered with 15 % of 2,4-dimethylsulpholane. RESULTS THE KINETIC CURVES Fig. 1 shows the kinetic curves obtained for the system NaH-Y zeolite + but-1-ene using preparations with various degrees of exchange of Na+ ions for protons. Fig. 1 6.5 t 36 of cc!ions exchanged min degree of ion exchange as shown in the figure. FIG. 1.--Kinetic curves for the system NaH-Y zeolite+ but-1-ene. Zeolites amples of varied148 ACTIVE CENTRES ON ZEOLITES 1 I 1.0; I 0 .5 1 FIG. 2.-Relation between % of cations exchanged the amount of but-l-ene consumed by the samples after (a) 5 and (6) 100 min and the degree of ion exchange. contains many intercrossing curves, making discussion difficult, so we direct attention to the simpler plots in fig. 2, derived from fig 1, which show the dependence of the amount of but-1-ene consumed upon the degree of ion exchange for two arbitrarily chosen times : 5 and 100 min. The shape of the plots is thelsame for any time between t 120 I temperature/"C FIG. 3 . T P D evolution spectra as they depend upon the exchangelextent1of zeolite NaH-Y samples.J. GALUSZKA, A. B A R A ~ S K I AND s. C ~ C K I E W I C Z 149 2 and 100min. The results at times < 2min are uncertain, for experimental reasons.From a comparison of fig. 1 and fig. 2 it can then be seen that the depend- ence presented in fig. 2 is valid for conditions close to equilibrium. It follows that diffusion phenomena do not affect the shape of the plots in fig. 2. We conclude that the shape of these plots illustrates an intrinsic property of the investigated system. The sequence of kinetic curves and the results presented in fig. 2 show the existence of three characteristic intervals related to the degree of ion exchange. From the SiOz/Al,03 molar ratio it was deduced that there are 53.2 sodium cations per unit cell. So interval I from 0 to 12 % is equivalent to 6.4 Na+ cations per unit cell : interval I1 from 12 to 68 % is equivalent to 29.8 cations per unit cell and interval 111 above 68 % is equivalent to 17.0 cations, assuming 100 % of exchange as the upper limit. TPD SPECTRA The TPD spectra are presented in fig.3, where a pronounced dependence of the size and shape of particular peaks upon the extent of exchange is easily seen. Trends in the temperature for the maximum of peak I can also be seen from this figure. The change of the trend is particularly marked at the end of interval I. The qualitative composition of the desorbate presented previously was confirmed by chromatographic analyses of desorption products. Especially important is the fact that in the temperature range of peak I only but-l-ene was found for unexchanged zeolite, whereas from the all other samples three butene isomers, and only these species, are desorbed in the same temperature range.DISCUSSION The sequence of consecutive reactions, given in the introduction, is taken as the starting point for the following discussion. It was inferred from chromatographic analyses of desorbed products after TPD experiments that, on unexchanged zeolite, only adsorption of but-l-ene takes place ; on 9.5 and 12 % exchanged samples isomerization also proceeds, whereas on samples exchanged to a higher extent further butene reactions additionally occur. These analyses agree with data from our previous paper.l It follows that at least three different types of active sites should be considered : adsorption sites (A), isomerization sites (B) and sites responsible for further butene reactions (C). Unexchanged zeolite contains only A sites; 9.5 and 12 % exchanged samples contain A and B sites; samples exchanged to greater extent contain A, B and C sites.The maximum at 68 % (fig. 2) may suggest the existence of a fourth type of active site. W'e consider first the adsorption step of the sequence (1). But-l-ene is adsorbed at the highest rate on unexchanged zeolite (fig. 1 and 2). When the degree of exchange increases, the amount of desorbed but-l-ene decreases as evidenced by chromato- graphic analyses. In addition, peak I, and hence the total amount of butene isomers, also decreases (fig. 3). Since Bronsted centres are absent from unexchanged zeolite, their amount increasing with the degree of exchange, it seems unlikely that they are responsible for the adsorption of but-l-ene.For the same reason, those Lewis centres which are produced from Bronsted centres via dehydroxylation can be ex- cluded, especially as the samples were not heated above 430°C.6* ' Possible species obtained after the adsorption step of the sequence (1) which must be considered are physically adsorbed but- l-ene and chemisorbed molecules on Na+-sites. Arguments against physical sorption are : (i) i.r. spectra of but-l-ene adsorbed on unexchanged zeolite are similar when measured at room temperature150 ACTIVE CENTRES ON ZEOLITES and 120°C.5 (ii) TPD spectra do not change if, instead of nitrogen flushing, butene initially sorbed is outgassed by diffusion pump (see Experimental section). From the data given in fig. 1 and the composition of the sample it is easy to calcu- late that there are about 33 molecules of but-1-ene adsorbed for each unit cell of zeolite, i.e., 4 for each super cage or 0.6 for each sodium ion.It is plausible that adsorbed butene molecules interact first of all with easily accessible Na+ cations located in the super cages. The literature indicates that a majority of exchangeable Na+ ions are located inside the super cages and that their number amounts to 30-40 cations per unit ce11.*-12 These numbers are in adequate agreement with the number 33 found in our present case. The importance of exchangeable cations in the surface chemistry of zeolites has already been emphasised.5. 3* l4 We, therefore, advanced the hypothesis that Na+ cations are active sites during adsorption of but-l-ene. We consider now the isomerization step of the sequence (1).Isomerization of but-1-ene proceeds on the exchanged samples; up to 12 % exchange only isomer- ization, additional to adsorption, of but-1-ene takes place. Exchanged samples differ from unexchanged ones in containing Bronsted sites. It is therefore reasonable to assume that B and C are Bronsted centres. In order to distinguish between the two kinds of Bronsted sites we focus attention on the shape of the plots given in fig. 2, where a sudden change of the catalytic properties of zeolite is seen to occur at the end of interval I, above 12 % exchange, equivalent to 6.4Naf per unit cell. It seems that the initially introduced H+ ions are responsible for isomerization of butenes (B sites). We consider now the further catalytic reactions of butene which proceed on zeolites exchanged to an extent > 12 %.The amount of but-1-ene consumed increases almost proportionally with the degree of exchange in interval I1 from 12 to 68 %, after which it falls (fig. 2). Sodium cations exchanged within interval I1 amount to 29.8 per unit cell. This number is in good agreement with 32 accessible sites at position SII suggested originally by Breck,’ confirmed experimentally by Eulenberger et al. O and since accepted by many authors.12* 13’ 15-18 The intensity of the 3650 cm-l i.r. band of OH groups in zeolites increases only up to - 70 % exchange, as can be seen from experimental plots given in ref. (14) and (19). It follows that for higher degrees of exchange the groups do not increase in number. Since these Bronsted centres are easily accessible 2o and are catalytically active in various reactions,5* 12* 21 it is assumed that they are identical with centres C.Olson and Dempsey 22 reported a single-crystal X-ray diffraction study of hydrogen faujasite from which they conclude that the O1 oxygen atoms (the bridging oxygen atoms of the hexagonal prism) are the most favourable sites for formation of hydroxyl groups having a 3650 cm-l i.r. band, Other authors 17* 18* 2 3 agreed with this suggestion. Taking into consideration the fact that in each hexagonal prism there are two O1 oxygen atoms which could be associated with 24 there will be 32 sites per unit cell for hydroxyl groups here discussed. This number agrees well with our results.The 17.0 sodium cations not yet discussed belong to the interval above 68 % exchange. During the process of exchange of the last 20-25 Na+ cations by protons 12* 1 4 9 25 the i.r. band at 3550 cm-l appears. It is postulated that the protons responsible for this band lie inside hexagunal prisms and, therefore, are not easily accessible.12* 21* 22 Uytterhoeven et aZ.17 point out that the classification of ion-exchange sites given by Breck and X-ray data published by Eulenberger et aZ.1° agree on “ the location of some cations up to a maximum of 16 per unit cell in the hexagonal prisms ”. These sites will be discussed in more detail later.J. GALUSZKA, A. B A R A ~ S K I AND s. C ~ C K I E W I C Z 151 This point of view accords with other experimental data.g* 26 Thus a number of authors agree that there are 16 or so cations per unit cell having essentially different properties in that they are not easily acce~sible.~~ 11* l6 This inaccessibility and the decrease of the number of Na+ cations postulated as active sites for adsorption may explain the decrease of the relevant parts of the plots in fig. 2 especially since the number of centres C responsible for further catalytic processes attains maximum value at the end of interval 11, so their density is constant within interval 111.Another possible explanation of this trend, namely the poisoning of the sample by products of the further butene reactions, was excluded experimentally. It was found that repetition of a typical experiment on NaH-Y samples which have been used once (they turn grey) yields the same kinetic curve and the same TPD spectrum.The present discussion is based on the comparison of our kinetic data with in- formation available in the literature. It follows from the data that 53.2 cations in the unit cell can be divided into three groups, namely 6.4, 29.8 and 17.0 cations for the intervals I, I1 and I11 respectively. The assignment for groups of 29.8 and 17.0 cations from intervals I1 and I11 has already been discussed. We now return to the 6.4 cations from interval I which are related in some way to sites B. These cations are easy to exchange. They are close in number to the 8 occupied positions in the SIII positions as suggested by Breck.* Perhaps B sites have this origin. We turn now to the TPD spectra presented in fig.3. Special attention will be given to peak I representing but-1-ene for NaY and all three butenes for NaAH-Y samples. It is seen from fig. 3 that peak I decreases with increasing exchange. As a result of the decreasing amount of desorbed butene isomers the temperature of the maximum of peak I increases from 120-136°C as the exchange increases from 0 to 12 % (within interval I), which agrees well with the commonly observed shift in the TPD spectra of a heterogeneous s ~ r f a c e . ~ However, from 12-68 % exchange (interval 11) decreasing amounts of desorbed species cause a decrease in the tem- perature of maximum of peak I from 136to 77°C. Such a trend (opposite to that observed in interval I) may be explained by the assumption that the more strongly bonded of the adsorbed butene isomers are consumed by further catalytic reactions of sequence (1).It is seen from fig. 3 that the decrease of peak I is further enhanced at higher exchange. Near 80 % exchange the peak I is almost invisible. It seems that for a sufficiently high exchange it will vanish completely. If this is true, no adsorbed but-1-ene on Na+ centres will be present on the sample. It follows that the first step of sequence (1) will not occur. As a result, zeolite should be inactive for butene-1 transformations provided that the consecutive sequence is valid for these highly exchanged samples. All the main conclusions concerning the active centres are summarized briefly by the following enlargement of sequence (1) already given in the introduction.Nas+111? J OiH Na+ H + but-1 -ene(gas) --4 but- 1 -ene(ads) --+ isomerization -----+ polymerization 0iH dismutation ----+ further reactions dehydrogenation. We are grateful to the referees for helpful comments and for assistance with the phrasing of some sentences.152 ACTIVE CENTRES ON ZEOLITES A. BaraAski, S. Cgckiewicz and J. Galuszka, Bull. Acad. polon. Sci., S6r. Sci. chim., 1976, 24, 645. D. J. V. Smith, Adv. Chem. Series, 1970,101, 171. R. J. CvetanoviC and Y. Amenomiya, Adv. Catalysis, 1967, 17, 103. S. Cgckiewicz and J. Galuszka, Inz. Chem., 1975, 5, 643. A. Bieladski, J. Datka, A. Drelinkiewicz and A, Mdecka, Bull. Acad. polon. Sci., Sir. Sci. chim., 1976, 24, 137. J. Datka, Bull. Acad. polon. Sci., Sir. Sci. chim., 1974, 22, 975. ' P. D. Hopkins, J. Catalysis, 1968, 12, 325. D. W. Breck, J. Chem. Ed., 1964,41,678. R. Beaumont, D. Barthomeuf and Y. Trambouze, Adv. Chem. Series, 1971, 102,327. lo G. R. Eulenberger, D. P. Shoemaker and J. G. Keil, J. Phys. Chem., 1967,71,1812. l1 H. S. Sherry, J. Phys. Chem., 1966,70,1258. l2 J. W. Ward and R. C. Hansford, J. Catalysis, 1969, 13, 364. l3 J. A. Rabo and M. L. Poutsma, Adv. Chem. Series, 1971, 102, 284. l4 V. Bosa&ek, V. Patzelowa, C. Hybl and Z. Tvaruikova, J. Catalysis, 1975, 36, 371. l5 J. W. Ward, J. Colloid Interface Sci., 1968, 28, 269. l6 K. Tsutsumi and H. Takahashi, J. Phys. Chem., 1970,74,2710. l7 J. B. Uytterhoeven, P. .lacobs, K. Makay and R. Schoonheydt, J. Phys. Chenz., 1968,72, 1768. l 8 T. R. Hughes and H. M. White, J. Phys. Chem., 1967,71,2192. l9 A. Bieladski and J. Datka, Bull. Acad. polon. Sci., Sbr. Sci. chim., 1974, 22, 341. 2o L. Moscou, Adv. Chem. Series, 1971, 102, 337. 21 P. E. Eberly, Jr., J. Phys. Chem., 1967, 71, 1717. 22 D. H. Olson and E. Dempsey, J. Catalysis, 1969, 13,221. 23 R. A. Schoonheydt and J. B. Uytterhoeven, J. Catalysis, 1970, 19, 55. 24 E. Dempsey, J. Catalysis, 1975, 39, 155. 25 J. W. Ward, J. Phys. Chem., 1969,73,2086. 26 P. Gallezot and B. Imelik, J. Chim. phys., 1977, 68, 816. . (PAPER 6/2121)

