摘要:

Coupled Fluxes in Electrochemistry Concentration Distributions near Electrodialysis Membranes BY JOHN F. BRADY-~ AND J. C. ROBIN TURNER* Department of Chemical Engineering, Pembroke Street, Cambridge CB2 3RA Received 15th February, 1978 In electrodialysis the gradient of electrical potential leads to coupling of the flux equations for all the different ionic species. When these ions have the same (numerical) valency, their concentra- tions close to and within the electrodialysis membrane can be obtained from a closed-form solution, as can the potential. When they do not, a series solution is derived, which requires that all the ionic fluxes be known. If they are not available (e.g. from measurement) they must be guessed at the start of an iterative scheme. Concentration and potential profiles are discussed for cases with and without “ water-splitting ”, which is the production of sizeable fluxes of H+ and OH- ions even though the feed solution is effectively neutral.The concept of “ limiting current density ” is discussed, and is shown to lose much of its usefulness in the presence of water-splitting. Many problems in electrochemistry require the solution of a set of coupled partial differential equations known as the transport equations. Obtaining such it solution poses impossible analytical, or numerical, problems for systems of arbitrary geometry and flow field. Fortunately, simplifications can often be made without damaging the usefulness of the results obtained. First, the geometry is frequently simple ; in electro-dialysis one can consider the ionic transport to be uni-dimensional, normal to the plane of the ion-exchange membrane.Secondly, the flow field may be a simple one; if the transfer is to such a planar interface, the flow velocities close to it may be well defined. The simplest model to take is the Nernst film model, in which the diffusive fluxes and concentration gradients are confined to a region close to the interface, where the flow velocity is assumed to be zero. A rather more complicated model is that appropriate to the rotating disc electrode, as described by Levich [ref. (l), p. 601. Here the rotation of the electrode produces a flow pattern which is not only well-defined, but also possesses some convenient characteristics. Given such a simplified geometry and flow pattern, we must choose a frame of reference with respect to which the motion of ions is defined.In electrodialysis an obvious candidate is a frame fixed with respect to the membrane interface. If the model allows for a net fluid velocity, there will be a convective term in the transport equation for each ion. Secondly, the gradients of chemical potential of the ions can be regarded as thermodynamic forces producing contributions to the flux of each ion. Not only will the gradient of chemical potential of any given ion affect its flux, an effect expressible by a coefficient Lit, but also those of all other ions, by the ‘‘ cross-coupling ” coefficients, Li j , of irreversible thermodynamics.2 A simplify- ing assumption frequently made is that these cross-coupling coefficients, Ltj, i # j are zero, and in dilute solutions one might hope that they would be so.However, T Present address : Department of Chemical Engineering, Stanford University, Stanford, California, U.S.A. 1-90 28392840 COUPLED FLUXES I N ELECTROCHEMISTRY one can do better than that. If specific interactions are not allowed, one is able to account for the effects of all other ions by allowing the Lii to be appropriate functions of the total concentration, or ionic strength. This having been done, it is nevertheless customary still to describe the transport equations as " coupled " since there is, in each of them, a term involving the electrical potential gradient which, because the ions are charged, will have an effect on the velocity of each species of ion.Since this potential gradient is itself dependent on the fluxes of the ions, the flux of each species of ion is dependent on the fluxes of all other species. If the solution is dilute (and in electrodialysis it is dilute solutions which give rise to the most interesting effects) it may be permissible to disregard the gradients of activity coefficients, and thereby to assume that the gradient of chemical potential is equal to the gradient of the logarithm of the concentration. Here again, if the activity coefficient is not affected by specific interactions, general interactions e.g., of the Debye-Huckel limiting type) may be accounted for by appropriate variation of the Lii coefficients. We shall thus consider the following particular form of the equation for the flux, ATf, of ions of species i : This is a special case of the equation discussed by Levich [ref. (1) p.2791. We are considering one-dimensional transport in a dilute solution. It has been assumed that the region of interest is in a steady state; any changes in the bulk phases bordering the stagnant film, or boundary layer are slow, and our region of interest can '' keep up " with such changes comparatively rapidly. In eqn (1) Nf is measured in kg ions m-2 s-l, Df (the Fick's Law diffusion coefficient) in m2 s-l, and x (the distance normal to the plane membrane surface) in m. The valency of the ion is zt and its concentra- tion ci kg ions m-3. The dimensionless electrical potential 4 is equivalent to volts x FIRT, where F is the Faraday constant. Finally v m s-l is the velocity normal to the membrane surface.This may often be negligible, but it can be retained without difficulty, and since in electrochemical systems there may be a net volume flow through the interface it should in principle be included. We have neglected in eqn (1) any term due to pressure gradient. This would not be acceptable if we were concerned with reverse osmosis, but in electrodialysis it is ju~tifiable.~ We have also not allowed for any chemical reaction term in eqn (I), and we will consider that possibility later. It is instructive to consider what magnitude of " diffusion film thickness ", S, is required to explain the observed behaviour, noting that 0 c x c 6. A review of the available data indicates that 6 is of the order of to m.How does this compare with the magnitude of the electrical double-layer at the surface of the membrane, on the one hand, and the fluid mechanical, or Prandtl boundary layer thickness on the other ? The Debye length gives a measure of the size of the double- layer. It varies as l/Jc, but even for normal solutions it is still only my which is very small compared to the to The relationship between the " diffusion film thickness " and the " boundary layer thickness " is discussed in many texts, e.g., Vetter.s For turbulent mixing in the bulk phase these two thicknesses are in the ratio of Sc", where Sc, the Schmidt number, is the ratio of the kinematic viscosity to the diffusion coefficient, and n is generally considered to be 3. For aqueous solutions the Schmidt number is of order 1000, and this means that the diffusion film thickness is about one-tenth of the boundary layer thickness.This makes the use of eqn (1) justifiable, since the processes m mentioned above for 6.J . F. BRADY AND J . C. R . TURNER 2841 are confined to a region very close to the membrane surface, where velocities parallel to the surface are small. The application of eqn (1) to isothermal conductance and diffusion is straight- forward; in fact the interest in these fields lies in considering those interactions, or “ non-idealities ” which have been “ simplified out ” of eqn (1). Its application to concentration polarisation at electrode surfaces is usually simple, too, since often only one electrode reaction occurs (or if more than one, at well-separated potentials) and the ion fluxes are therefore separately determined.With electrodialysis, however, two complications arise : (i) In general the fluxes of all the species in a mixture have to be considered, since ion-exchange membranes will allow many species to enter (unlike an electrode, at which frequently only one species may disappear). (ii) The known conditions are usually the bulk concentrations on either side of the membrane. Between these known compositions lie the ion-exchange membrane, with two “ boundary-layers ”, one at each surface, as shown in fig. 1. DILUATE BULK SOLU TI 0 N CONCENTRATION PROFILE CONCENTRATE BULK SOLUTION t I? C O N V EC TI 0 N 1 ’ 7 DIFFUSION j i ~ I F F X ~ O N j 8 MIGRATION 8 MIGRATION I E R i T I O N ] FIG. 1 .-Predominant transport processes occurring close to an electrodialysis membrane.The problem is thus a split boundary-value problem in a multicomponent system, where there is really only one independent variable, the total voltage applied from bulk solution to bulk solution. The answer to the question “ what will be the fluxes and the intermediate concentrations and potential ? ” will involve trial and error. This paper attempts to define a sensible approach to such an answer. SOLUTION OF THE EQUATIONS We seek a solution to the set of eqn (l), which apply in the three regions shown in fig. 1. Eqn (1) can be rewritten as dc, d4 C ~ U - +ziti - -- = -ai, dx dx Di2842 COUPLED FLUXES I N ELECTROCHEMISTRY where a, = N,/Di. We also have the electroneutrality condition Here we note that zi includes the sign of the valency, and that Xis the total equivalent concentration of immobile ionic groups including their signs. For a free solution X is zero, but in an ion-exchange membrane it is given by the capacity of the resin.If the system, which is non-reacting, is in the steady-state, then the Ni are constant with x. In our analysis we shall assume that these fluxes, and hence the ai, are known. The solution of eqn(2) requires integration constants, which we will take to be the concentrations, co, and potential, Sbo, at x = 0. Other sets of conditions are possible, but this choice allows a straightforward solution of eqn (2). Schlogl also obtained a solution of eqn (2) in the case o = 0. He was primarily interested in the fluxes within an ion-exchange membrane, and he assumed the concentrations ci to be known at each face of the membrane and 4 at one of them.His solution then generated the concentration and potential profiles, and the fluxes. Schlogl's method treats the problem as a 2-point boundary value one; our method treats it as an initial value problem. The choice depends on whether one has a better estimate of the fluxes or of the concentrations. In our electrodialysis problem there are three contiguous layers and the concentra- tions at the inner boundaries cannot be known beforehand. One could use Schlogl's analysis and guess these concentrations, but this requires a double set of guesses. Our procedure is to assume that one set of bulk concentrations is known and to guess the different fluxes (these fluxes can often be measured, e.g., as described by Makai).' Since each flux has the same value in all three layers, only a single set of guesses is needed.From our solution, the potential distribution and the concentra- tions in the other bulk liquid can be obtained, which provides the necessary checks for an iterative scheme. The equilibrium relationships at each membrane-solution interface must be known, but that would be so with Schlogl's method also. If we take eqn (2), multiply by zi and sum over i, we obtain (assuming Xconstant with x) - [-c ziai+v ( z i c i / D i ) ] ( ~ z:ci)-' d# dx i I i - - in which we note the factor Substitute eqn (4) in eqn (2), obtaining $ci, which is twice the ionic strength. i (4) in which the summation index has been change to j .v = 0 in eqn (5) and summing over i, we obtain It is helpful to consider first the case of a free electrolyte with v = 0. Putting However, in eqn (6) we note that zicI = 0 for a free solution, X = 0. I Thus eqn (6) simplifies greatly to d -(C ci) = -C ai. dx i iJ . F . BRADY AND J . C. R . TURNER 2843 Since the ai must be constant with x, there is a linear gradient of the total concen- tration. (i) n-n ELECTROLYTE We now put lzrl equal for all the ions and shall call this case " the n-n electrolyte ''. It is the same as Schlogl's case of " nur eine Gegenionen-Klasse ". Solving eqn (7), and substituting in eqn (4), again taking v = 0, one can easily solve for 4 (since the in z?c, can be taken out of the summation). The result is t 4 = #O+M1 In [l -M2x], where and +O, cp are the values at x = 0.All of these quantities are assumed known. By substituting the solution of eqn (7) into eqn ( 5 ) we obtain, for z1 = 0 and X = 0, For the special case M1zi = - 1 eqn (9) takes the form - - 1 though, as can be seen from the definition of M I , this is an unlikely occurrence [e.g., the case of aj = 0 for all ionsj of the same sign as ion i, in which case Ci = (1 - M,x)cf for i and all those j ] . Eqn (8) and (9) are equivalent to Schlogl's for the case z1 = X = 0, though such is not instantly apparent since the quantities assumed known are different in our case from in his. In eqn (8) it can be seen that as M2x + 1, 4 3 - 00. As we show in the next section, which deals with the general case, this condition (that x < l/MJ gives the radius of convergence for the series representation of the concentration distributions.Values of x larger than 1/M2 have no meaning. For values of x smaller than 1/M2 d, may be finite, but one or more of the cf from eqn (9) may be negative. This is clearly not allowable, and indicates that a physically impossible set of ai have been chosen. We shall revert to this point when considering some results later. (ii) GENERAL CASE We have not been able to obtain a solution in closed form for the case where lzll is not the same for all i. The difficulty arises because we cannot use the solution of eqn (7) in the summations in (4) and ( 5 ) since z: is not a common factor for all ions. The situation is thus inherently much more complicated, as Schlogl stresses in his paper. However, there is a series solutions of the form ci = C A~,,x".n=O Substitution into eqn (4) and (5) and comparing coefficients gives2844 COUPLED FLUXES IN ELECTROCHEMISTRY The boundary condition cf = cp at x = 0 gives (13) The succeeding coefficients can then be calculated from eqn (12). This is rendered simpler if 11 = 0, which can often be assumed. It should be noted that involves the A j , n for all lower n and all speciesj. Also, eqn (12) does not explicitly involve X , which can be regarded as entering when all the cp are included ; it is, however, general in that non-zero v and X are allowed for. The Ai,n can be computed for given cp and ai, and substituted in eqn (11) to give Ci at any given x. The radius of convergence, within which this series representation is valid, is not simple to calculate.For the n-n free electrolyte, eqn (8) shows that x must be less than 1/M2. The ratio of Ai,n+l to Ai,n as n + 00 is the reciprocal of the convergence ratio. This ratio is given by Z ; A ~ , ~ / E Z ~ A ~ , ~ [the first term in the series in eqn (12)] under rather special conditions, which seem to be that all the positive ions have equal valency and all the negative have equal valency (not necessarily the same as the positive ions). This certainly includes the n-n electrolyte, for which this ratio equals Mz. In general, however, the convergence ratio has to be estimated numerically from the computed coefficients. This is no real problem if a computer is available, though we have found that it may be necessary to compute a hundred or so coefficients before a precise estimate of the radius of convergence can be made.The graphical method of Domb and Sykes, as described by Van Dyke,8 is very helpful, and since the extrapolation is often linear, the computer can simply provide the answer without resort to graphing. We did come across cases where the af and cp gave oscillating which oscillations showed no sign of dis- appearing up to n = 1000. Here the convergence ratio was estimated from the heights of successive peaks in the &. Programming and computing time can be shortened if the following relationships are noted : first 0 Ai,o = ~ i . j i n m = l Pl 2 1. This comes from the electroneutrality condition, eqn(3), and even holds for n = 0 if X = 0.We also have Aj,n = 0, n 3 2, if X = 0 and v = 0. (1 5 ) J This comes from eqn (7), which requires a linear gradient of total concentration for a free solution. The potential t$ can be calculated from the following integral form, which follows directly from eqn (4). As x is increased towards the radius of convergence, 141 increases without bound. The concentrations ci, given by eqn (ll), will also show similar behaviour, and some of them will be negative for a range of x, which of course has no physical meaning. It should be remembered that the ai were taken as known. The production of nega- tive values of some cf for mathematically-acceptable values of x indicates that the chosen ai are physically impossible for these values of x . Fig. 2 shows an example of this.J .F . BRADY AND J . C. R . TURNER 2845 As the potential difference across a given film thickness, 6, is increased without limit, the fluxes, and hence the at, increase towards limiting values. These give the " limiting current density ", and represent the unique set of ai which simultaneously give 6 as the radius of convergence of eqn (1 1) md give zero values of all the ci at x = 6. For lower values of the potential difference across 6 an infinity of sets of the ai can be applied to eqn (12) and give bounded values of the ct from eqn (1 1). Even rejecting those sets of at which give any negative values of cl: within S leaves an infinite number remaining, for each of which there is a set of positive ci at x = 6. There is, however, only one set of ai which will make all cf zero at x = 6, and for these A+ across the film will be infinite.In practice, under conditions of severe polarization, " water-splitting " will occur. This is the production of fluxes of H+ and OH- ions which can become greater than 0.2 0.L 0.6 0.8 ?.O 1.2 1.4 1.6 XIS FIG. 2.-Concentration and potential profiles in the diluate diffusion layer near a cation membrane. NaCl+KCI mixed solution, each at 0.01 kmol m-3. Profiles are calculated, assuming no water- splitting. the fluxes of the ions present in the feed solution. These fluxes can be measured, if not predicted, and the associated as included in the above treatment. Their presence requires modification of the simplified conclusions of the last paragraph, which would fully apply only in the absence of the dissociation of water.The major effect is that as + is increased many of the ai continue to increase, that for the H+ or OH- ion most markedly. An example is given later. RESULTS AND DISCUSSION We start by showing three specimen results, all of which refer to the free electrolyte case, with II and X both equal to zero. The first case is the electrodialysis of an equimolar solution of NaCl and KC1, each at bulk solution concentration kg mol m-3. We are interested in the profiles in the diluate compartment close to the surface of a selective cation exchange membrane. The transfer of Cl- ions is small, but not negligible, in comparison with those for the Na+ and K+ ions. For calcula- tion purposes, we take a total current density of 150 A m-2 (which is a reasonable figure for electrodialysis).Taking appropriate transport numbers and diffusion2846 COUPLED FLUXES I N ELECTROCHEMISTRY coefficients, we obtain values of the ai equal to -38.3 kg ions for the Cl- ion, 643 for the Na+ ion and 319 for the K+ ion. The stagnant film thickness, 6, is taken as 0.02 mm. Fig. 2 shows the profiles of the ion concentrations and the potential. Values of X I S greater than 1 have no physical meaning. We note the minimum in the K+ concentration, which increases close to the membrane surface even though the ion is moving towards the membrane. Since this is an n-n electrolyte, there is a linear fall in the ionic strength, which means that the C1- ion concentration falls linearly. The radius of convergence for this example is at x/6 = 2.17, where the C1- concentra- tion would be zero, the K+ concentration + a, the Na+ concentration - 00, and the potential, 4 would also be negative infinite.This is a calculated example, and in practice there would have been some “ water-splitting ”. Greater current density (or a thicker boundary layer) would lead to more water-splitting, and such a case is shown in fig. 3, for which fluxes and bulk concentrations experimentally measured by FIG. 3.