ISSN:0300-9599    

DOI:10.1039/F19787400146

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

p-V-T Studies on Molten Alkali Nitrates Part 1 .-Thermal Pressure Coefficients and Compressibilities BY JOHN E. BANNARD*? AND ALLAN F. M. BAR TON^ Department of Chemistry, University of Southampton, Southampton SO9 5NH Received 20th December, 1976 The pV-T relationships of the molten alkali nitrates (Li, Na, K, Rb and Cs) were determined using a gas-pressurised, externally-heated pressure vessel. The cells were the pyknometer type with a long capillary barrier between the pressurising gas and the bulk of the liquid. Data, over a range of temperature from 500 to 800 K and a range of pressure to 1400 bar, were plotted in the form of isochores and used to determine densities and compressibilities of the liquids. The measurement of the conductance of melts at atmospheric pressure is a well established and precise procedure but its extension to conditions of high pressure is not.The case in favour of determining the density-dependence as well as the temperature-dependence of transport processes was argued previously and to obtain significant pressure coefficients requires a range of pressures up to and beyond 1 kilobar.' The analysis of the p and T dependence of conductance requires that the p-V-T relationship of the system also be known over the same range of conditions arid the determination of the isothermal compressibilities of melts becomes necessary. Furthermore, p-V-T data are necessary for an exhaustive study of the liquid state. The externally heated apparatus described earlier was used in the present work on the p-V-T relationships for the molten alkali nitrates.Few other p-V-T measurements at high static pressures have been made. Density measurements at 1 bar are straightforward involving either the Archimedean or the pyknometric methods, but their extension to high pressures involves considerable difficulties. Compressibilities of some alkali nitrates have been reported by Owens,2 a piston- cylinder method being used to 9 kbar and 520°C. Fray used the more elegant pyknometric method to study the nitrates, but, as with his conductance measurements? suffered large uncertainties associated with the solubility of the pressurising gas and with the influence of pressure on the dimensions of his cells. That work has been published together with the preliminary findings of the work reported here. The measurement of the velocity and attenuation of sound waves, however, is a more precise procedure even in molten salts.It gives rise to adiabatic compres- sibilities which can, via known values of Cp or Cv lead to the corresponding isothermal values, because Ps/BT = Cv/C,. Such applications are only feasible at present at 1 bar and both Bockris and Richards and Higgs and Litovitz have evaluated /IT in this way. Their results are in fair agreement and are close to those of Owens. The equilibrium, thermodynamic properties of any system can be described by the density [p = M/v], the thermal expansivity + Present address : Department of Metallurgy and Materials Science, University of Nottingham, NG7 2RD. Present address : Department of Chemistry, Murdoch University, Perth, Australia.153154 MOLTEN ALKALI NITRATES and the isothermal compressibility i av [ P T = -v(ap),l at all real values ofp, T and V. These last two coefficients, the so-called mechanical coefficients, are the experimental quantities normally determined in the studies of liquids; their significance has been discussed in detail by Ro~linson.~ The ratio of the mechanical coefficients is equal to the thermal pressure coefficient : The aim of the present work was to determine the p-V-T relations over a range of pressure and temperature out of contact with the pressurising fluid. The volumes of the salts were measured pyknometrically using a conductance bridge to detect the meniscus of the liquid. Hence constant-volume plots could be obtained at various values of p and T.The externally heated apparatus was used to study the nitrates over a range of temperatures to 52OoC, and pressure to 1400 bar. The practical data were obtained in the form of thermal pressure coefficients, and the data for the molten nitrates of Li, Na, K, Rb and Cs are given. These data are then used to derive densities and compressibilities which are discussed in terms of the possible structure of the melts. EXPERIMENTAL APPARATUS The externally heated pressure vessel, used up to a temperature of 520°C and a pressure of 1500 bar, has been described previous1y.l The vessel was pressurised with argon. The types of cells used are shown in fig. 1. They are all liquid-barrier pyknometers, the capillary presenting a long path to the diffusing gas.Fig. l(a) is a multiprobe cell, 25 cm long, which was used with the sodium and potassium nitrates. Fig. l(6) is a smaller version of this cell, w 14 cm long, with only two probes and was used for the expensive (4 (b) (4 FIG. 1.-The " constant-volume " cells : (a) and (b) Pyrex, (c) quartz.J . E . BANNARD A N D A . F. M. BARTON 155 rubidium and caesium nitrates. As lithium nitrate reacts with glass it was necessary to construct a cell from quartz, fig. l(c). Because of the difficulty in bringing metallic seals through quartz, the probes were introduced from the lip of the cell where they were rigidly held. The position of the meniscus of the melt was determined with the use of a conductance bridge circuit. In the case of the multiprobe cell, contact with a probe would be indicated by a change in the overall resistance; with the other two cells either short circuit or open circuit was identified.The bridge used was of the Wien type. MATERIALS Sodium nitrate and potassium nitrate (B.D.H., AnalaR grade) and rubidium nitrate and caesium nitrate (Johnson Matthey, 99.9 %) were dried at 180°C in an air oven for several days, maintained at just above their melting points for 10-12 h and filtered through sintered glass of porosity 3 before use. Lithium nitrate (Hopkin and Williams, Reagent grade) was recrystallised from distilled water, dried at 180°C for several days and again filtered before use, but this time in silica apparatus. PROCEDURE The alkali nitrates are fairly stable as melts in dry air, so the filling of the cell was made a simple process with the use of a small open furnace.Pure solid salt was placed in the top of the cell which was then lowered into the furnace. Subsequent to the salt melting the cell was alternately cooled and heated, and filling was brought about as a result of this heat-pump effect. The cell, connected to the pressure vessel cap, was then lowered into the hot vessel and removed after the plotting of each isochore in order that a quantity of salt could be added or taken away. Frequent cell failures occurred during this manipulation because freezing of the salt almost invariably led to breakage of the cell. The pressure was raised until the melt made contact with a probe. Time was allowed for thermal equilibrium to be achieved and then gas was allowed to leak out at a rate of one or two bars per min.When open circuit was again reached, the temperature and pressure were recorded. For the multiprobe cells, further rises in pressure allowed the use of another probe and another circuit resistance to identify the position of the meniscus. The duration of a pressure run was limited to -2 h and a similar time was allowed between runs at reduced pressure for the melt to degass. A correction to pressure ; which amounted ACCURACY was applied to the resistance of the platinum resistance thermometer due rT)T = -1.91 x bar-', to mO.7OC per kbar. The temperatures were measured to a tenth of a degree and as every care was taken to allow thermal equilibrium to be reached inside the vessel before the temperature or volume were recorded, the temperature is probably accurate to within k0.2"C.The pressures were recorded to 1 bar. The gauge was calibrated and the pressure will certainly be accurate to + 2 bar. These limits of imprecision are too small to be recorded on the isochores. A possible source of inaccuracy would be brought about by dissolution of the pressurising gas ; however thep-Tpoints always fell on straight lines except when gas was known to have penetrated the bulk of the melt. The effect of decomposition was also considered to be negligible as the experiments were generally carried out at temperatures below the recorded decomposition temperatures and because formation of nitrite was considered to have very little effect on the density and compressibility of the melt.After long periods (several days) at high temperatures, evidence of salt condensation was frequently found around the cell leads in the cooler part of the vessel. However the isochores were not plotted in any specific order and no hysteresis was found on jumpingI56 MOLTEN ALKALI NITRATES from low temperature to high temperature measurements and back again. The effect of salt evaporation was thus considered negligible. The measured slopes of the isochores were corrected for the expansivity and the compressibility of the cell material, Pyrex or quartz. The available values are very old and, as the correction accounted for w10 % of yv, it is possible that an unavoidable error has been introduced, say F 1 % of yv. The correction took the form P C d l + 1 1 %alt - = -(obs.) Yv Yv E s a ~ t + ace11 asalt + ace11 where o! is the volume expansivity of the melt or of the cell material./ ,,/ 0 2 4 6 8 Ib {2 tr pressure/bar x 1 O-’ FIG. 2.-Isochores for molten LiN03. TABLE 1 .-DENSITIES OF FUSED NITRATES a --bx 104 ref. LiN03 2.068 5.46 (lo), (11) NaN03 2.329 k0.017 7.14 (12), (13), (14) KN03 2.315 k0.032 7.29 (14), (10) RbN03 3.099 +0.050 10.10 (12), (14) CSNO~ 3.6206k0.065 11.66 (lo), (15) Calculation of compressibilities from isothermal pressure coefficients (and vice versa) and also from velocity of sound measurements relies greatly on accurate values of expansivity. This quantity is the temperature derivative of density and so any uncertainties in density will give rise to amplified uncertainties in compressibility. Similarly values of y y calculated from ref.( 5 ) and (6) will contain these same uncertainties which are probably _+3 % of a, (see table 1). A total uncertainty of k4.5 % of PT is therefore possible.J . E . BANNARD A N D A . F. M. BARTON 157 Compressibilities were here calculated from the thermal pressure coefficient and cc, from - b / p in the equation p = a+bT. The literature density data on molten salts vary considerably ; those considered the most accurate are given in table 1. RESULTS The raw data in the form of isochores are given in fig. 2-4 for the melts Li, Rb, and Cs nitrates respectively. Those for Na and K nitrates were presented ear lie^.^ Values of the slopes, extrapolated to zero pressure, were corrected for the compres- sibility and expansivity of the cell material as described above, and yv (corrected) was I I 1 I , 1 0 2 4 6 8 10 1 2 1 4 pressure/bar x lo-’ FIG.3.-Isochores for molten RbN03. platted against the temperature of the intercept as given for CsNO, in fig. 5. Also obtainable from the isochores are plots of yv against pressure for various temperatures, and one such family of lines is given for CsNO, in fig. 6. The scatter of the points .~~ ~~-~ _ _ _ _ ____.__ ___ take; from the yv againit T plot, fig. 5. Values of change of compressibility with temperature (fig. 7) may be obtained from plots like fig. 5, and a plot of compressibility against pressure (fig. 8) may be obtained from graphs of the type given in fig. 6. One-atmosphere compressibilities obtained from this work are compared with values obtained by other workers in table 2.158 MOLTEN ALKALI NITRATES The densities calculated at 1000 bar are compared with those of Owens2 in table 3.The atmospheric pressure values in the Gesent work were calculated from the equations given in table 1. Those in Owens' work were taken from ref. (16). The two sets of figures are in good agreement regarding the effect of pressure on density although in some cases different atmospheric pressure values were used. 700 800 12.01 ' - - * ' . - . ' " temperature/K FIG. 5.-yV against temperature for CsN03. Pressure, x , 1 bar ; 0, 1000 bar.salt LiN03 NaN03 KNOj RbNOj CsN03 J . E . BANNARD AND A , F . M. BARTON 0 2 4 6 a 10 pressure/bar x FIG. 6.-yy against pressure for CsN03. TABLE 2.-cOMPRESSIBILITY OF FUSED NITRATES tempPC 300 350 400 500 300 350 400 450 500 300 350 400 450 500 350 400 450 500 400 450 500 106jlT/bar-1 at zero pressure , DISCUSSION ref.(5) 19.6 23.4 28.9 17.8 21.6 26.8 19.3 23.4 29.4 - - - - - - I - - - - ref. (6) 20.3 23.3 27.4 18.1 21.3 26.3 19.3 23.5 28.9 - - - - - - - I - - - - ref. (2) - - 18.5 19.5 - - 21.5 24.5 - - - 20.5 30.5 - - 22.5 24.5 - - - - I59 this work 17.8 19.5 20.8 - 17.9 19.2 20.8 22.5 - - 20.0 22.5 25.7 - 22.4 24.7 27.3 - 28.6 31.7 34.8 The raw isochores (fig. 2-4) and the corrected isochores were all straight lines. The compressibilities of the fused alkali nitrates were determined principally because the thermodynamic analysis of the conductance data, especially the evaluation of the160 MOLTEN ALKALI NITRATES isochoric energies of activation, requires a knowledge of the p- V-T relationship for each of the ~a1ts.l~ Additionally, the data are of interest in their own right and.can be made the basis of an examination of the liquid-state properties of these systems. The corrected data for this study and for the calculation of compressibilities are in the form of values of yv, i.e.(ap/aT),, as a function of the absolute temperature, one of which is shown in fig. 5. These apparently linear plots are general for all liquid^.^ TABLE 3.-DENSITIES COMPARED WITH THE DATA OF OWENS LiN03 1 bar 1000 bar NaN03 lbar 1000 bar KNO, 1 bar 1000 bar RbN03 lbar 1000 bar Owens densities/g cm-3 400°C 500°C 1.700 1.646 1.735 1.682 1.853 1.783 1.892 1.826 1.824 1.753 1.861 1.801 2.395 2.298 2.448 2.352 this work densities/g cm-3 400°C 500°C 1.700 1.646 1.848 1.777 1.886 1.818 1.825 1.752 1.864 1.796 2.420 2.319 2.476 2.381 1.735 - That the relationship cannot be linear over the whole accessible temperature range is evident from the fact that an asymptotic rise in yv is to be expected as the solid or glassy state is approached. Also, as the temperature rises towards T,, yv rapidly approaches zero.The critical temperatures for these materials are not known, nor do we have data in the low temperature (glassy) region, but it may be concluded that the yv against T functions are virtually linear segments of an otherwise curved relationship. I 1 . 1 1 600 700 temperature/K FIG. 7.-Change of isothermal compressibility with temperature for the molten alkali nitrates.It is evident from fig. 7 that the compressibilities rise with temperature. This is to be expected as the gas-like character of the liquid develops and it also follows that with decreasing temperature, the compressibilities would decrease to the low values characteristic of the glassy and solid states. The compressibilities also change significantly with applied pressure (fig. 8) even over the relatively small range of pressures studied here. The values themselves are typically those of other liquidJ . E . BANNARD A N D A . F . M. BARTON 161 systems and it is possibly worth noting that the equilibrium, steady state and non- steady state properties of almost all liquids are very similar even though the melting points differ considerably.This is because the liquid state reflects short range interactions and the solid state long range interactions, Short range van der Waals interactions are similar for most molecular ensembles as are the repulsive interactions. Only if there is extensive structuring in the liquid state are the normal liquid properties significantly changed. * Whilst, therefore, the molten alkali nitrates have typical values of yv and PT, there are systematic differences between the values for the five salts (fig. 7 and 8) which correspond to the variation in cation radius. This does not seem to be a geometric factor because even in crystals there is a corresponding increase in PT. An increase in cation radius attenuates the residual coulombic attraction, and to a lesser degree the van der Waals attraction, without significantly affecting the repulsive interactions and a corresponding increase in compressibility is to be expected.The fact that NaNO, and LiN03 have similar compressibilities and that (fig. 7 and 8) I 1 20 - Li oo 10 pressure/bar x lo-" FIG. S.-Change of isothermal compressibility with pressure for the molten alkali nitrates. their T and p functions intersect arises simply because there is a lower limit of effective cation size at which anion-anion contact limits the further shrinkage of molar volume and at which anion-anion coulombic repulsion begins to attenuate the compres- sibility. The small size of the lithium ion will allow it to find refuge in a position bidentate to the nitrate ion thus hindering rotation and thus producing a large temperature derivative of the rotational degree of freedom.The similar steep rise in PT with temperature for CsNO, indicates a similar phenomenon, but this time low rotational order has probably been retained by virtue of large r++ repulsions rather than large r+- attractions. We thank the S.R.C. for financial support (to J. E. B.) and Prof. G. J. Hills for making laboratory facilities available. A. F. M. Barton, B. Cleaver and G. J. Hills, Trans. Faraday Sac., 1968, 64,208. B. B. Owens, J. Chem. Phys., 1966,44,3918. D. J. Fray, Ph.D. Thesis (London, 1965). A. F. M. Barton, G. J. Hills, D. J. Fray and J. W. Tomlinson, High Temp. H&h Press., 1970, 2,437-452. 1-61 62 MOLTEN ALKALI NITRATES J. O'M. Bockris and N. E. Richards, Proc. Roy. SOC. A, 1957,241,44. R. W. Higgs and T. A. Litovitz, J. Acoust. SOC. Amer., 1960, 32, 1108. C. Duval, Inorganic Thermogravimetric AnaZysis (Elsevier, 1963). G. W. Morey, The Properties of Glass (Rheinhold, 1938). ' J. S. Rowlinson, Liquids and Liquid Mixtures (Butterworths, London, 1959). lo G. J. Janz, Molten Salts Handbook (Academic, 1967). l 1 F. M. Jaeger and B. Kapma, 2. Anorg. Chem., 1920,113,27. l2 B. C. J. Neil, P h B . Thesis (Southampton, 1968). l3 B. de Nooijer, Ph.D. Thesis (Amsterdam, 1964). l4 W. J. McAuley, E. Rhodes and A. R. Ubbelohde, Proc. Roy. SOC. A , 1966,299, 151. l5 I. S . Yaffe and E. R. van Artsdalen, J. Phys. Chem., 1956,60,1125. l6 A. Klemm, in Molten Salt Chemistry, ed. M. Blander (Interscience, 1964). I'J. E. Bamard, A. F. M. Barton and G. J. Hills, High Temp. High Press., 1971, 3, 65-80. S. D. Hamann, Physico-chemical Efects ofPressure (Butterworths, 1957). (PAPER 6/2315)