-Concentration profiles during electrodialysis. Measured fluxes, including those due to water-splitting, have been used to calculate these profiles. Makai have been used. These refer to the electrodialysis of NaCl solution which has become slightly acid as a result of the electrodialysis.We are interested this time in the profiles in the solutions close to the anion membrane. The total current density was 330 A m-2 and the values of ai on the diluate side were 738 kg ions m-4 for the C1- ion, - 11.2 for the Na+ ion, and -228 for the H+ ion. The Na+ transport arises from the imperfect selectivity of the anion membrane. The much larger H+ flux comes from water-splitting and, in fact, leads to a greater Cl- flux than would be cakulated in its absence, see Makai.’ These are severely polarising conditions, as shown by the convergence radius, 0.0449 mm, being very close to the boundary layer thickness, 6. Fluid mechanical correlations gave a value of 6 equal to 0.045mm. It cannot in fact be greater than the convergence radius, but the correlations are not precise enough for this purpose.If, for illustration, we take a value of 6 equal to 99 % of the convergence radius, then fig. 3 can be drawn, using the equations for the n--y2 electrolyte. The H+ concentration shows a marked maximum within the boundary layer. We stress again that we have used the measuredJ . F. BRADY AND J . C . R. TURNER 2847 flux of H+ ion ; how that flux arises will be discussed in another paper devoted to the water-splitting phenomenon. Fig. 3 also shows the profiles in the boundary layer on the concentrate side. Here the Hf flux (which was to the Zeft on the diluate side) is replaced by an equivalent OH- flux to the right (within the membrane and the boundary layer on the concentrate side). We have not attempted to calculate the concentration profiles within the ion exchange membrane, but have worked back from the known bulk concentration on the concentrate side.This is because (a) the concentrations within the ion-exchange membrane are very much higher than in the solutions, and therefore the potential drop across the membrane is comparatively small, and (b) in a case of severe water- splitting such as this the concentrations at the surface of the membrane in the diluate are extremely sensitive to the thickness chosen for the boundary layer. For a less severely polarising situation, e.g., with a mixture of KCl and NaCl, this can be done, and a check of the equilibrium at either membrane surface can be used to check the calculated procedure. 3.0 I I I I 0.2 0.L 0.6 0.8 1.0 x /(convergence ratio) FIG.4.-Concentration profiles during electrodialysis of a mixed-valency system, CaCI, + NaCl. Measured fluxes, including those due to water-splitting, have been used to calculate these profiles. The plot of 6 against x/S is not shown on fig. 3. In the concentrate diffusion layer 4 varies roughly linearly with x/S, rising from left to right. In the diluate diffusion layer, however, it resembles the curve on fig. 2, but curving sharply upwards (instead of downwards) as the anion (cation in fig.2) membrane is approached. This sharp increase in the potential gradient repels the H+ ion to the left away from the membrane, against its concentration gradient. At lower values of x/S the potential gradient is smaller, and the concentration gradient switches sign ; the two gradients now act in concert to provide the Hf flux.The H+ concentration thus passes through a maximum; in fig. 2 the K+ concentration showed a minimum for similar reasons. The final example also uses experimental information and relates to the electro- dialysis of a mixture of CaC1, and NaCl under polarising conditions. This is not an n-n electrolyte, so the series solution had to be used in the calculations. The bulk2848 COUPLED FLUXES I N ELECTROCHEMISTRY concentrations in the diluate near the anion-exchange membrane were 5.60 x kg ions m-3 for Ca2+, 1.17 x for H+ (which arose from water-splitting). The currents carried by the different ions through the diluate boundary layer were 259 A m-2 for the C1- ion and 195 A m--z for the H+ ion. Since the anion-exchange membrane was very selective the currents carried by the Ca2+ and Na+ ions were about 1 A ~ll-~ each.Though the H+ ion is present to only a comparatively small extent in the bulk solution it contributes nearly half the current, due to water-splitting at the membrane surface. The OH- ion so formed passes through the anion-exchange membrane with the C1- ion. Using appropriate diffusion coeficients, the values of the ai for the four ions were 1530 kg ions m-4 for C1-, -246 for Hf, - 9 for Ca2+ and - 11 for Na+. The fluid mechanical estimate of the boundary layer thickness was 0.035 mm. The convergence radius, obtained from the coefficients Ai,n, was 0.033 mm. The simple expression c z : A ~ , ~ / ~ Z ~ A ~ , ~ would have given 0.026 mm, considerably lower than the true value and outside the limit of error of the fluid mechanical estimate.Fig. 4 shows the concentration profiles in the boundary layer up to values of x equal to 95 % of the convergence radius. Once again the H+ concentration shows a maximum. The other concentrations give slightly curved distributions. This approach can be applied to situations of practical interest. The operation of a desalinating electrodialysis stack under conditions of severe polarisation can lead to acid and alkaline product streams. Calculating the concentrations of ions close to the membrane surfaces, as the driving voltage is increased, could lead to an understanding of the water-splitting process. Since the fluid mechanics within an electrodialysis stack are complex (see Makai) ' we carried out similar experiments using a rotating-disc set-up.The fluid mechanics in this situation are well established, see Levich [ref. (1) p. 601, and the results are thereby more easily interpreted. The solutions used in all these experiments involved single salts, and mixtures, of n-n and of mixed-valency types. The results are described in a later paper.9 The procedure described here requires the values of Ni or ai to be known (or to be guessed as a first step of an iterative scheme). In our apparatus, where water- splitting occurs, these fluxes can be measured and used in the theoretical equations. There remains the problem of why and how large fluxes of hydrogen or hydroxyl ions can arise when the feed solution is essentially neutral. It is for this problem that the equation for the distribution of 4 is required; this matter will be discussed elsewhere.for Na+, 2.37 x for Cl-, and 8 x 2 i CONCLUSIONS A procedure for calculating the compositions and potential close to and within ion-exchange membranes in an electrodialytic stack has been described. The analysis of eqn (1) given here may also be useful in other problems of ion transport, such as electrochemical reactions at solid electrodes, the electrolysis of multicomponent systems and transport in biological membranes. The fluxes of the ions have to be assumed, but they are in practice measurable, whereas interface compositions are not. For predictive purposes they can be guessed, and the calculations carried out from one bulk solution to the other, iteratively, until agreement is obtained. If severe water-splitting occurs, which is the most interesting case, the boundary layer thick- nesses have to be very precisely known if the concentrations, and particularly the potential, are to be closely defined. It is for this reason that we carried out measure- ments with a rotating-disc system, for which the fluid mechanics are more accuratelyJ . F . BRADY AND J . C. R . TURNER 2849 defined than in an electrodialysis stack. These experiments are described in the following paper. J. F. B. wishes to thank the Winston Churchill Foundation for the award of a scholarship during the tenure of which this work was carried out. V. G. Levich, Physicochemical Hydrodynamics (Prentice-Hall, Englewood Cliffs, New Jersey, 1962). D. D. Fitts, Non equilibrium Thermodynamics (McGraw-Hill, New York, 1969, chap. 1-5,7. L. H. Shaffer and M. S. Mintz, inPrincipZesof Desalination, ed. K. S . Spiegler (Academic Press, 1966), chap. 6. N. Lakshminarayanaiah, Transport Phenomena in Membrams (Academic Press, New York, 1969), chap. 4. K. J. Vetter, Electrochemical Kinetics (Academic Press, New York, 1967), chap. 1-2. R. Schlogl, 2. phys. Chem. (Frankfurt), 1954, 1, 305. ’ A. J. Makai, Ph.D. Thesis (University of Cambridge, 1977). M. Van Dyke, Quart. J. Mech. Appl. Math., 1974, 27, 720. A. J. Makai and J. C. R. Turner, J.C.S. Faraday I, 1978, 74,2850 (next paper). (PAPER 81264)

ISSN:0300-9599    

DOI:10.1039/F19787402839

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Polarisation in Electrodialysis Rotating-disc Studies BY ALEXANDER J. MAKAI-~ AND J. C. ROBIN TURNER* Department of Chemical Engineering, Pembroke Street, Cambridge CB2 3RA Received 15th February, 1978 Polarisation in an electrodialysis stack is difficult to examine because of the complex flow of the fluids through the stack. A rotating-disc apparatus has been built, for which the fluid mechanics are both well-known and convenient. Using anion-exchange membranes, the current against voltage plots at high voltages show considerable " water-splitting ", with large fluxes of H+ and OH- ions. The results correlate well with the treatments of Levich and of Cowan, and the enhancement of the anion fluxes by the water-splitting can be explained. Using cation-exchange membranes there is much less water-splitting, with anomalously high cation flux enhancement. Though these character- istics are desirable in practice, they are not explicable by the same approach as seems successful with anion membranes.In an earlier paper,' Brady and Turner have discussed the concentration and potential distributions close to an electrodialysis membrane. These depend upon the current density being passed through the membrane, the independent variable of the practical process of electrodialysis being the total voltage applied across a stack consisting of perhaps a hundred or more compartments separated by alternate anion and cation exchange membranes. In practice, the membranes are separated by inert spacers, which also define the compartments and the volume of liquid in them.Through these compartments the liquid is passed, and the fluid mechanics are not simple. As well as this, the composition of the fluid changes significantly as it passes through the compartnient (which is indeed the object of the process). Therefore, measurements of voltages, currents, and demineralization rates will be average over the membrane surfaces of a compartment. As the current density is raised (by raising the applied voltage) polarisation will set in, of variable severity depending on the position within the corn- part ment . We wished to examine whether the polarisation behaviour observed in practice could be described by our theoretical resu1ts.l To make a clear-cut comparison, it was necessary to we a system in which the fluid conditions were better defined than they are in an electrodialysis stack.The rotating-disc electrode, which is a well-known tool in electrochemical kinetics, see e.g., Albery,2 involves fluid-mechanics which are well-defined, and which have been solved theoretically by Levich. We have constructed a similar device using an ion-exchange membrane in place of the electrode, and have measured fluxes through such a membrane under conditions of increasingly severe polarisation. THEORETICAL When a horizontal disc electrode suspended from above is rotated, fluid spinning near the surface is flung out centrifugally, and is replaced by a stream flowing upwards t Present address : Union Carbide Corporation, Charleston, W. Virginia, U.S.A. 2850A . J . MAKAI AND J . C.R. TURNER 2851 towards the disc. Levich solved the problem of the convection combined with diffusion of a substance reacting at the surface of the electrode. He showed that provided that the Schmidt number (the ratio of the kinetic viscosity to the diffusion coefficient of the reacting solute) was high (and it is of order 1000 for liquid-phase solutions) then the gradient of the reactant's concentration was uniform over the whole surface of the disc. It is this property of" uniform availability " which makes the rotating-disc system a valuable tool. The value of this surface concentration gradient, when the surface concentration is reduced to zero by reaction, is given by (aCilaX)surface = WD; 3,- WC;. (1) Here D, is an appropriate diffusion coefficient, m2 s-l, and cp the concentration far from the electrode surface, kg mok3.v is the kinematic viscosity of the solution, m2 s-l, and O.I is the rotational speed of the electrode, rad s-l. W is a constant equal to 0.621. Gregory and Riddiford have shown that W in fact depends upon the Schmidt number, being 0.603 for the Schmidt number equal to 1000, and 0.582 when it is 100. The value 0.621 is that for infinitely large Schmidt number. It is common in boundary layer theory to describe a surface gradient in terms of that thickness, 6, across which the given surface gradient would produce the required change from surface to bulk-fluid conditions. In electrochemistry this is what we mean by the Nernst film thickness, used in the earlier paper.l For the rotating disc, this thickness, 6, is given by The current density i is given by S = W'1D$~*~-3.(2) and its maximum value, ilim, is hence, from eqn (l), where Z, is the valency of the reacting ion, and DR its diffusion coefficient, m2 s-l. These expressions as they stand apply to electrochemical measurements in the presence of " supporting electrolyte ", an excess of an electrolyte which is inert to the electrode reaction and which serves to maintain a very low potential gradient within the solution for any given i. In that case Ds is the same as DR. But in electrodialysis nothing is added ; the aim is to remove electrolyte. Again in electro- chemistry one may be able to consider an electrode reaction involving only one reagent ion, at least over a certain range of potential. In electrodialysis all the ions of a given sign, say cations, can pass into a cation membrane, thereby disappearing at the surface of the " electrode " as if by reaction, andithis will happen whatever the potential range involved.For a solution of a single salt, containing ions of valency Z, and z,, the absence of supporting electrolyte increases the limiting current density across a film of thick- ness 6, the appropriate expression being The question then arises as to which value of Ds to use in eqn (2) to determine 6, or in eqn (5). Eqn (2) was derived for the single reagent ion in supporting electrolyte, with a negligible potential gradient throughout 6. With a dilute solution of even a2852 FOLARISATION I N ELECTRODIALYSTS single electrolyte [the situation is more complicated with mixtures, see ref.(l)] it is not clear which value of Ds should be used in eqn (2) or (5). We have chosen to use the diffusion coefficient of the salt rather than of the ion constituent which passes into the membrane. If the ions have similar diffusion coefficients, the diffusion coefficient of the salt will also be the same. If the ions have different mobilities, that of the salt will be between them. Since Ds occurs only to the 1/3 power in eqn (2), the effect of an incorrect choice of Ds is smaller than might be feared, Albery6 has examined the problem, and has obtained results for the cases (a) where the supporting electrolyte is at an “ appreciably larger ” concentration than that of the reactant, or (b) where all the ions have the same D.The general case, of arbitrary concentrations and values of D, remains for further examination. We have used eqn (5) to estimate the “ Levich limiting current density ”. In practice, as this value is approached water splitting occurs, and fluxes of H+ and OH- ions cannot be ignored. If these fluxes are measured, we can then use the methods of ref. (1) to calculate the concentrations and potential profiles close to the membrane, but a value of 6 is still required. For this purpose we have used eqn (2) with D, equal again to the value for the salt. Cowan and Brown ’ suggested that a plot of A@/i against l / i could be used to locate the onset of severe polarisation. Here A@ is the potential drop across a membrane, or a cell pair if in a stack, in volts. When the current density, i, is raised, the apparent cell resistance, A@/i, stays approximately constant until a value of i is reached above which A@/i starts to increase sharply.In an electrodialysis stack, the sharpness of this change is blunted by the averaging of non-uniform effects across the membrane surface. With a rotating disc, a sharper transition was expected, and it was hoped to relate this to the “ limiting-current density ” given by eqn (4). As the current density is increased, the onset of water-splitting complicates the picture. Assuming the value of 6 given by eqn (2), the concentrations of the ions can be calculated, at any point within 6, from the measured fluxes, including those due to water-splitting. If there is only one type of ion, for example Cl-, (other than OH- ion from water splitting) which can enter the (anion) exchange membrane, it becomes possible to estimate the enhancement of its flux due to water splitting.It should be noted that both the total current density i, and the Cl- current density are increased by water splitting and the resultant fluxes of H+ and OH- ions. Then the calculation is done again, with the H+ and OH- ion fluxes set to zero, and with a trial reduced Cl- flux, until that C1- flux which produces the same concentration of C1- ion at the electrode surface is obtained. It was found that the choice of position within the boundary layer at which this matching was made had virtually no effect on the result. This is the “ unenhanced ” flux, and it can be compared with the ‘‘ en- hanced ”, or actual, flux.The potential distributions are not directly comparable. It will be seen that the “ unenhanced ” fluxes do show plateaux, even under severely polarising conditions leading to large water-splitting currents. To do this, the concentration of the C1- ion is calculated as just described. EXPERIMENTAL Fig. 1 shows the general layout. The membrane