ISSN:0300-9599    

DOI:10.1039/F19787400153

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

p- V-T Studies on Molten Alkali Nitrates Part 2.-Internal Energies and Equation of State BY JOHN E. BANNARD Department of Metallurgy and Materials Science, University of Nottingham, Nottingham NG7 2RD Received 20th December, 1976 p-V-T data for the molten alkali nitrates to a temperature of 800 K and a pressure of 1400 bar were used to derive internal pressures and specific heats, and to derive the parameters of the van der Waals equation. p-V-T measurements on liquids have been restricted to studies at relatively low temperatures.' However for a study of liquids of simple internal structure one must look to the molten salts, particularly to the molten alkali halides.2 The melting points of these salts are usually in excess of 750°C although much lower temperatures may be used in a study of the nitrates.Part 1 presented the p-Y-T data for the alkali nitrates in the form of thermal pressure coefficients, and used this data to derive densities and compressibilities. The isobaric expansivity, %[ = ;(g)J and the isothermal compressibility, are related to the isochoric thermal pressure coefficient, yv, by Furthermore these so-called mechanical coefficients may be related to a number of other properties, e.g. to the internal pressure, n, where and to the heat capacities, Cp and Cv because T Va; C,-C, = - BT The individual values of Cp and Cv can be resolved either by independent calori- metric means or in terms of the corresponding adiabatic compressibility, BS, which is defined by the relationship 163164 MOLTEN ALKALI NITRATES and which is normally determined from measurements of the velocity of sound in the liquid, W, i.e.( p is the density). Evidently, CPICV = PTlPS (3) and from experimental values of /IT and Ps together with eqn (2), values of C3 and Cv can be found. The acoustical method is precise and has been used in the evaluation of heat ~apacities.~ The determination of isothermal compressibilities is more difficult and probably less precise (especially at high temperatures and pressures). The Equation of State interrelates p-V-T data. There are a large number of empirical equations such as the Virial Equation of State, where, using modern computers, any number of virial coefficients may be used to cause the equation to fit experimental data. The approach taken here is to examine how closely the measured thermodynamic values fit equations of the van der Waals type. THE EQUATION OF STATE A comparison of the yv values obtained earlier can be made by plotting V,/y, against molar volume, V,.It is seen from fig. 1 that the mechanical coefficients of a large range of liquids are comparable. If these plots are considered to be straight, the slope is [l/y, x V,/(Vm- VO)], a type of free-volume function where Yo wiil be the intercept on the molar volume axis. This relationship is based on an application of the van der Waals equation of state, e.g. (P+n)(V-b) = BRT (4) where n is the internal pressure, b the excluded volume and B is a constant equal to 1 for a monatomic fluid and 2 for a polyatomic fluid. Recalling that the internal pressure may be written as (a,/P,)T-p, then eqn (4) may be modified to where the excluded volume term is considered as a definition of the free volume, VF = V-b.According to this simple relation yv should be a linear function of V and at l/y, = 0, V = b. That this is very nearly so is evident from fig. 1 which illustrates the relation for the five alkali nitrates. The precision of the results is not sufficient to establish the linearity of the relationship; indeed for all the liquids there seems to be a slight systematic curvature. The dashed-arrows indicate the values of solid volume, V,, for the five nitrates. (Virtually no measurements have been made of cc, or /IT for ionic solids, but yv for solid NaNO, has been estimated at = 100 bar K-I). The slopes of the lines show some scatter but a reasonably constant B value of 2-3 indicates that the p-V-T properties of these melts are approximately expressed by a van der Waals equation for a polyatomic f l ~ i d .~ Some exception would be expected for RbNO, because it is known that the solid undergoes a number of phase transitions to low density structures. As a consequence rubidium nitrate belongs to that rare group of solids which contract on fusion. But it is seen that none of the values of AVfUsi,, correspond to the value V - b, so that the relevance of the V, values in fig. 1 should not be overemphasised. The value of the intercept, b, together with the molar volume of the solid at 25°C and the molar volume of the liquid at the melting point are summarised in table 1.J .E . BANNARD 165 The approximate nature of the van der Waals equation can be illustrated by evaluating b from a rearrangement of eqn (3, i.e. and b values for the nitrates may be deduced from table 2. Their average values can be compared with the volume change on fusion tabulated in table 1, and the differences illustrate the difficulty of unambiguously defining Fg. b = V-BR/yy, Vm/cm3 FIG. l.-Vm/yv against molar volume for various liquids. TABLE 1 .-FREE-VOLUME FOR THE ALKALI NITRATES m.pt. Vo = b V V S AVtus /K /cm3 mol-1 /cmJ mol-1 /cmJ mol-1 /cm3 mol-1 B LiN03 525 29 38.7 29.0 6.84 2.1 NaN0, 580 32 44.6 37.6 4.32 3.1 KNO, 611 45 54.0 48 .O 1.73 2.2 cSNO3 679 53 69.2 52.9 7.48 2.9 RbN03 581 45 59.3 47.5 -0.14 3.2 TABLE 2.-PARAMETERS OF THE VAN DER WAALS EQUATION T/K LiN03 600 700 NaNO, 600 700 KN03 600 700 RbN03 600 700 CsN03 600 700 YV V Y-b /bar K-1 /cm3 mol-1 /cmJ 16.57 39.6 10.6 20.01 44.8 12.8 17.87 46.5 14.5 20.20 53.9 8.9 16.58 56.0 11.0 17.97 59.2 14.2 15.42 61.7 16.7 13.58 69.4 16.4 - - - - - - (Rlyv) /cm3 mol-1 4.96 4.1 1 4.59 4.10 4.99 4.61 5.36 6.10 - - (BR/rv)- /cm3 mol 1 10.4 12.7 14.2 9.0 11.0 14.7 17.1 - p+n/bar 9940 11990 12440 12130 11590 10780 10820 - - - 17.7 9500 p+n/RT 0.183 0.244 0.217 0.247 0.202 0.219 0.189 0.165 - -166 MOLTEN ALKALI NITRATES The other term of the van der Waals equation, the internal pressure term, can be calculated directly from the observed values of yv and p , i.e.from the relation n = T(dp/aT)v-p. The various parameters of the van der Waals equation are compared in table 2.SPECIFIC HEATS The mechanical coefficients are also related to the specific heats. Thus, the difference in the isobaric and isochoric heat capacities is related to the expansivity and compressibility by the thermodynamic relation given in eqn (2). The determina- tion of individual values of Cp and C, requires either that one be known from other sources or that another relation be used ; i.e. eqn (3), where ps is the adiabatic com- pressibility which is normally found from measurements of the velocity of sound in the system. Adiabatic compressibilities have been derived for Li, Na and K nitrates from the measurement of sonic velocities in these melts. The heat capacities obtained 1 bar 1000 bar 1 bar 1000 bar 1 bar 1000 bar TABLE 3 .-TEMPERATURE AND PRESSURE DEPENDENCE OF SPECIFIC HEATS temp./K /I(-1 /cm2 dyne-1 /cm2 dyne-1 Cp/Cv /J mol-1 K-1 /J mol-1 K-1 /J mol-1 K-1 104ap 1 0 1 2 ~ ~ 1012BP BTlPS = c p - c v CP c v LiN03 573 3.111 17.7(18.3) 18.3 1.034 11.88 360 348 673 3.210 21.0(20.9) 22.0 1.052 12.76 259 246 573 3.059 16.2 16.6 1.026 12.47 493 481 673 3.149 18.8 19.6 1.043 13.56 332 318 NaN03 573 3.719 15.6(15.6) 17.9 1.147 19.58 154 134 673 3.864 18.7(18.4) 21.3 1.139 21.67 177 155 773 4.018 22.6(22.1) 25.2 1.115 23.68 229 205 573 3.660 14.4 17.2 1.194 19.50 120 100 673 3.790 17.0 19.8 1.165 22.09 156 134 773 3.927 19.5 22.8 1.169 24.52 171 146 KN03 573 3.843 16.1(16.1) 18.3 1.143 24.69 197 172 673 3.997 19.3(19.1) 23.0 1.195 25.90 1 60 134 773 4.163 23.3(22.8) 30.4 1.300 25.48 109 84 573 3.781 14.8 16.8 1.136 25.52 214 188 673 3.915 17.6 20.4 1.159 27.36 199 172 773 4.057 20.4 25.4 1.245 28.24 141 113 by means of eqn (6) and (3) are compared at different temperatures and pressures in table 3.The values of BS in brackets are those derived from the ultrasonic data of Higgs and Litovitz;' they differ slightly from the other values taken from similar work by Bockris and Richards.* For lithium nitrate and for potassium nitrate the ratio pT/Ps rises with rise in temperature and falls with rise in pressure. However for sodium nitrate the ratio falls with rise in temperature and rises with rise in pressure. Subsequently the trend in specific heats is reversed in the case of NaNO,. The actual values obtained for the specific heats are of the magnitude expected for highly associated liquids, cf.the values quoted by Rowlinson for liquids of various kinds and those for the oxyanion melts quoted by L~msden.~ In general Cp and Cv converge as the triple point is approached, i.e. as ( p + z) AV becomes zero,J. E . BANNARD 167 and at higher temperatures Cp usually rises but sometimes only after a low temperature fall. However the apparent variations of Cp and Cv with temperature as given in table 3, may or may not be significant because the precision with which the absolute values of these quantities is known is really very low. The fact that they are large and similar in magnitude makes their evaluation from their ratio and difference a very uncertain procedure and the data derived here (and elsewhere) from com- pressibility measurements must be regarded with reserve.The two groups of ultrasonic workers ' 9 * disagree in their adiabatic compressibilities by up to 4 %, so an uncertainty in ps of +4 % is immediately introduced. This together with an uncertainty of up to +4.5 % in PT produces the large errors. Once, of course, a I 1 6h ' *6 40 390 39.5 Vm/cm3 temperatures in K. FIG. 2.-Change of internal energy with molar volume for LiN03. 