disc was formed by fixing a piece of ion-exchange membrane across the end of a flanged glass tube, g. The flat ground flange had an outside diameter of 30 mm and the open area in the centre had a diameter of 10 mm. The 10 mm annular flange minimises the fluid-mechanical edge effects, as required for Levich’s treatment. The glass tube, g, sealed at one end by the membrane disc, was fiiled with a solution, c, called the “ concentrate ” solution.The current-carrying electrode, p, was a flat spiral ofA . J . MAKAI AND J . C. R. TURNER 2853 platinum wire. The membrane/tube assembly could be rotated, at speeds from 150 to 1300 r.p.m., by the motor, m, while dipping into the “ diluate ” solution, d. This was con- tained in a plastic vessel in the wall of which was an ion-exchange membrane i. On the other side of this membrane was a further volume of “ diluate ” containing the other current- carrying electrode, a nickel plate, n, of area ~ 5 0 0 0 mm2. The purpose of the membrane was to prevent electrode products from n reaching the solution d, and it was of anion or cation-exchanger as appropriate for any experiment. Thus using the current-carrying electrodes n and p, connected to a d.c.power source, a controlled current could be passed through the membrane disc r. To measure the potential drop across r two Ag/AgCl potential electrodes, s, were used. One of these was located inside the rotating tube close to the membrane surface, and the other was placed in the diluate solution about 5 mm below the disc. This positioning could not be repeated very precisely in different runs, so the absolute values of the potential-drop measurements were rather arbitrary. However, we shall see that often we are mainly interested in the change in the potential-drop during a run, and provided the electrodes did not move duringlan experiment, such a change could be accurately measured. FIG. 1 .-Rotating-disc apparatus. Before a membrane was used, it was equilibrated with the solution employed in the experiment.During the experiment, readings of the potential drop and current were taken over a range of current densities extending to well above the “limiting” ion transport conditions, for several speeds of rotation. The concentration and temperature of the diluate bulk solution were also measured. The ionic transport rates were measured in separate experiments under similar conditions using constant currents for accurately timed periods with known solution volumes. The probe electrodes were not in place since it was found that these electrodes were damaged by the products of lengthy electrodialysis. Chemical analysis at the end of these runs gave the transport of each ion (as an average over the run).Solutions of single salts, and of mixtures, involving Na+, Ca2+, C1- and SO:- ions were used. The cations were analysed by atomic emission spectrophotometry, the anions by titration. The current carried by H+ (or OH-} ions through the cation (or anion) membrane was estimated by the difference between the total current and the current carried by the other cations (or anions). This could be done since the membranes were almost ideally perm-selective, i.e., there was negligible transport of cations through an anion membrane, and vice versa.2854 POLARISATION I N ELECTRODIALY SIS RESULTS Runs were performed using Neosepta AV-4T anion exchange membrane with solutions containing Cl-, SO%- and a mixture of the two. Another series of runs used Neosepta CL-2.5T cation exchange membrane with solutions containing Naf, Ca2+ and a mixture of the two.I00 200 (1/i)/cm2 A-l FIG. 2.-Cowan plots at different disc rotational speeds. 0.0174 rnol dm-3 NaCl solution. C1- transport through anion disc membrane. Disc speed : 0, 175 ; 0,250 ; x , 350 ; 8 , 740 ; +, 1200 Fig. 2 shows Cowan plots for the electrodialysis of C1- ions from a 0.0174 mol d r 3 NaCl solution at 19°C through the anion disc membrane. Five different disc rotation speeds were used. The results are in two groups, the potential electrodes having been changed between groups. It can be seen that a sharp break is observed r.p.m. 5 i0 15 AcpW FIG. 3.----Ion transport rates. 1290 rpm. 0.0174 rnol dm-3 Na C1 solution. Anion exchange in the curves, the position of which on the l/i axis can be closely estimated by extrapolatioii of the straight line sections of each plot.We shall call this break point the Cowan limiting current density. Timed electrodialysis runs were also carried out, from which the ion transport rates could be calculated. Using a rotation speed of 1290 r.p.m. the results in fig. 3 membrane. x , total ; 0, C1-; 8, OH-.A . J . MAKAI A N D J . C . R . TURNER 2855 were obtained. As A@ was increased above 5V the OH- ion transport, arising from water-splitting, rose from virtually zero to equal the C1- ion transport. The total current showed no plateau, and the C1- transport continued to increase up to the highest voltages employed, though not as markedly as the OH- transport. Using the procedure previously outlined, the " unenhanced " C1- transport can be calculated; the results are shown in fig.4. The unenhanced C1- transport does show a plateau at about 390 A m-2. This value is compared with the Cowan break- I I I / I I I 5 10 15 A(PP FIG. 4.-E!nhanced, or actual, C1- current density (0). Unenhanced, or corrected, C1- current density( 0). Conditions as for fig. 3. point value and with the Levich value, eqn (5), in table 1, which also shows the Cowan and Levich values obtained at a series of rotation speeds. Fig. 5 shows the Cowan values from table 1, together with a similar set for Na+ transport into a cation membrane, compared with the Levich plots, using eqn (5). It can be seen that these approaches correlate well, and that the Cowan breakpoints (which are comparatively easy to obtain) can be used to estimate boundary layer thicknesses, 6, in electro- dialysis stacks.TABLE l.-COMPARISON OF THE COWAN BREAKPOINT CURRENT DENSITY WITH THE LEVICH AND " UNENHANCED " LIMITING CURRENT DENSITIES 0.0174 mol dm-j NaCl SOLUTION AT 19°C ANION-EXCHANGER DISC disc rotation speed/r.p.m. 175 250 350 740 1200 1290 Cowan breakpoint current density/A m-' 145 175 200 270 350 370 Levich ilh/A m-' eqn (5) 134 161 191 277 353 366 - 390 " unenhanced " plateau - - - - Results of the other systems are given in detail by Makai.* The following conclusions can be stated : (i) for 0.0098 mol dm-3 Na,SO, at 1300 r.p.m., anion exchange membrane ; results very similar to fig. 3, giving an " unenhanced " SO;- plateau of 380Anr2, which can be compared with 350Am-2 from eqn (4), and 330 A m-2 from the Cowan breakpoint. (ii) For 0.0174 mol dm-3 NaCl at 1 300 r.p.m., cation exchange membrane. Results different in nature from those in fig.3.2356 POLARISATION IN ELECTRODIALYSIS The Hf ion transport from water-splitting is much smaller than OH- transport through the anion membrane. The Na+ transport continues to increase with increasing A@, but at a much greater rate than was the case with Cl- transport. The “ unenhanced ” Na+ transport does not show a plateau, but increases for the whole range of A@ covered. Cowan plots for a range of rotation speeds do show break- points, which correlate fairly well with eqn (5), see fig. 5. (iii) For 0.0092 mol dm-3 CaCl, at 1300 r.p.m., cation exchange membrane. Results very similar to (ii).Again low H+ transport and no “unenhanced” plateau. (iv) For mixed NaCl, Na,SO,, each at about 0.01 kg equivalents m-3 at 1300 r pm., anion exchange membrane. Results similar to the anion transport of C1- or SOi- individually. The 400 800 1200 10 ‘ 200 disc speed1r.p.m. FIG. 5-Dependence of Cowan limiting currents on disc rotation. 0 , Na+ transport through cation membrane ; 0, C1- transport through anion membrane. The lines are Levich plots, eqn (5). procedure for calculating ‘‘ unenhanced ” transports is not applicable for mixtures. (v) For mixed NaCl, CaCI,, each at about 0.01 kg equivalents m-3 at 1280r.p.m., cation exchange membrane. Results similar to the cation transport of Na+ or Ca2+ individually, including the anomalously high increasing transports with increasing voltage.DISCUSSION For transport through the anion exchange membrane a consistent picture can be drawn, in which the Levich treatment, the calculation procedure outlined here and the empirical Cowan plot approach are all in good agreement with each other. The high OH- transport produced at high A@ is undesirable in practice, but at least the theory and attendant calculations seem to be able to cope with the experimental data. For transport through the cation exchange membrane the situation is reversed. The Hf transport produced at high A@ is much lower (a desirable thing in practice), but the theory cannot explain the experimental data. The small H+ transport which did arise in our experiments came almost entirely from the bulk-phase H+ ion concentra- tion resulting from water-splitting at the anion-exchange membrane between the compartments, and not from water-splitting at the rotating cation-exchange membrane disc.Tests with alkaline bulk solutions showed this to be the case, the H+ transport through the disc being almost undetectable in that case. The real problem is how the anomalously high Naf, or Ca2+, transports can be explained. These high transports also occur in practical electrodialysis stacks, see Makai, * and Spiegler et al. There was the possibility that the fluid mechanics in theA . J . MAKAI AND J . C. R. TURNER 2857 cell compartments could be affected by the transport. But if so, why not at the anion membrane also ? The well-defined forced-convection flow field in a rotating-disc apparatus would seem to be able to dominate over free convection effects, which