0, 1 bar ; x , lo00 bar; single calorimetric value of Cp or Cv is known, then the other data could be evaluated with confidence. Assuming, however, that both the magnitude and the sense of the specific heats given in table 3 do have some significance, differences between the structures of the three melts are suggested. The minimum in Cp for water between the melting point and the boiling point has been explained in terms of the large change in the configurational contribution to the specific heat,lo which itself is a measure of the distortion and breaking of the cohesive bonds in the liquid. The maximum contribution to the specific heat from the " internal " energies, i.e.from rotational, vibrational and translational modes, can be roughly estimated. If the melt consists of fully-liberated, non-planar M-N03 species then the maximum contribution from these sources is 12R, i.e. approximately 100 J mol-1 K-l. The difference between168 MOLTEN ALKALI NITRATES this quantity and the values given in table 3 must be accounted for by the configura- tional specific heat, the specific heat arising from dynamic changes in structure. The values for LiN03 are considerably above 100 J mol-1 K-l at the temperatures calculated, although showing a very large temperature-dependence and a very large pressure dependence.This supports the anticipated highly-structured configuration for that melt. INTERNAL ENERGY The internal pressure of the liquid, n, may be obtained from the thermal pressure coefficients according to eqn (I), hence plots of yv T-p may be made as a function of molar volume and these are given in fig. 2 to 6. Values are also given for 1000 bar, and although it is immediately apparent that this range of pressure produces very 45 46 Vm/cm3 temperatures in K. FIG. 3.-Change of internal energy with molar volume for NaN03. 0, 1 bar; x , 10o0 bar; little change in the derivative (aE/dY), for the alkali nitrates, the slopes of the isotherms between the 1 and 1000 bar points may be estimated and are given in table 4.The 6: 12 potential of Lennard-Jones and Devonshire for simple liquids leads to the well known energy/distance relationship ( E / r ) , and for any simple material, r will have a direct dependence on molar volume, V,, and the derivative i3E/i3Vm will take the form shown in fig. 7. It will be noted that the experimental plots of (yv T-p) given in fig. 2 to 6 take the form of the curve at positions A, B and C depending on the cation. There is a clear tendency for these isobaric plots to go through maxima. The isothermal slopes, between two pressure points, are given in table 4. Gibson et aZ.11-14 made a number of similar plots for a series of liquids ranging from water and ethylene-glycol to hydrocarbons, and obtained plotsJ .E. BANNARD 1 1 1 169 11.2 54 55 56 Vm/cm3 temperatures in K. FIG. 4.4hange of internal energy with molar volume for KNO+ 0, 1 bar; x , 10oO bar; t 0.m Vmlcm3 temperatures in K. Fro. 5.-Change of internal energy with molar volume for RbN03. 0, 1 bar; X , lo00 bar;170 MOLTEN ALKALI NITRATES of similar shape with, in some cases, maxima. Again, the simple model of a liquid describes the internal energy as a sum of the attractive potential energy and the repulsive : and E = EA-ER In the studied region - (a2ER/aV2)T increases more rapidly than (a2EA/a V2)= as the volume is diminished until at the maximum (dE,/aV), = -(aE,/aV,). Further decrease in volume increases the repulsive term causing (aE/aV), to fall.Of particular interest is the effect of temperature on the internal pressure, n. For most liquids n falls with rise in temperature which might be considered as a criterion of I I I 1 740 x - I I I 78ov\, 69 70 71 Vm/cm3 FIG. 6.-Change of internal energy with molar volume for CsN03. 0, 1 bar; x , lo00 bar; temperatures in K. " normality ".15 Two notable exceptions were the plots for water and ethylene glycol,14 where the n values rose with rise in temperature over the experimental range; there appears to be a tendency for this to be the case for the molten alkali nitrates although the experimental scatter is considerable. Gibson suggested that this is due to directional interaction, very strong in the case of water.For the effect of cation on the properties of a series of liquids like the alkali nitrates to be studied, acceptable corresponding states should be compared. However in broad terms the plots show that the internal pressure decreases markedly with increase in density (having passed through a maximum) which follows from the corresponding increase in the repulsive component of the intermolecular potential. Support for this is found in the data for LiN03 for which the total internal pressures are comparatively low and in which we suppose the repulsive contributions are the greatest. The internal pressure data for the salts just above the melting point isJ . E . BANNARD 171 given in table 5. A strict comparison on the basis of corresponding states is not easily achieved and here the simple expedient is followed of comparing the total internal pressure for the five salts at " corresponding " free volumes, i.e.at values of V, 25 % in excess of V,. As expected the increase in coulombic attraction with \ vm FIG. 7.-The Lennard-Jones potential for a liquid. decreasing cation radius systematically enhances the value of n from Csf to Na+, but thereafter coulombic repulsion becomes significant to the degree that n for lithium nitrate is reduced considerably. reveals that the alkali nitrates with the exception of lithium nitrate all have similar structures with the force constant K shown in An assessment of spectral data TABLE 4.-sLOPES OF ISOTHERMS (0-IOOO bar) temperature/K slopc/kbar cm-3 temperature/K slope/kbar cm-3 temperature/K slope/kbar cm-3 LiN03 NaN03 KN03 580 +0.30 660 +0.31 660 -0.18 600 + 0.20 680 + 0.30 680 - 0.26 620 +0.10 700 + 0.22 700 -0.33 640 + 0.08 720 +0.16 720 -0.38 660 0 740 +0.14 740 - 0.48 RbN03 CsN03 660 + 0.06 740 -0.11 680 - 0.02 760 -0.13 700 - 0.03 780 -0.17 720 - 0.08 740 - 0.08 760 -0.12172 MOLTEN ALKALI NITRATES fig.S(a) decreasing with the surface charge density of M+. The magnitude of this force constant implies a very definite association between the metal ion and the nitrate ion. Interactions of this type have been further suggested from heats of mixing of ternary halide + nitrate mixtures,l the excess enthalpies undoubtedly arising from departures from random mixing, and from n.m.r. studies.l* The force constant for Li was found to be lower than that for Na, in fact about the same as the K-N03 value.Preferential penetration of the small Lif into the bidentate position may well account for this, fig. 8(b). TABLE 5.-~”I’ERNAL PRESSURES FOR THE NITRATE MELTS density at m.pt. Vb m 125 %YO =i MIV; 1g cm-3 ;r/bar Volcm3 / a 3 / g cm-3 n at pb/bar LiN0, 1.781 9680 32.5 40.6 1.700 9940 NaNO, 1.91 5 11850 36 45 1.889 12100 KN03 1.872 12110 45 56.3 1.796 1 1490 RbN03 2.510 10740 49 61.3 2.406 10820 CSNO, 2.820 9530 57 71.3 2.734 9270 Ion association of this type seems to be important in nitrate melts and an increase in such association with increase in pressure, followed by loss of rotational freedom and increase in the amount of bidentate structure is probably the reason for the enhanced repulsive term.This also helps to explain why the entropy change for the fusion of lithium nitrate is so much higher than the corresponding chloride salt (see table 6). The low value of ASfusion for RbN03 indicates that the anion is:possibly freely rotating in the solid below the melting point. FIG. &-The structure of an M+ nitrate.J . E. BANNARD 173 TABLE 6.-THERMODYNAMIC FUNCTIONS RELATING TO THE FUSION OF ALKALI CHLORIDES AND NITRATES l9 LiCl NaCl KCl RbCl CSCl LiN03 NaN03 KN03 RbN03 CSNO~ m.pt./K 883 1073 1043 995 91 8 525 580 61 1 581 679 AHruaion /Id mol-1 19.92 27.99 26.53 23.72 18.74 25.61 14.73 11.72 4.64 14.10 A Vfwion A S ~ o l l / % /cm3 mol-1 /J mol-1 K-1 +26.2 5.8 22.55 f25.0 7.5 26.07 +17.3 7.3 25.44 +14.3 6.8 23.85 +10.0 5.6 22.05 +21.4 6.84 48.79 +10.7 4.32 25.52 +3.32 1.73 19.16 -0.23 -0.14 7.99 + 12.1 7.48 20.75 CONCLUSIONS The non-hard-sphere nature of the nitrate ion adds an element of complication to the interpretation of the mechanical coefficient data.In this respect a study of the technically more difficult halide melts would be useful. The significant results of this work on nitrate melts lie in the computed internal energy changes with density of the liquids, and the relative effects of different cations. However, for this analysis to be carried further, studies must be made over much greater ranges of T and p than those reported here. The author thanks the S.R.C. for financial support and Prof. G. J. Hills for making laboratory facilities available. J. S. Rowlinson, Liquids and Liquid Mixtures (Butterworth, London, 1959). J. E. Bannard and A. F. M. Barton, J.C.S. Faraday I, 1978,74,153. K. F. Herzfeld and T. A. Litovitz, Adsorption and Dispersion of Sonic Waues (Academic Press, New York, 1959). R. N. Haward, Trans. Farahy SOC., 1966,62,828. B. Cleaver and J. F. Williams, J. Phys. Chem. Solids, 1968,29, 877. ’ R. W. Higgs and T. A. Litovitz, J. Acoust. SOC. Amer., 1960,32, 1108. a J. O’M. Bockris and N. E. Richards, Proc. Roy. SOC. A, 1957,241,44. J. Lumsden, Thermodynamics of Molten Salt Mixtures (Academic Press, 1966). ‘ G. Goldmann and K. Todheide, 2. Naturforsch., 1976,31a, 656. l o D. Eisenberg and W. Kauzmann, The Structure and Properties of Water (Oxford, 1969). l 1 R. E. Gibson and J. F. Kincaid, J. Amer. C/zem. SOC., 1938,60,511. l2 R. E. Gibson and 0. H. Loeffler, J. Phys. Chem., 1939,43,207. l3 R. E. Gibson and 0. H. Loeffler, J. Amer. Clzem. SOC., 1939,61,2515. l4 R. E. Gibson and 0. H. Loeffler, J. Amer. Chem. SOC., 1941, 63, 898. l 5 J. H. Hildebrand, J. Chem. Phys., 1939,7, 233. l6 S. C. Wait, A. T. Ward and G. J. Janz, J. Chem. Phys., 1966,45, 133. l9 G. J. Janz, Molten Salts Handbook (Academic Press, New York, 1967). S. V. Messchel and 0. J. Kleppa, J. Chem. Phys., 1965,43,4160. S . Hafner and N. H. Nachtrieb, J. Chem. Phys., 1964, 40,2891 ; 1965, 42, 631. (PAPER 6/2316)

ISSN:0300-9599    

DOI:10.1039/F19787400163

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Hydrogenolysis of Cyclopent ane and Hydrogenation of Benzene on Palladium Catalysts of Widely Varying Dispersion BY SERGIO FUENTES AND FRAN YOIS FIGUERAS* Institut de Recherches sur la Catalyse du C.N.R.S., 79, boulevard du 11 Novembre 1918, 69626 Villeurbanne Cedex, France Received 26th January, 1977 Adsorption of hydrogen and oxygen, titration of preadsorbed oxygen, benzene hydrogenation and cyclopentane hydrogenolysis were performed on palladium catalysts of widely varying dispersion. A change in the stoichiometry of oxygen adsorption is found when increasing the metallic dispersion. On clean supports, free of sulphur and iron, the turn over number for both reactions of hydrocarbons is constant and independant of, the dispersion of Pd. After a suitable reduction, sulphur may preferentially inhibit hydrogenolysis, while iron preferentially inhibits hydrogenation.On an industrial silica support the turn over for hydrogenolysis changes with dispersion ; this effect is attributed to contamination of Pd by iron from the support. Much recent work on catalysis by metals has been concerned with the question of how the degree of dispersion, or crystallite size, of a metal influences the specific catalytic activity. An effect of the dispersion has been found on hydrogenolysis of ethane on nickel and Rh/Si02,2 on hydrogenolysis of neopentane on platinum and of cyclohexane on Ru/S~O,.~ Maurel et aL5 recently demonstrated that trace impurities from the support could noticeably alter the selectivity of platinum, and the notion of selective poisoning was thus introduced.Only a few metals have been studied from this point of view, and the data are rather scarce, in particular for palladium. The present work is devoted to the study of the influence of metallic dispersion on the catalytic activity of palladium. We compared the rates for benzene hydrogenation, which is known to be structure insensitive, and for ring opening of cyclopentane which could be structure sensitive on platinum and which can be considered as a parent reaction of the hydrogenolysis of cyclohexane used recently by Lam and S i ~ ~ f e l t . ~ EXPERIMENTAL The catalysts were prepared by ion exchange using Pd(NH3)2C12 in basic solution with silica, and PdC12 in acidic medium with alumina. The solids were left overnight in contact with the solution, then filtered and dried at 110°C.The supports were : Davison silica gel (grade 70), of surface area 350 m2 g-I, containing traces of iron (0.02 % Fe) and sulphur (0.17 %) in the form of sulphate SO, ; Degussa alumina (llOC), a non porous alumina of surface area 180 m2 8-l containing only small traces of iron (< 200 p.p.m.) ; and y-alumina made in the laboratory with care taken to avoid introducing impurities, containing only weak traces of iron (< 100 p.p.m.) undetectable sulphur. The study of the possible influence of sulphur and iron impurities was undertaken on samples of y-alumina impregnated by known amounts of contaminants in the form of (NH4)2S04 and FeCl,. The chemical compositions of these solids are listed in table 1. A good dispersion (90-100 %) was obtained after calcination at 400°C in oxygen and reduction in dry hydrogen at 300°C.Changing the temperatures of calcination and reduction resulted in gradual sintering of the metallic phase, probably due to variation in the water content of the sample as underlined by Boudart.' 1 74S. FUENTES AND F. FIGUERAS 175 The dispersion of palladium was measured by H2-02 titration, using volumetry at 0.1 Torr and a temperature of 70°C for H2 adsorption, to avoid H2 dissolution, as in a similar procedure used by other authors.8* The dispersion was defined in the usual way by Pd,/Pd, Pd, being the iiumber of superficial palladium atoms which adsorb hydrogen and oxygen and Pd the total number of palladium atoms of the sample. Catalytic activities were determined on an aliquot of the sample used for dispersion measurements.The sample was reactivated for 1 h at 300°C under hydrogen. The rates were measured in a conventional flow reactor at low conversion (<2 %) to avoid heat and mass transfer limitations. For benzene the conditions were : temperature 14O"C, partial pressure of hydrocarbon 56 Torr, partial pressure of H2 704 Torr ; under these conditions the reaction order relative to benzene is zero. TABLE 1 .-CHEMICAL COMPOSITION OF THE SAMPLES CONTAMINATED WITH SULPHUR AND IRON catalyst % Pd wt % S wt % Fe - 200 1 0.1 310 0.35 0.1 - 320 0.35 0.2 - 330 0.25 0.4 - 340 1 I 0.035 350 1 I 0.056 360 1 - 0.078 For cyclopentane the reaction temperature was 290"C, pressure of hydrocarbon 100 Torr and pressure of hydrogen 660 Torr.The reaction order relative to cyclopentane was found to be close to zero. The sole product of hydrogenolysis was n-pentane on Pd/alumina. On Pd/Si02 n-pentane was the major product (selectivity 75-80 % initially) but small amounts of methane,butane and isopentane were detected during a short initial period ; they decreased rapidly with time. The catalysts suffer deactivation in the reaction of conversion of cyclopentane. A better correlation between activity and metallic dispersion may be expected when using the initial activity. Therefore an effort was undertaken to determine the law of deactivation. The hyperbolic law first proposed by Germain and Maurel lo fits the data well as illustrated in fig. 1. The equation is a particular form of a more general rate law proposed by Szepe and Levenspiel.ll Levenspiel l2 recently discussed the means of testing several models of deactivation by a proper choice of the reactor and pointed out the usefulness of the well-mixed reactor for that purpose.A flow reactor operated at low conversion can be assimilated to a well mixed reactor. If we assume that the reaction proceeds by a parallel scheme : n-pentane I' 4 cyclopent ane coke and that resistance to diffusion is small, the theory developed in ref. (12) predicts for a second order of deactivation, the following law : where C A ~ is the inlet concentration of reactant, CA is the outlet concentration of reactant, k is the zero order rate constant of the reaction, ki is the deactivation constant, with k'd = k d ( C ~ ) ~ and z' = WCA,/FA, with W the mass of catalyst and FA^ the flow rate of reactant.176 HYDROGENOLYSIS OF CYCLOPENTANE timejmin 0, Pd+ S/Al203-334 ; A, Pd/Si02-47 ; + , Pd/A1203-21 1.FIG, 1 .-CataIyst deactivation in cycIopentane hydrogenolysis. Plots of the reciprocal conversion as a function of time for 3 samples. Therefore the plot of the reciprocal conversion CAJCA~- CA against time gives a straight line which allows the determination of the initial activity and of the rate constant of deactivation. A detailed discussion of the data gathered on deactivation will be given elsewhere, but we can note here that, a good representation of the kinetics can be obtained by this type of analysis after making some reasonable assumptions. RESULTS AND DISCUSSION DETERMINATION OF THE PALLADIUM AREA Benson et al.have extended to palladium the technique of titration of preadsorbed oxygen previously applied to ~1atinum.l~ On a sample characterized by a moderate dispersion of palladium ( D N 0.2), the stoichiometry observed for the ratios : hydrogen adsorption/oxygen adsorption/titration of preadsorbed oxygen was 1-1 -3. The reactions which represent the processes of adsorption are : Pd++H2 -+ PdH Pd++02 + PdO PdO + +H2 --+ PdH + H20. The results of the present work, summarized in table 2, are in agreement with this stoichiometry when the dispersion is low ( D < 0.4). However at higher disper- sion ( D fi 0.9), the stoichiometry shifts to 1-0.5-2 as illustrated in table 2 by the samples 176 and 211. In the intermediate range (sample 1702) the stoichiometry is not a simple one.This situation is quite similar to that described by Wilson and or Dalla Betta and Boudart for well dispersed platinum catalysts. The interpretation now accepted in the case of platinum is a change in the stoichiometry of adsorption of oxygen which give PdO on large particles and Pd20 on small particles. The reaction of titration of oxygen on small particles is then : Pd2O + 2H2 --+ 2PdH + HZO.S . FUENTES AND F . FIGUERAS 177 The shift from PtO to PtzO was attributed in ref. (7) to the fact that small metallic particles are electron deficient and cannot therefore supply enough electrons to complete a monolayer of oxygen. We can suppose, since the behaviour of palladium is the same as that of platinum as a function of the dispersion, that the assumptions made previously for platinum are valid for palladium.The change in stoichiornetry of oxygen adsorption with particle size is thus a general phenomenon and not restricted to platinum. TABLE 2.-sTOICHIOMETRIC RATIO FOR THE ADSORPTION OF HYDROGEN, ADSORPTION OF OXYGEN AND HYDROGEN TITRATION OF PREADSORBED OXYGEN ON DIFFERENT SAMPLES OF SUPPORTED PALLADIUM H/Pd O/Pd H/Pd Pd/A1203 176 1 .o 0.48 1.85 Pd/A1203 21 1-1 1.03 0.59 2.13 Pd/A1203 1702 0.73 0.51 1.62 Pd/Al2Oj 185 0.31 0.28 0.82 catalyst adsorption adsorption titration Pd/A1203 211-2 0.96 0.49 1.99 Pd/Si02 1 0.35 0.34 1 .oo Pd/Si02 66 0.47 0.39 1.22 CATALYTIC ACTIVITIES OF CLEAN SAMPLES The catalytic activities of some samples, supported on y-alumina or silica are reported on table 3 in terms of turn over number N (number of molecules reacting per hour and per surface palladium atom).On a clean support like y-alumina, the turn over numbers N for both reactions are constant and independent of the dispersion. The values for benzene hydrogenation are in good agreement with previous determinations on palladium catalysts. 9 l6 This behaviour is character- istic of structure insensitive reactions, and of the absence of any support effect. In contrast, Davison silica supported palladium gives a low but constant turn over for benzene hydrogenation and a striking increase in the turn over for hydrogenolysis with dispersion. The results obtained with Pd/Si02, if they were taken separately, could lead to the conclusion that the hydrogenolysis of cyclopentane is structure TABLE 3.