was a reason for these experiments.Nevertheless, the high cation transports remain and we can offer no explanation. From the results for mixed solutions it is possible to define a '' preference factor ", which is the ratio of the proportion of ions transferred to the proportion of ions in the bulk solution. As A@ is increased, the preferential removal of divalent ions (a desirable characteristic in the electrodialysis of brackish water) is reduced. It is easy to see that at high A@ the preference factor should fall towards unity, which is a subsidiary disadvantage of high-current-density electrodialysis. We would like to thank Dr. J. N. Agar for much helpful discussion, and J. M. Sturton for assistance with the atomic absorption analyses. A. J. M. is grateful to the Shell Co. of Australia for a Postgraduate Scholarship which supported this work. J. F. Brady and J. C. R. Turner, J.C.S. Farahy I, 1978,74,2839. W. J. Albery, Electrode Kinetics (Clarendon Press, Oxford, 1975). V. G. Levich, Physicochemical Hydrodynamics (Prentice-Mall, N.J., 1962), sect. 11, p. 60. D. Gregory and A. C. Riddiford, J. Chem. Soc., 1956,3756. V. G. Levich, Physicochemical Hydrodynamics (Prentice-Hall, N.J., 1962), sect. 52, p. 293. W. J. Albery, Trans. Faraday Soc., 1969, 61,2063. ' D. A. Cowan and J. H. Brown, Ind, and Eng. Chem., 1959,51,1455. * A. J. Makai, Ph.D. Thesis (Cambridge University, 1977). K. S. Spiegler, J. Sinkovic and 3. Leibovitz, Desalination Report No. 62, University of Cali- fornia, July 1975, p. 19. (PAPER 81265)

ISSN:0300-9599    

DOI:10.1039/F19787402850

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Calculations on Ionic Solvation Part 2.-Entropies of Solvation of Gaseous Univalent Ions using a One-layer Continuum Model BY MICHAEL H. ABRAHAM" Department of Chemistry, University of Surrey, Guildford, Surrey JANOS LISZI Department of Physical Chemistry, University of VeszprCm, 8201 Veszprkm, Hungary AND Received 17th February, 1978 The electrostatic entropy of solvation of gaseous ions has been calculated using our previous model in which the ion is surrounded by a local solvent layer, immersed in the bulk solvent. The calculate6 electrostatic entropy is combined with the nonelectrostatic entropy of solvation, obtained from experimental data on entropies of solution of gaseous nonpolar solutes, to yield the total entropy of solvation of a gaseous ion. The only new parameters involved in the calculations are the variations with temperature of the solvent bulk dielectric constant (a known property) and the dielectric constant in the local solvent layer.We found that if the latter parameter is taken as - 0.001 60 (a reasonable value for a region of low dielectric constant), there is excellent agreement with experiment for entropies of solvation of univalent cations and anions in a wide variety of aprotic solvents, and for entropies of transfer of these ions between aprotic solvents. Since no adjustable parameters are used in the calculations, the method can be used to predict entropies of solvation or of transfer in aprotic solvents. Agreement with experiment is no+ 'xnd for solvation entropies of ions in hydrogen bonded solvents.We have reported previously results of calculations on the free energy of solvation of gaseous univalent ions, using a continuum model in which an ion of radius a and dielectric constant ci is surrounded by a local solvent layer of thickness (b-a) and dielectric constant For all ions in all solvents we took a as the ionic crystal radius, (b -a) as the solvent radius, E~ = 1, c1 = 2 (since for many solvents the value of E: is approximately 2) and c0 as the bulk dielectric constant. We then calculated the electrostatic free energy of solvation by the method of Beveridge and Schnuelle and combined this with the nonelectrostatic free energy, obtained from experimental data on nonpolar solutes, to obtain the total free energy of solvation via eqn (1) Agreement between calculated and observed free energies of solvation was so good as to encourage us to carry out similar calculations on ionic entropies of solvation.Our aim was first to see if the simple one-layer model could be used to reproduce known entropies of solvation of cations and anions in a variety of solvents; several workers have suggested that for ions in water and highly structured solvents a second, disorganised, solvent layer will make an additional contribution to the total entropy. Secondly, we hoped to be able to predict entropies of solvation and immersed in the bulk solvent of dielectric constant c0. AG," = AG,"+AG,". (1) 2858M. H . ABRAHAM AND J. LISZI 2859 entropies of transfer of ions, especially for the less polar solvents where there are very great experimental difficulties in obtaining such data.Few calculations on entropies of solvation of ions in nonaqueous solvents have been reported. Although Eley and Evans successfully calculated entropies of solvation of ions in water using a discontinuous method, a similar calculation failed completely to account for ionic entropies of solvation in methanol. More recently, de Ligny et aL7 have carried out calculations on the division of entropies of transfer of electrolytes into single ion contributions, but to date calculations on actual entropies of solvation in nonaqueous solvents have met with but little success. THEORY OF THE METHOD In terms of the one-layer model, the electrostatic part of the ionic solvation equation with entropy is obtained by differentiating Beveridge and Schnuelle's respect to temperature.Taking only the relevant first term, we derive eqn (2) in which 2 is the charge on the ion (taken as unity in all the present calculations), 6c1/6T is the variation of local dielectric constant with temperature, and 6co/6T is the like variation of the bulk dielectric constant. We retain exactly the same values of a, b, c1 = 2 and c0 as used in our free energy calculations,' so that the only para- meters left are 6c1/6Tand 6c0/6T. The latter is a property of the bulk solvent and the values we have used are given in table 1. Thus the only variable parameter that remains is 6cl/6T, and even this is constrained because for all solvents we have fixed cl = 2. For many solvents with c0 N 2, values of 6c0/6T are about -0.001 60; furthermore, although data are rather limited, it seems as though for most aprotic solvents values of 6&2/ST also approach the same value.We, therefore, set for all solvents Sc,/6T = -0.801 60, and thus retain no variable parameter at all. This procedure has the advantage of rendering the method entirely predictive; the dis- advantage is that the method may well not apply to hydrogen bonded solvents where ~EJBT is usually numerically smaller than 0.001 60. We now combine the calculated AS," values with a nonelectrostatic contribution, AS,", obtained from experimental data *-12 on the entropy of solution of gaseous nonpolar solutes (table 2), via eqn (3) AS: = AS,O+AS,O. (3) As with the corresponding free energies of solvation, this procedure removes all difficulties over standard states.There is some uncertainty in the experimental values given in table 2, and within this uncertainty, values for solution of nonpolar solutes in many nonaqueous solvents can be fitted to the linear eqn (4)" AS: = mr+c (4) where nz = 6.96, Y = solute radius in A, and c = 1.2 cal K-l mol-l, (see table 3). For solvents other than water, the AS," values are quite small (e.g., -8.5 cal K-I mol-1 for Naf and - 16.5 cal K-l mol-' for I- in many solvents)? and to a large * Data we have used to calculate m and c in eqn (4) cover only the range r = 1.29 (He) to 2.03 (Xe). Recent work l4 on the entropy of solution of tetramethyltin, with r = 3.07, indicates that for many nonaqueous solvents eqn (4) does in fact hold up to at least r = 3.07 A.t 1 cal = 4.184 J.2860 CALCULATIONS ON IONIC SOLVATION extent reflect the change in standard state from 1 atm (gas) to unit mol fraction (soln) ; for most nonaqueous solvents such a change corresponds to an entropic contribution of about - 12,cal K-l mol-l. For water,12 the AS," values are much more negative. TABLE 1.-vALUES OF -6&,/6T USED IN THE CALCULATIONS solvent - 68016T water methanol ethanol 1-propanol formamide N-methylformamide N,N-dime thy1 formamide dimethylsulphoxide acetoni trile nitromethane acetone nitro benzene ammonia 1 ,Zdichloroethane 1 , 1 -dichloroethane tetrahydro furan 1 ,Zdimethoxyethane ethyl acetate chlorobenzene bromo benzene pentyl acetate diethyl ether di-isopropyl ether butyl stearate di-isopen t yl ether benzene cyclohexane 0.3595 * 0.197 b* C 0.147 0.142 0.72 (0.40) 1.62 0.178 d 0.106 (0.126)f 0.160 C 0.161 0.0967 (0.0977) C 0.180 0.0780 g 0.0560 C 0.0480 0.0299 (0.0298) 1 0.0410 0.015 0.0168 C 0.0143 0.012 0.020 0.018 0.0053 C 0.0050 0.001 99 0.001 5 5 ' a R.L. Kay, G. A. Vidulich and K. S. Pribadi, J. Phys. Chem., 1969,73,445 ; B. B. Owen, R. C. Miller, C. E. Milner and H. L. Cogan, J. Phys. Chem., 1961, 65, 2065. b From data quoted in Landolt-Bornstein, Zaklenwerte und Funktionen, Sechste Auflage, Band 11, Teil6, and by J. Timmer- mans, Physico-chemical Constants of Pure Organic Compounds (Elsevier, Amsterdam, 1950 and 1965). C A. A. Maryott and E. R. Smith, Table of Dielectric Constants ofPure Liquids, NBS Circular 514, Washington D.C., 1951. d G. R. Leader and J.F. Gormley, J. Amer. Chem. Soc., 1951, 73, 5731. e R. Garney and J. E. Prue, Trans. Faradar Soc., 1968, 64, 1206. f J. F. Casteel and P. G. Sears, J. Chem. Eng. Data, 1974,19,196. 