-TURN OVER NUMBER FOR THE HYDROGENATION OF BENZENE AND HYDROGENOLYSIS OF CYCLOPENTANE AS A FUNCTION OF THE DEGREE OF DISPERSION, ON TWO SUPPORTS N2 cyclopentane at 290°C N1 benzene at 140°C samples wt %Pd Tnd/OC % D / molecules h- 1 / molecules h-1 R = N I I N ~ y-aluminas 2202 0.25 21 1 0.5 1702 1 .o 176 1 .o 184 1 .o 185 1 .o 122 1.2 252 8.0 83 0.7 41 1.3 4 1.3 2 1.3 72 1.3 400 300 400 300 400 500 300 500 650 500 400 300 400 100 305 100 286 55 243 80 270 40 264 28 220 12 282 12 230 Davison silica gel 4 115 26 104 33 101 48 94 32 140 33 24 30 24 28 21 30 22 8 15 19 34 50 9.2 12 9.7 11.2 9.4 10.5 9.4 10.5 14.4 6.9 5.3 2.5 2.5178 HYDROGENOLYSIS OF CYCLOPENTANE sensitive.Indeed, Lam and Sinfelt interpreted their results obtained on the hydrogenolysis of cyclohexane on Ru/SiO, in this way.However, such an effect should appear on alumina as well as on silica and this is not the case. Therefore we cannot conclude that there is a particle size effect on the catalytic properties of palladium, but rather may suspect possible influence of an impurity. Sulphur and iron being common impurities of silica, we investigated the influence of these contaminants. INFLUENCE OF CONTAMINANTS The relative ratio poison/palladium was varied by changing the chemical composi- tion (table 1) and the temperature of reduction. The contaminant content was kept low in order to simulate a catalyst supported by a commercial carrier. The palladium area of contaminated samples is difficult to determine by H2-02 titration owing to the presence of impurities, the behaviour of which is unknown in the process of adsorption.In this case a useful procedure is that proposed by Maurel which consists in the comparison of the rates of hydrogenation and of hydrogenolysis on the same sample. The hydrogenation of benzene is known to be insensitive to structure and the activity for this reaction can be a measure of the palladium area. If the hydrogenolysis of cyclopentane proceeds on the same sites as hydrogenation, the selectivity ratio, rate of hydrogenation/rate of hydrogenolysis must be constant ; a variation of this selectivity ratio must reflect that hydrogenolysis and hydrogenation proceed on different sites, therefore that hydrogenolysis is structure sensitive.INFLUENCE OF SULPHATES The results are summarized in table 4. The blank runs on clean samples give the activities in the absence of sulphates. The temperature of reduction has a strong influence on activity and selectivity, TABLE 4.-cATALYTIC ACTIVITIES ( X lo5 mol S-l 8-l Pd) OF Pd/Al203 SAMPLES, CONTA- MINATED WITH SULPHATE AND REDUCED AT 300 OR 400°C catalyst 176 2201 1702 201 204 311 321 331 202 205 312 332 loading activity A 1 for temperature of cyclopentane %Pd % S reductionl'c at 290°C 1 0 300 (2 h) 5.3 0.25 0 300 (2 h) 9.3 1 0 400 (2 h) 4.2 1 0.1 300 (1 h) 2.9 1 0.1 300 (1 h) 2.9 0.35 0.1 300 (1 h) 6.7 0.3 0.2 300 (1 h) 8.0 0.25 0.4 300 (1 h) 4.32 1 0.1 400 (12h) 0.85 1 0.1 400 (2 h) 1.08 0.35 0.1 400 (12h) 0.52 0.25 0.4 400 (2 h) 0.08 activity A2 for benzene at 140°C S = A21Ai 62 11.7 100 10.7 50 11.9 average 11.4 45 16 35 12 90 13 87 11 60 14 average 13.3 20 24 22 20 13 25 2.5 31 average 25 -S .FUENTES AND F. FIGUERAS 179 The samples contaminated with sulphate and reduced at 300°C have activities close to the blank, the maximum decrease is by a factor of 2. The selectivity of these catalysts is slightly higher than that of clean samples, but the difference is not large. In contrast, after reduction at 400°C, the activities for hydrogenation and hydrogenolysis decrease sharply, but the selectivity for hydrogenation increases : this increase means that sulphur preferentially represses hydrogenolysis. These results obtained on palladium are similar to those published for platinum : 5 on both metals a selective poisoning of hydrogenolysis can be observed in the presence of sulphate after reduction at 400°C, while reduction at 300°C yields a non-selective poisoning.It is gratifying that palladium has the same behaviour as platinum in this respect, and the same interpretations may be valid in both cases. Maurel et aL5 concluded that hydrogenolysis proceeds on metal sites with specific geometrical requirements as suggested by Boudart,l while hydrogenation would proceed on all superficial metallic atoms. Elementary sulphur produced by the reduction of sulphates at 400°C would be preferentially adsorbed on these sites of low coordination and could specifically poison hydrogenolysis. The reduction at 300°C would yield H2S or SO2 which are non-selective poisons of platinum.The influence of sulphate poisoning on palladium is thus in good agreement with literature data concerning platinum. However it seems clear that the results obtained in the present work with Pd/Si02 cannot be attributed to sulphur poisoning since the selectivity ratio is shifted in the opposite way. INFLUENCE OF IRON The experimental data are listed in table 5. In the case of iron the temperature of reduction of the sample also influences the results. Moreover, an effect of the iron content can be noted. Iron contamination has only a small effect on the catalytic properties when the sample is reduced at 300°C. The selectivity ratio is slightly lowered but the difference is not large. This is consistent with the results obtained on Fe+Pt/A1,03 samples reduced at 300"C,18 which show that iron is a non selective and low toxic poison of platinum under such conditions.When the sample is reduced at 400"C, the activities for hydrogenation and hydrogenolysis decrease, and the selectivity ratio tends to decrease also. The decrease in this ratio is more pronounced at low iron loading than at high contents. With 0.035 % Fe the effect is clear since an average value of 6 is obtained, compared with 10 for Pd/A12Q3 and >20 for Pd+S/A120,. A reduction at 500°C yields a further decrease in the activities of both reactions, but a relative increase in the selectivity, since the activity ratio of pure Pd/A1203 catalysts is reobtained. From the literature we know that the degree of reduction of iron increases with temperature and with platinum loading on Pt + Fe/A1203 cata1ysts.l 9 9 2o After reduction at 500"C, Mossbauer spectroscopy detects PtFe clusters in a sample characterized by a low Fe/Pt atomic ratio of 0.2; when the iron content increases PtFe clusters are formed but some iron remains as ferrous ions.In a separate study of the reduction as a function of temperature it was observed that a ferrous ion spectrum was obtained at low reduction temperatures, namely 300°C. Similar results were reported by Garten 21 on PdFe/A1203 catalysts. The catalytic activity of well characterized PtFe clusters has been also measured. Bartholomew and Boudart 22 reported that PtFe clusters supported on carbon have a lower activity but the same selectivity for isomerisation of neo-pentane as180 HYDROGENOLYSIS OF CYCLOPENTANE Pt/C.Similarly, Vannice and Garten 2o observed that PtFe/A1203 have the same selectivity pattern for methanation of CO as Pt/A1203. Therefore, it may be concluded that metallic iron is not a selective poison of platinum. It is reasonable to suppose that the reduction of our Pd+Fe/Al,O, samples at 500°C yields PdFe clusters which have a low activity, but the same selectivity as Pd/A1203 . From what is known from the literature on platinum+iron catalysts, the possible interpretation of the results obtained by reducing the sample at 400°C could be a selective poisoning by Fe2+. This hypothesis could explain the influence of tempera- ture and iron content. It is supported by the results obtained by reducing the catalysts TABLE 5.-cATALYTIC ACTIVITIES (lo5 XI01 S-' 8-l Pd) OF IRON CONTAMINATED SAMPLES loading catalyst % Pd % Fe 176 1 0 1702 1 0 341 1.0 347 1.0 342 1.0 345 1.0 351 1.0 354 1.0 362 1.0 365 1.0 343 1.0 346 1.0 348 1.0 3401 1.0 352 1.0 3501 1.0 364 1.0 366 1.0 0.035 0.035 0.035 0.035 0.056 0.056 0.078 0.078 0.035 0.035 0.035 0.035 0.056 0.056 0.078 0.078 344 1 0.035 354 1 0.056 363 1 0.078 activity A 1 for activity A2 for conditions of reduction1OC cyclopentane at 290°C benzene at 140°C R = A 2 / A 1 300 (2 h) 5.3 62 11.7 400 (2 h) 4.2 50 11.9 300 (1 h) 300 (2 h) 300 (4 h) 300 (3 h) 300 (1 h) 300 (2 h) 300 (1 h) 300 (3 h) 400 (12 h) 400 (1 h) 0,400-H 2400 H2fH20 400 (12 h) 400 (12 h) 400 (12 h) 400 (2 h) H2+H20 400 (12 h) 5.0 3.6 6.6 3.6 6.2 3.6 4.8 4.1 3.4 3.2 1.5 0.94 2.9 1.5 1.26 3.05 43 34 56 35 62 30 38 42 17 16 10 23 12 24 5.4 5.0 8.6 9.4 8.5 10.5 10 8.4 7.9 10.2 average 9.1 5 .o 5.0 6.6 5.8 7.9 8 .O 4 7.9 500 (12 h) 0.9 11 12 500 (12 h) 0.5 6.5 13 500 (12 h) 0.7 10 14.3 by wet hydrogen (hydrogen saturated with water at room temperature) : the activities but not their ratio are changed.That treatment will probably not yield metallic iron, but could give Fe2+. In every case, it has been demonstrated here that the selectivity of Pd/A1,03 may suffer large variations in the presence of small amounts of contaminants like iron and sulphur. When an inhibition occurs, these two contaminants have opposite effects on selectivity. Since the phenomenon observed on Pd/Si02 is a relative increase in hydrogenolysis, it may be interpreted by a contamination of the support by iron, which is commonly present in industrial carriers or has been introduced accidentally.The larger effect obtained with silica could be explained by a different reducibility of Fe3+ on silica compared with alumina.23S. FUENTES AND F. FIGUERAS 181 CONCLUSIONS The hydrogenolysis of cyclopentane on palladium catalysts appears to be insensi- tive to metallic dispersion when the support is clean. In the presence of impurities like SO$- or Fe3+, the catalytic properties may be noticeably modified. Both contaminants behave as poisons, but their influence depends mainly on the temperature of reduction of the sample. When an inhibition occurs these two contaminants have opposite effects : sulphur preferentially poisons hydrogenolysis and iron preferentially represses hydrogenation.It is worth recalling here that all the methods which are used to change the disper- sion of a supported metal, like modification of the metal loading or reduction temperature, can also affect the chemical state of the contaminants. The partial pressure of water can also play an important role in this respect. In the case of an impure support it would not be easy to discriminate between the effect of particle size and that of contamination of the metallic surface ; therefore it is useful to compare the results for several supports to obtain safe conclusions on the influence of the particle size for a given reaction. S. F. thanks the " Consejo Nacional de Ciencia y Tecnologia de Mexico " for a scholarship. D. J. C. Yates, W. F. Taylor and J. H. Sinfelt, J. Amer. Chem. Soc., 1964, 86, 2996. D. J. C. Yates and J. H. Sinfelt, J. Catalysis, 1967, 8, 348. M. Boudart, A. W. Aldag, L. D. Ptak and J. E. Benson, J. Catalysis, 1968,11,35. Y . L. Lam and J. H. Sinfelt, J. Catalysis, 1976, 42, 319. R. Maurel, G. Leclercq and J. Barbier, J. CataZysis, 1975, 37, 324. R. Maurel and G. Leclercq, Bull. SOC. chim. France, 1971, 1234. R. A. Dalla Betta and M. Boudart, Vth Int. Congr. Catal., ed. J. W. Hightower (North Holland, Amsterdam, 1973), vol. 2, p. 1329. P. C. Aben, J. Catalysis, 1968, 10,224. J . E. Benson, H. S. Hwang and M. Boudart, J. Catalysis, 1973,30, 146. lo J. E. Germain and R. Maurel, Compt. rend., 1958,247,1854. l 1 S . Szepe and 0. Levenspiel, Proceedings of the Fourth European Symposium on Chemical l2 0. Levenspiel, J. Catalysis, 1972, 25,265. l3 J. E. Benson and M. Boudart, J. Catalysis, 1965,4,705. l4 G. R. Wilson and W. K. Hall, J. Catalysis, 1970,17, 190. l6 F. Figueras, R. Gomez and M. Primet, Ada Chem. Ser., 1973,121,480. l S J. Barbier, mesis (Poitiers, 1975). * O M. A. Vannice and R. L. Garten, J. Mol. CataZysis, 1975,1,201. 22 C . H. Bartholomew and M. Boudart, J. Caralysis, 1972,25, 173. 23 T. Yoshioka, J. Koezuka, H. Ikoma, J. Catalysis, 1970, 16, 264. Reaction Ewineering (Brussels, 1968, Pergamon, London, 1971), p. 265. P. C. Aben, J. C. Platteuw and B. Stouthamer, IVrh Ini. Congr. Catalysis (Moscow, 1968), paper 31. M. Boudart, Adv. Catalysis, 1969,20, 153. R. L. Garten and D. F. Ollis, J. Caralysis, 1974,35,232. R. L. Garten, J. Catalysis, 1976,43, 18. (PAPER 7/133)

ISSN:0300-9599    

DOI:10.1039/F19787400174

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Stability of Metal Uncharged Ligand Complexes in Ion Exchangers Part 2.-The Copper + Ethylenediamine Complex in Montmorillonite and Sulphonic Acid Resin BY ANDRE MAES, PAUL PEIGNEUR AND ADRIEN CREMES* Centrum voor Oppervlaktescheikunde en Colloi'dale Scheikunde, Katholieke Universiteit Leuven, de Croylaan 42, B-3030 Heverlee, Belgium Received 14th February, 1977 The thermodynamic stability of the copper + ethylenediamine complex is determined in montmoril- Ionite and a macroreticular sulphonic acid resin using ion exchange data and the complex formation function. Both methods are in excellent agreement and correspond to an increase in the overall stability constant of lo3 for montmorillonite and for the resin. The stepwise formation constants FIE2 of the ion exchanged complex are for montmorillonite and and logm7 for the resin, as compared with values of 10'0.7 and logs3 for the bulk solution.Some tentative interpretation is offered in terms of a decreased ligand solvation in the ion exchanger. and In Part 1 it was shown that metal + uncharged ligand complexes in ion exchangers can be characterized thermodynamically by two different methods: a new ion exchange method which leads to the overall stability constant and the extension of the Bjerrum complex formation function, which specifies the stability in terms of the stepwise formation constants of the intermediate complexes. It is the purpose of this paper to illustrate these methods through a thermodynamic study of the copper + ethylenediamine complex in an inorganic and an organic ion exchanger.The bulk of the experimental work in the area of adsorption of metal uncharged ligand complexes is confined to organic ion exchangers (see Part 1) and very few data are available for inorganic ion exchangers. Moreover, it appears that in some cases such as montrnorill~nite,~-~ mechanisms other than ion exchange may be involved, judging from the fact that such complexes are adsorbed in amounts far greater than the cation exchange capacity. Most of the data are therefore of a semi- quantitative nature in that little effort has been, or could be, made to characterize these adsorbed species thermodynamically. Owing to its very high stability, the copper + ethylenediamine complex is well suited for such a quantitative study: first, the fully coordinated two-complex is formed in a very small excess of the ligand, a condition which eliminates any significant ion exchange involvement of the protonated ligand; secondly, it is possible to carry out the distribution measurements of both the aqueous copper ion and the complex at identical, fairly low pH values, a condition which, in the case of montmorillonite with a pH-dependent charge, is most advantageous. EXPERIMENTAL Two ion exchangers are used in this study : the strong acid macroporous 100-200 mesh sodium lewatite SP 1080 (Merck, analytical grade) and the <0.5 pm fraction of the sodium form of the Camp Berteau montmorillonite, prepared as described ear lie^.^ These materials 182A .MAES, P . PEIGNEUR AND A . CREMERS 183 are converted into the calcium form by equilibrium dialysis with 0.5 rnol dm-3 calcium nitrate.The resin is washed free of electrolyte and air-dried prior to use, its water content being measured over P205. The montmorillonite suspensions, adjusted to a clay content of 10 g dm-3 are equilibrated with 0.005 mol dmd3 Ca(N0,)2, prior to use. The cation exchange capacity of the resin is 4.2meqg-I dry resin, as measured by isotopic dilution methods, using 45Ca or 22Na. The ion capacity of the clay, as measured by similar methods varies between 1.05 (Na) to 1.15 (Ca) meq g-l at a pH N 6. The calcium-copper ion exchange equilibria are measured at 25°C by shaking overnight known amounts of resin or clay with Ca + Cu nitrate solutions of varying composition at a pH 2~ 5.5 and at constant total concentration (0.005 mol dm-3). The effect of ethylene diamine (en) is studied using similar methods, adjusting the en/Cu molar ratio to 2 and adding an excess of en corresponding to 2.10-4 niol dm-3.The pH in all systems is in the range 6.8-7. These conditions are consistent with the formation of the Cu(en), complex, the stepwise stability constants of which are : log K1 = 10.72 ; log K2 = 9.31.5 Distribu- tion data are obtained by analysing the equilibrium solutions for copper, using (carbon rod) atomic absorption and 45Ca isotopic dilution. The coordination of the copper ion in the ion exchanger is obtained radiometrically from a duplicate experiment, using 14C-labelled en. The formation function of the copper complex in the resin is measured by equilibrating known amounts of Ca-resin with a constant amount of copper (11 meq g-l), the en/Cu molar ratio varying between 2 and 0.4.The pH values vary in the range of FZ 5 (low molar ratio) to 7 (high molar ratio). A different procedure is used in the case of montmorillonite. 10 cm3 of Na-clay suspension, dialytically equilibrated with 0.01 rnol dnr3 NaN03 are added to 40 cm3 of a solution of sodium + copper nitrate at constant total normality (0.01) containing 1.2 meq copperlg clay. The en/Cu molar ratio is varied from 0.3 to 1.9 and the pH of all systems is adjusted to 7. All systems are shaken overnight and centrifugated. Equilibrium solutions are analysed for copper and en (radiometrically) and their pH measured, using an extended scale potentiometer. dool spacings for the montmorillonite saturated with Cu(en), are measured on an air-dry film, using a Philips PN 1051 diffractometer (CuK,, 3, = 1.5418 A).Heats of exchange are measured at 25°C with the LKB-10700-2 batch microcalorimeterY6 using the mono- protonated enH+-montmorillonite as starting material. The enHf clay is prepared by dialysing Na-montmorillonite with a 0.1 rnol dm-3 en solution at pH = 10.3 and asubsequent series of equilibrations with a 0.05 mol dm-3 en solution at pH = 11.5. Heats of exchange are obtained for the reactions enH+-Cu(en)z+ and enHf-Cs+ and the term for the hypo- thetical reaction Cu2+-Cu(en)$ i- is obtained indirectly from the Na+-Cu2+ and Na+-Cs+ data.6* RESULTS ION EXCHANGE DISTRIBUTIONS The experimental ion exchange isotherms of the aqueous copper ion and the copper + ethylenediamine complex in Ca-montmorillonite and Ca-lewatite are shown in fig.1. In order to afford a quantitative illustration of the effect of en on the ion exchange behaviour of the copper (complex) ion, the data have been represented as the equivalent fraction of copper in the ion exchanger as a function of the logarithm of its equivalent fraction in the equilibrium solution. It was established from the en distribution data that the copper species involved is in fact the Cu(en);+ complex: the average ligand number of copper in the ion exchanger varies between 2 and 2.05. In assigning all of the adsorbed ligand to the copper ions, it is assumed that the ion exchange adsorption of the protonated forms enH+ and enH$+ is occurring to a negligible extent.Such an assumption is easily justified in view of the fact that the equilibria were carried out at 0.005 mol dm-3 total concentration and in the presence of a very low excess of en (2 x mol dm-3). Under the pH conditions used, such an excess would correspond to a concentration of enH+ and enH$+ of fi: mol dm-3, which is nearly two orders of magnitude184 STABILITY OF COMPLEXES IN ION EXCHANGE 4 3 2 1 0 - log SC" FIG. 1.-Ion exchange isotherms in montmorillonite and lewatite for Ca2+-Cu2+ (0, A) and Ca2+-Cu(en)~'(0, A). The dashed curve refers to unit selectivity coefficient : the lower curve refers to lewatite. 7 6 5 4 3 2 I 0 0 . 2 0.4 0.6 0.8 ZC" FIG. 2.-Logarithm of selectivity coefficient against exchanger composition in montmorillonite and lewatite for Caz+-Cu2+ (0, and Ca'+-Cu(en);+ (0, A); the lower curve refers to lewatite.A .MAES, P . PEIGNEUR AND A . CREMERS 185 below the calcium concentration in the equilibrium solution. This assumption is further verified by two experimental facts: first, the adsorption of en in calcium montmorillonite, measured under similar conditions as for the ion exchange isotherms but in the absence of copper, amounts to 1 to 2 % of the exchange capacity ; secondly, throughout the composition range of the exchanger, the sum of adsorbed copper (complex) and calcium ions equals the exchange capacity, pointing to a one-to-one stoichiometry. In montmorillonite, these values vary in the range 1.10-1.1 5 meq g-l and for lewatite, 4.10-4.35 meq g-l. Therefore, we are confident that no mechanisms, other than ion exchange, are involved in the process.TABLE THERMODYNAMIC DATA FOR THE Ca-CU AND Ca-CU(en)z EQUILIBRIA equilibrium In K AGo/kJ mol-1 mon t morilloni t e 0.25 - 0.63 lewatite -0.60 +1.50 7.00 -17.15 2.50 - 6.27 { ZCa+ Cu sZCu+Ca Z C ~ + ~u(en), + ~ ~ u ( e n ) ~ + ~a Fig. 1 shows that for both ion exchangers, there is only a very small difference in ion exchange affinity between calcium and copper. In montmorillonite, Cu is slightly preferred, while the opposite is true for lewatite. The addition of en has a significant effect on copper adsorption: in lewatite, this effect corresponds to a one-order-of-magnitude lowering of the copper concentration in the equilibrium solution, whereas in montmorillonite the shift amounts to three orders of magnitude.All selectivity data are summarized in fig. 2, which shows the variation of the logarithm of the selectivity coefficient against ion exchanger composition. In view of the low total concentration, we may identify the molar ratios with activity ratios. are summarized in table 1, along with the standard free energy values. These results correspond to In K values of 6.75 (montmorillonite) and 3.1 (lewatite) for the reversible displacement of the aqueous copper ion by the Cu(en);+ complex. If we assume a unit partition coefficient for the uncharged ligand, then these data correspond to stability enhancements of - 16.8 (montmorillonite) and - 7.8 (lewatite) kJ mol-l, i.e. the overall stability constant of the Cu(en)$+ complex is raised by factors of The thermodynamic equilibrium constants, as obtained by graphical integration and respectively.COMPLEX FORMATION FUNCTIONS The complex formation data for the Cu(en)$+ complex in both exchangers are represented in fig. 3. In the case of montmorillonite, the data correspond to copper loadings 2,- in the range 0.85-0.95, these values referring to the lower and higher ligand concentrations. For lewatite, these loadings vary in the range 0.3 to 0.9. The free ligand concentration in the equilibrium solution is calculated from pH values, copper concentration and overall ligand content using the formation constants for the copper complex, reported in the experimental section and log K,(enH+) = 10.17 ; log K,,(enH;') = 7.49.5 The ligand numbers E of copper in the ion exchanger are net values, i.e.the ion exchange contribution from the protonated forms (predominantly enH5-t) is sub- tracted from the overall ligand content. This correction was estimated using an upper limit of K, = 10 for the Ca-enH;+ equilibrium in the resin and K, = 30 for the Na-enH$+ equilibrium in the clay. This correction is quite small, at most 0.05 in 5, as can be expected on the basis of the very low concentrations of enH$f, relative186 STABILITY OF COMPLEXES I N I O N EXCHANGE to Na+ and Ca2+ ions (enH;+ M mol dm-3 for the clay) and the fact that enH2f has to compete with either Na+ or Ca2+ for a small fraction of the ion capacity. Fig. 3 shows that complex formation is initiated at much lower ligand concentra- tion, compared with the bulk solution, the effect being most pronounced in mont- morillonite.The curves in fig. 2 afford a direct estimate of the average formation constant, defined as K1K2, from the reciprocal of the free L concentration, corres- ponding to a degree of complex formation of 0.5; the result is for mont- morillonite and for lewatite. mol dm-3 for the resin and M 20 1.6 12 - n 0.8 0.4 - - - - - - -l 8 9 I 0 r i 12 PL FIG. 3.