0 H. M. Grubb, J. F. Chittum and H. Hunt, J. Amer. Chem. Soc., 1936,58,776. h J. T. Denison and J. B. Ramsey, J. Amer. Chem. Soc., 1955,77,2615. i C. Carvajal, K. J. Tolle, J. Smid and M. Szwarc, J. Arner. Chem. Soc., 1965,87,5548. f D. J. Metz and A. Glines, J. Phys. Chem., 1967, 71, 1158. k V. Viti and P. Zampetti, Chem. Phys., 1973, 2, 233. * R. H . Stokes, J. Chem. Thermodynamics, 1973,5, 379. Experimental ionic solvation entropies can be obtained from known entropies of hydration,13 together with ionic entropies of transfer from water to the given solvents. In order that the same standard states apply to both nonpolar solutes and ions, we recalculated the entropies of hydration tabulated by Noyes l3 to standard states of 1 atm (gas) and unit mol fraction (as).There is still the problem that the assignment of cationic and anionic contributions carried out by Noyes l3 may not be appropriate or consistent with the constants we use in the present calculations. We find, however,M. H. ABRAHAM AND J . LISZI 2861 that it comparatively small adjustment (by 5 cal K-l mol-l)* to Noyes’ cationic and anionic values yields single ion entropies of hydration that are perfectly compatible with the present calculations for all solvents, provided that single ion entropies of transfer are assigned by the correspondence plot method. We used ionic entropies of transfer, given by Abraham and c o - ~ o r k e r s , ~ ~ - ~ ~ that are based on the mol fraction TABLE 2.-sTANDARD ENTROPIES OF SOLUTION OF GASES (1 atm GAS AND UNIT MOL FRACTION SOLUTION) IN cal K-l mol-1 AT 298 K a gas : solvent hexane cyclohexane benzene toluene iodobenzene bromobenzene chloro benzene nit robenzene acetone average value calc.from eqn (4) b N-met hylacetamide isobutanol ethanol methanol water -AS,” He Ne Ar Kr Xe 10.0 9.9 10.6 10.0 9.4 10.4 12.3 9.4 8.6 10.1 10.2 11.2 12.3 8.3 11.1 11.1 12.1 10.6 8.5 9.9 10.6 10.9 14.1 13.7 13.0 12.7 12.9 13.7 13.7 13.4 13.3 13.4 13.1 13.6 13.4 13.3 14.3 14.1 14.6 14.2 13.3 13.9 13.8 13.7 12.4 15.0 14.3 14.1 15.4 16.8 13.0 13.4 15.5 16.3 14.6 14.9 17.3 17.4 24.1 26.3 30.9 32.3 a From data in ref.(8)-(12) ; b -AS,” = 6.96r+ 1.2. 14.5 15.0 15.0 15.1 15.7 14.3 14.9 15.3 34.2 TABLE 3.-cONSTANTS IN THE EQUATION - AS,”(cal K-l mol-l) = mr+ c solvent m C aprotic solvents in table 2 6.96 1.2 N-met h ylace t amide 6.96 3.2 isobutanol 6.96 4.0 ethanol 6.96 4.0 methanol 6.96 5.3 water b 10.51 12.84 a With AS: in cal K-I mol-l and Y in A. b Using data for Ar, Kr, Xe, Ch, C&, C3Ha, n-C4H1,, and neo-CsH12. standard state, separated into cationic and anionic contributions by the correspon- dence plot method, and smoothed out by the method developed by Abraham.15 Details of the final solvation entropies are in table 4. Absolute solvation entropies of ions have also been tabulated by Criss and Salomon.l * Taking into account the change in standard state, there is quite good agreement with the values we give in * This adjustment leads to ASg(H+) = -42.3 cal K-I mol-1 with our standard states.Given that F(H+, gas, 1 atm) = 26.0 this yields for $(H+, aq) a value of - 16.3 on the mol fraction scale or - 8.3 cal K-l mol-’ on the usual molal scale. The latter value is quite close to suggested ‘‘ abso- lute ” values on the molal scale ( - 5 to -6 cal K-l mol-1).18 Noyes original single ion division corresponds to a value of -3.3 cal K-l mol-l for the “absolute ” value of*(€€+, aq, I mol kg-l).2862 CALCULATIONS ON IONIC SOLVATION table 4. We also give in table 4 ionic solvation entropies obtained by Strong and Tuttle for the solvents tetrahydrofuran (THF) and 1,2-dimethoxyethane (1'2- DME). Unfortunately there seem to be rather large uncertainties in the data that preclude any exact application of the correspondence plot method.Na+ K+ Rb+ cs+ Me4Nf Et4N+ c1- Br- I- c l o y TABLE 4.-ENTROPIES OF SOLVATION OF IONS (1 atm GAS + UNIT MOL FRACTION SOLUTION) IN Cal K-' M01-l AT 298 K waterb MeOHc EtOHC h"OHC F NMF DMSO DMF MeCN -36.8 -47 -50 -54 -42 -45 -50 -57 -55 -28.7 -42 -45 -49 -37 -40 -45 -52 -50 -25.9 -38 -41 -45 -33 -36 -41 -48 -46 -25.2 -36 -39 -43 -31 -334 -39 -46 -44 -39.8 -35 -38 -42 -30 -33 -338 -45 -43 -59.1 -43 -42 -46 -34 -37 -42 -49 -47 -23.3 -33 -36 -4-0 -28 -31 -36 -43 -41 -19.3 -31 -34 -38 -26 -29 -34 -41 -39 -14.3 -27 -30 -34 -22 -25 -30 -37 -35 -13.0 -26 -29 -33 -21 -24 -29 -36 -34 acetone - 60 - 55 - 51 - 49 - 48 - 52 -46 -44 - 40 - 39 ammonia 1,2-DCE 1,l-DCE d THF * 1,2-DME Na+ - 60 - 69 - 70 - 82 - 58 K+ - 55 -64 - 65 - 67 - 74 Rb+ - 51 -60 - 61 -70 -104 cs+ - 49 - 58 - 59 - 89 - 88 Me4N+ - 48 - 57 - 58 Et,N+ - 52 - 61 - 62 c1- - 46 - 55 - 56 - 63 - 73 BI- -44 - 53 - 54 I- -40 - 49 - 50 ClO; - 39 - 48 - 49 Obtained from the values for water, together with the smoothed entropies of transfer from the correspondence plot method in ref.(15). b From ref. (13) adjusted as stated in the text. C Ref. (17). d Ref. (16). e From ref. (19) ; the cationic and anoinic contributions have been roughly assigned by the correspondence plot method. There is also a difficulty over AS," for Me4N+ and Et4N+. Quite discordant values have been calculated for S0(Me4N+, gas) : 73.7 cal K-l rnol-1 by Boyd 2o and 58.2 by Ladd.21 Combination with Johnson and Martin's value 22 for S"(Me,N+, aq, 1 mol kg-l) yields on our scale values for ASi(Me4N+) of -39.8 (Boyd) or -24.1 (Ladd).We chose the value based on Boyd's calculation on the grounds that Ladd's value for So(Me4N+, gas) seems too low by comparison with that for neo- pentane (73.2 cal K-l r n ~ l - l ) . ~ ~ For Et4N+ we used data by Boyd 2o and Johnson and Martin,22 but there is no independent value to confirm these calculations. RESULTS AND DISCUSSION We give in table 5 details of the calculations for 1,l-dichloroethane (1'1-DCE). Agreement between calculated and observed AS," values is in general excellent, except for Me,N+ and Et4N+ where the calculated values are too positive by 10 and 15 cal K-l mol-1 respectively. These differences are not special to 1,l -dichloroethane, and we find for a11 the solvents studied that on average the calculated AS," values areM. H.ABRAHAM AND J . LISZI 2863 too positive by 8 and 12 cal K-I mol-l. This cannot be due to incorrect values of the solute radii, a, because wide variations in a still do not yield agreement. We think it possible (see discussion above) that the observed values of AS,” are in error ; this would then result in exactly the same error in AS,” for all the solvents studied. In the event we have incorporated in all the calculated AS,” values, extra entropic contributions of - 8 and - 12 cal K-l mol-1 for Me4N+ and Et4N+, respectively. TABLE 5.-RESULTS FOR SOLVENT. 1,I-DICHLOROETHANE IN Cal K-’ mOl-’ AT 298 K ion ASZ(ca1c) AS: a ASi(ca1c) ASi(obs) Na+ - 67.3 - 8.5 - 75.8 - 70 K+ - 53.7 - 10.5 - 64.2 - 65 Rb+ -50.1 -11.1 - 61.2 - 61 a+ -43.5 - 12.7 - 56.2 - 59 Me4N+ -28.6 - 19.1 - 47.7 (- 55.7) - 58 Et4N+ - 24.0 - 22.8 -46.8 (-58.8) - 62 c1- - 40.1 - 13.8 - 53.9 Br- - 37.3 - 14.8 - 52.1 I- - 33.3 - 16.5 -49.8 c104 - 30.1 - 18.2 -48.3 - 56 - 54 - 50 - 49 Q -AS: = 6.96r-1- 1.2.b Extra entropic contributions included (see text). TABLE 6.xALCULATED AND OBSERVED VALUES OF IONIC SOLVATION ENTROPES, As,” IN i on Na+ Kf Rb+ cs+ Me4N+ EtdN+ c1- Br- I- Clog ion Na+ K’ Rbf Csf Me4N+ Et,N’ c1- Br- I- ClO, d K-’ rno1-l AT 298 K 1, 1-DCE 1,2-DCE THF 1,2-DME ammonia calc. obs. calc. obs. calc. obs. calc. obs. calc. obs. -75.8 -70 -78.0 -69 -78.4 -82 -87.7 -58 -64.3 -60 -64.2 -65 -66.2 -64 -66.6 -67 -75.3 -74 -53.7 -55 -61.2 -61 -63.2 -60 -63.6 -70 -72.1 -104 -51.0 -51 -56.2 -59 -58.0 -58 -58.4 -89 -66.5 -88 -46.6 -49 -55.7 -58 -57.2 -57 -57.5 - 64.3 -48.0 -48 -58.8 -62 -60.2 -61 -60.5 - 66.6 -51.9 -52 -53.9 -56 -55.6 -55 -56.0 -63 -63.8 -73 - 4 .6 -46 -52.1 -54 -53.8 -53 -54.2 - 61.8 -43.2 -44 -49.8 -50 -51.4 -49 -51.8 - 59.0 -41.4 -40 -48.3 -49 -49.8 -48 -50.1 - 57.0 -40.3 -39 acetone acetonitrile DMF 2MSO calc. obs. calc. obs. calc. obs. calc. obs. -63.8 -60 -57.7 -55 -59.3 -57 -55.1 -50 -53.0 -55 -47.5 -50 -48.9 -52 -45.0 -45 -50.3 -51 -45.0 -46 -46.3 -48 -42.5 -41 -45.9 -49 -40.9 -44 -42.1 -46 -38.5 -39 -47.3 -48 -43.3 -43 -44.1 -45 -41.3 -38 -51.2 -52 -47.6 -47 -48.3 -49 -45.7 -42 -43.9 -46 -39.1 -41 -40.2 -43 -36.8 -36 -42.5 -44 -37.9 -39 -38.9 -41 -35.6 -34 -40.7 -40 -36.3 -35 -37.3 -37 -34.2 -30 -39.6 -39 -35.5 -34 -36.4 -36 -33.4 -29 a Calculated values using -AS: = 6.96rf 1.2.All values for Me4N+ and Et4N+ include the extra entropic contributions (see text). Observed values from table 4. b If - 8eo/8T is taken as 0.126 (see table l), the calculated values would be more negative by only about 0.3 cal K-’ mol-‘.2864 CALCULATIONS O N IONIC SOLVATION In table 6 are summarised results of our calculations on the non-hydrogen bonded solvents for which data are available. As before,l we define two parameters as measures of agreement between calculated and observed AS," values : the standard deviation 0 = { lAS,"(calc) - AS:(obs)J2/(n - 1)}&, and a measure of systematic devia- tion = AS,"(calc, average) -AS,"(obs, average).For the solvents l,l-DCE, 172-DCE, ammonia, acetone, acetonitrile, dimethylformamide (DMF) and dimethyl- sulphoxide (DMSO) there is good agreement between calculated and observed values, with CJ averaging about 2 cal K-l mol-1 and 5 averaging numerically only about 1 cal K-1 mo1-I for the 7 solvents (see table 7). These deviations are of the order of experimental errors in AS," values, and it seems therefore that we have achieved one of our main aims, namely to reproduce observed entropies of solvation, especially in the less polar solvents, by a method that is predictive in nature. The two aprotic solvents for which there is only poor agreement are THF and DME (table 7), but it TABLE VA VALUES OF 5 AND Q JN cal K-' rnoP solvent no. of ions 5 d 1, 1-dichloroethane 1 , 2-dichloroethane tetrahydrofuran 1,2-dirnethoxyethane ammonia acetone acetoni trile dimet hyl