-Ligand number of the copper ion in montmorillonite (0) and lewatite (0) against logarithm of free ligand concentrations in solution. The dashed curve refers to complex formation in solution, The stepwise stability constants were obtained by least-squaring the data in the equation - K, L + 2K,R,L2 n = I+K,L+R,R,L~ and the resulting data are summarized in table 2.The overall stabilization factor, obtained by this method is lO3*I and 101.32 for montmorillonite and lewatite, a result which is in excellent agreement with the result obtained from ion exchange measurements. TABLE ~.-STABLLITY CONSTANTS FOR THE Cu(en), COMPLEX, OBTAINED FROM COMPLEX FORMATION DATA overall stabilization IogF, logZ2 log82 factor montmorillonite 11.60 11.50 23.10 10 3.1 lewatite 11.65 9.70 21.35 bulk solution 10.72 9.31 20.035 10 1.32 X-RAY AND CALORIMETRIC DATA The 001 spacing of Cu(en)2-montmorillonite, obtained as the average of the first, third and fourth-order reflections is 12.6 AA. MAES, P. PEIGNEUR AND A. CRBMBRS 187 The microcalorimetric results are summarized in fig. 4. The ion exchange displacement of enH+ by Cu(en):+ is slightly endothermic, whereas the reaction with Cs+ is considerably exothermic.Using literature values for the AH terms for Na+-Cs ’ and Na+-Cu2+ exchange, we estimate the heat terms for the hypothetical reaction Z Cu,, + Cu(en):+ + Cu(en)$+ + Cu&+ to be about - 12 kJ mol-l. cnH Ns I i I I 1 I I f+ 18.0 t t f I I I I I I I I 1 I cs FIG. 4.-Heats of exchange reactions (kJ/equivalent) in montmonionite. DISCUSSION The foregoing data show that the Cu(en)$+ complex is strongly stabilized in both ion exchangers. The stabilization corresponds to a factor of loo0 (montmorillonite) and 20 (lewatite) in the overall stability constant, assuming a partition coefficient of unity for the ligand. The result found for the resin is in excellent agreement with the value which may be calculated from the data of Cockerel1 and Walton on the same complex (logp, N 21.5) and with the data reported by SkoroWood and Kalinina lo (log Kl = 11.39 to 11.77 and log K2 = 9.74-9.83, depending on moss linking).The nature of the agreement tends to show that the stabilization is insensitive to the degree of cross linking, since our data were obtained on a highly porous resin, whereas the literature data refer to resins of 10-16 % DVB. The good agreement between the ion exchange method and the complex formation indicates that the reaction is thermodynamically reversible. It should be emphasized that the extent of stabilization, as obtained from ion exchange, refers to the actual stability constant in the equilibrium solution and is therefore independent of its absolute value.A similar situation occurs in the complex formation method where the ratio of stability constants in the ion exchanger and the solution f12/b2 (or K l / K ,188 STABILITY OF COMPLEXES I N ION EXCHANGE and K2/K2) is quite insensitive to the absolute values of the bulk solution stability constant, chosen for calculating the free ligand concentration : for example, a 10 % difference in Kl and K2 has an insignificant effect on the extent of stabilization obtained. No unambiguous interpretation has been given for the stabilization of such complexes in ion exchangers. Cockerel1 and Walton proposed that the diamine acts as a " cross linking ligand " binding metal ions in a chain or network. Sometimes the " organic solvent character " differences of ion hydration and lower dielectric constant in the ion exchanger 1 1 * l2 are thought to be responsible for complex stabilization.In montmorillonite, the X-ray data, which agree with literature data,l 3-1 point to the presence of a 3 A monolayer between the lamellae and it appears that Cu(en)$+ is present as a monomeric species: its conformation is essentially identical to the one in aqueous solution, i.e. a square planar complex in which the two axially coordinated water molecules are replaced by oxygens from the clay lattice. Such a configuration was also proposed for the hydrated copper ion under conditions where a monolayer is present in the interlamellar region and where the copper ion is thought to be coordinated to four water molecules in X Y plane.It appears therefore that the cross-linking hypothesis is not tenable, at least not for mont- morillonite, since monodentate ligands such as thi~urea,~ pyridine and alkylamines, to be discussed in subsequent papers, are found to lead to equally important stabiliza- tion phenomena in montmorillonite. Invoking the organic solvent character as an explanation does not seem realistic either, in view of the much larger effects in montmorillonite. If, of course, the observations were limited to organic resins, it could be argued that the stabilization is merely an artefact resulting from a very favourable distribution of the ligand between the ion exchanger and the liquid phase, causing the adsorbed metal ion to be " exposed " to higher ligand concentrations than those measured in solution.However, the less pronounced effect in resins fails to support this idea. In any case, the fact that the metal complex exhibits a very high preference for the ion exchanger under conditions (high ligand concentration) corresponding to the formation of the same coordinatively saturated complex in both cases, clearly points to a thermo- dynamically real effect. It is likely that the reason for the enhanced stability of ion exchanged complexes is related to the effect of the ion exchanger on solvent properties and that the more important effects in montmorillonite are connected with surface geometry. Of course, bulk properties such as dielectric constant can be invoked, but it seems preferable to search on the molecular level. It is apparent that the origin of the stabilization in montmorillonite is mainly energetic in origin.The experimental value of - 12 kJ mol-1 is subject to some uncertainty, being determined indirectly. However, unpublished data from this laboratory, obtained for the same complex on hectorite from temperature dependence of ion exchange equilibria show a similar trend. That the origin was predominantly energetic was also shown from diffuse reflectance spectroscopy on dehydrated Cu(en)$+ montmorillonite and interpreted in terms of an increase in crystal field stabilization energy. Turning our attention to the solvent effect, it is customary to analyse differences in solvent properties in the two phases exclusively in terms of possible effects on cation hydration. In view of arguments presented above, it seems that the primary hydration of the copper ion in montmorillonite is hardly affected, apart from the replacement of the axially coordinated water molecules by lattice oxygens.In the present contextA . MAES, P . PEIGNEUR AND A . CREMERS 189 however, it is important to examine possible effects of a change in solvent properties on the ligand. In a series of papers connected with the enhancement of thermodynamic stability of metal complexes with tetramine macrocyclic ligands, compared with the non- cyclic analogues, Margerum and co-workers 17-19 introduced the concept of a " macrocyclic effect ". It was shown that the stability enhancement was mainly due to a more favourable enthalpy term which could not be described in terms of stronger metal-nitrogen bonds in the cyclic ligands.The enthalpic differences were attributed to a decreased ligand solvation of the cyclic ligand which has less amine hydrogen bonded water to be displaced in the complex formation reaction. Such ligand solvation effects " will hold for any ligands where the donor groups are forced to be close to one another or in some way are shielded from solvation ".18 This situation could occur in the shallow interlamellar region in montmorillonite where hydrogen bonded solvent-lattice structures have been demonstrated.20 Admittedly, these are little more than conjectures, the value of which remains to be further tested, perhaps through a thorough thermodynamic study of a number of complexes in various ion exchangers.The financial support of the Belgian Government (Programmatie van het Wetenschapsbeleid) is acknowledged. A. Maes, P. Marynen and A. Cremers, J.C.S. Faraday I, 1977, 73, 1297. S. L. Schwartzen Allen and E. Matijevic, J. Colloid Interface Sci., 1975, 50, 143. M. H. El-Sayed, R. G. Burau and K. L. Babcock, Soil Sci. Soc. Amer. Proc., 1971, 35, 571. J. Pleysier and A. C. Cremers, J.C.S. Furaday I, 1975, 71, 256. G. L. Sill6n and A. E. Martell, Stabllity Constants of Metal Ion Complexes (The Chemical Society, London, 1964-1971). ti A. Maes, P. Peigneur and A. Cremers, Proc. Int. Clay Conf. (Applied Publishers, 1975), p. 319. A. Cremers and H. C. Thomas, Israel J. Chem., 1968, 6, 949. G. L. Gaines and H. C. Thomas, J. Chem. Phys., 1953,21,714. L. Cockerell and H. F. Walton, J. Phys. Chem., 1962, 66, 75. lo 0. R. Skorokhod and A. A. Kalinina, Russ. J. Phys. Chem., 1975,49,187. l1 0. R. Skorokhod and A. G. Varawa, Russ. J. Phys. Chem., 1972,46,980. l2 Y. Marcus and A. J. Kertes, Ion Exchange and Solvent Extraction of Metal Complexes (Wiley, l 3 W. Bodenheimer, L. Heller, B. Kirson and S. Yariv, Clay Min. Bull., 1962, 5, 145. 14R. Laura and P. Cloos, Proc. Rkunion Hispano-BeIga de Minerales de la Arcilla (Madrid, l5 F. Velghe, R. A Schoonheydt, J. B. Uytterhoeven, P. Peigneur and J. H. Lunsford, J. Phys. l6 D M. Clementz, T. J. Pinnavaia and M. M. Mortland, J. Phys. Chem., 1973,77, 195. l7 D. K. Cabbiness and D. W. Margerum, J. Amer. Chem. SOC., 1969,91,6540. '* F. P. Hinz and D. W. Margerum, J. Amer. Chem. SOC., 1974,96,4993. l9 F. P. Hinze and D. W. Margerum, Inorg. Chem., 1974,13,2941. 2o V. C. Farmer and J. D. Russell, Trans. Faraday SOC., 1971,67,2737. London, 1969). 1970), p. 76. Chem., submitted for publication. (PAPER 7/243)

ISSN:0300-9599    

DOI:10.1039/F19787400182

出版商:RSC    

年代:1978

数据来源: RSC