formamide dimethylsulphoxide 10 10 5 5 10 30 10 10 10 0.8 9.6 6.3 -0.1 0.6 0.3 1.3 - 1.9 - 2.4 2.7 3.4 16.1 24.7 1.9 2.0 1.9 2.2 3.2 N-methylformamide 10 - 8.9 9.6 formamide 10 - 13.5 14.4 1 -propano1 C 10 - 12.4 13.2 ethanol 10 - 12.5 13.3 methanol d 10 - 13.4 14.3 water 10 - 26.8 29.6 0 -AS: = 6.96~+3.2.b -AS: = 6.96r+5.3 ; if - Sq,/ST was taken as 0.40 (see table l), 5 would be about - 12.5 cal K-' mol-'. C -AS: = 6.96~+4.0. d -AS: = 6.96rf5.3. AS: ob- tained from a plot of AS: against Y (inert gas) for Na+, K+, Rbf and Cs+ and from -AS: = 10.51r + 12.84 for the larger ions. must be borne in mind that for these rather nonpolar solvents there is considerable technical difficulty in obtaining experimental values and that this is reflected in possible errors.Attempts to apply the correspondence plot method to experimental values for THF and DME suggests that they are self-consistent to only about 10 cal K-I mol-l (THF) and 20 cal K-I mol-1 (DME). Strong and Tuttle I9 themselves estimate that errors in enthalpies of solution could be " several kcal mol-I ", so that it is not surprising that there is only poor agreement between calculated and observed entropy values. Indeed, we might claim that for solvents as nonpolar as THF and DME, with dielectric constants 7.4 and 7.3, our method of calculation yields entropies of solvation that are more reliable than those that can be obtained by present experi- mental methods. Since values of 5 and CT (table 7) are so small for the solvents that are not hydrogen bonded, it seemed that we might also achieve the aim of calculation of reliable ionic entropies of transfer.Details are in table 8 for the typical case of (Rbf and Br-), with acetonitrile taken as an arbitrary reference solvent. Agreement with observedM, H. ABRAHAM AND J . LISZI 2865 values is quite remarkable, and the present method clearly can be used for the calcula- tion of entropies of transfer between non-hydrogen bonded solvents to within present experimental error (usually about 1-2 cal K-l mol-l). For hydrogen bonded solvents, we found substantial differences between calcu- lated and observed AS," values. We do not give details for these solvents, but list in table 7 the [ and 0 values found.Since the arithmetical value of [ is approximately the same as 6, the differences are systematic and not merely random. One cause TABLE CALCULATION OF ENTROPIES OF TRANSFER ON THE MOL FRACTION SCALE, M cal K-l mol-1 AT 298 K solvent Rb+(calc) Br-(cab) Rb+(obs) Br-(obs) a dimethylsulphoxide + 3 +2 +5 + 5 dimethylformamide - 1 - 1 -2 -2 acet onitrile 0 0 0 0 acetone - 5 - 5 - 5 - 5 ammonia -6 -5 - 5 -5 1,2-dichloroethane - 18 - 16 - 14 - 14 1, 1-dichloroethane - 16 - 14 - 15 - 15 a Table 4. of the discrepancies is that for this type of solvent -6~,/6T is probably less than 0.001 60, the value used in all calculations. This cannot be the only cause because choice of other values for 6~,/6T still does not yield agreement. We hope to investi- gate the hydrogen bonded solvents in more detail later.We also thought it useful to investigate for the aprotic solvents the effect of altering the thickness of the local layer, (b -a), whilst keeping a constant value of Q. It is not obvious how changes in b will influence the calculated AS," (and AS,") values, because both terms in eqn (2) include the b parameter. We find that an increase in b TABLE CA EFFECT OF THE THICKNESS OF THE LOCAL LAYER (b-a) ON CALCULATED ENTROPIES AND FREE ENERGIES OF SOLVATION OF Rb+ AT 298 K a solvent (b-a)lA AS,"(calc) AS,"(obs) AGXcalc) AG:(obs) 2.0 4.444 MeCN { 2,222 4.970 C 5.200 -44.1 - 45.0 - 49.7 - 49.3 - 50.3 - 53.1 -61.9 - 61.2 - 59.8 - 73.0 - 71.9 - 75.3 - 46 - 73.9 - 64.3 - 74.9 - 51 - 72.2 - 64.7 -72.1 - 61 - 69.3 - 62.8 - 68.4 I a Free energies (in kcal mol-') calculated as in ref.(1) ; entropies in cal K-l mol-I. In all calculations a = 1.43 A. b Solvent radius. C Solvent diameter. results in a more negative entropy of solvation in the case of polar solvents, but in a more positive calculated entropy of solvation for the less polar solvents ; results for acetonitrile, acetone and 1, 1-DCE are in table 9 where Q has been taken as 1.43 A (Rb+). Small variations in (b-a) produce only minor changes in AS," (calc), so that it requires a quite detailed analysis to show that choice of (b-a) as the solvent radius leads to better agreement with observation than choice of (b-a) = 2 A for all solvents. It does seem, however, that with the values of and ael/aTused in our2866 CALCULATIONS ON IONIC SOLVATION calculation, the choice of solvent radius does lead to slightly better agreement in terms of entropies of solvation.The free energy term is affected more by the change in (b -a) than is the entropy term (table 9), and in this case also the choice of solvent radius leads to better agreement than does that of taking (b -a) = 2 A. We also give in table 9 results of calculation with (b-a) taken as the solvent diameter, but there is now considerable discrepancy between calculated and observed values, especially in the free energy term. Many previous workers have modified the simple Born equation for the eiectro- static free energy or enthalpy of solvation of an ion, by means of empirical adjust- ments, A, to the ionic radii. For cations, these adjustments are often ~ 0 .8 A not only in water 24 but also in propylene carbonate 2 5 and DMF,26 but for anions the adjustments range from 0.1 A in water 24 to 1.00 A in DMF.26 These empirical, non-predictive, modifications to the simple Born equation are in no way equivalent to the present one-layer model. In the modified Born equation a charged region of radius (a+A) and E = 1 is surrounded by the bulk solvent with E = E ~ , whereas in the present model a charged region of radius a and E = 1 is surrounded by the solvent layer of thickness (b -a) and E = 2, followed by the bulk solvent with E = co. Our calculations indicate that this solvent layer must be of about a solvent radius thick; for most solvents this corresponds to 2-3 A. CONCLUSION Ionic entropies of solvation and ionic entropies of transfer can be calculated accurately by use of a simple one-layer model in which the dielectric constant of the solvent layer is taken as = 2 and in which &/8T is taken as -0.001 60, provided that the solvents investigated are non-hydrogen bonded.The method uses no adjustable constants and is predictive in nature. As an example, entropies of solva- TABLE 1o.-PREDICTION OF ENTROPIES OF SOLVATION IN Cal K-' m01-l AT 298 K " solvent Rb+(calc) Br-(calc) Rb+(obs) Br-(obs) dimethylsulphoxide ni tromethane dimet hylformamide acetonitrile nitro benzene acetone ammonia 1 ,2-dichloroethane 1,l-dichloroethane tetrahydro fur an 1 ,Zdimet hoxyet hane ethyl acetate chlorobenzene bromobenzene pentyl acetate diethyl ether di-isopropyl ether butyl stearate di-isopentyl ether benzene cyclohexane - 42 - 45 - 46 - 45 - 48 - 50 - 51 - 63 - 61 - 64 - 72 - 58 - 63 - 61 - 63 - 86 - 87 - 61 - 66 - 57 - 57 - 36 - 38 - 39 - 38 -40 - 43 - 43 - 54 - 52 - 54 - 62 - 49 - 53 - 52 - 54 - 74 - 75 - 52 - 56 - 48 - 48 =Values of AS: taken as -11.1 for Rb+ and -14.8 for Br- -AS: = 6.96r+ 1.2.b Table 4. - 41 - 34 - 48 - 41 - 46 - 39 - 51 -44 - 51 -44 - 60 - 53 - 61 - 54 - - - - in all cases, as calculated fromM . H . ABRAHAM AND J . LISZI 2867 tion for Rbf and Br- are predicted in table 10 for a very wide range of aprotic solvents? The only solvent properties required for such predictions are the solvent density (used to obtain the thickness of the solvent layer) and the bulk properties co and SEO/~T. Even for the aprotic solvents of some degree of solvent structure, such as DMSO and DMF, a one-layer model is sufficient to account for the ionic entropies of solvation.This result is slightly at variance with the suggestions of Parker and co-workers that entropic contributions from a second (disorganised) solvent layer are negligible in the case of acetonitrile (as we would also suggest) but are not negligible for DMSO and DMF. Parker and his coworkers stress, however, that any contributions from this second layer in the case of DMSO and DMF are much less than with highly structured hydrogen bonded solvents. M. H. Abraham and J. Liszi, J.C.S. Furuduy I, 1978,74,1604. D. L. Beveridge and G. W. Schnuelle, J. Phys. Chem., 1975,79,2562. H. S . Frank and M. W. Evans, J. Chem.Phys., 1945,13,507. B. G. Cox, G. R. Hedwig, A. J. Parker and D. W. Watts, Austral. J. Chem., 1974, 27, 477. D. D. Eley and M. G. Evans, Trans. Furaduy SOC., 1938,34,1093. D. D. Eley and D. C. Pepper, Trans. Furuduy SOC., 1941,37,581. ' D. Bax, C. L. de Ligny and M. Alfenaar, Rec. Trav. chim., 1972, 91,452; D. Bax, C. L. de Ligny and A. G. Remijnse, Rec. Truv. chim., 1972,91,965. * H. L. Clever, R. Battino, J. H. Saylor and P. M. Gross, J. Phys. Chem., 1957, 61, 1078; R. Battino, F. D. Evans, W. F. Danforth and E. Wilhelm, J. Chem. Thermodynamics, 1971, 3, 743; J. H. Saylor and R. Battino, J. Phys. Chem., 1958,62,1334. R. H. Wood and D. E. DeLaney, J. Phys. Chem., 1968,72,4651. (University of New Zealand, 1949). mentuls,) 1976, 15, 336. lo A. J. Beckwith, MSc. Thesis (University of New Zealand, 1949) ; J. T. Law, PhD. Thesis l1 C. L. de Ligny, N. G. van der Veen and J. C. van Houwelingen, Ind. and Eng. Chem. (Fundu- l2 E. Wilhelm, R. Battino and R. J. Wilcock, Chem. Rev., 1977,77,219. l3 R. M. Noyes, J. Amer. Chem. SOC., 1962,84, 513. l4 M. H. Abraham and A. Nasehzadeh, unpublished observations. l5 M. H. Abraham, J.C.S. Furuduy I, 1973,69, 1375. l 6 M. H. Abraham, A. F. Danil de Namor and R. A. Schulz, J. Solution Chem., 1976,5, 529. l7 M. H. Abraham and A. F. Danil de Namor, J.C.S. Faraday I, 1978, 74, 2101. l 8 C. M. Criss and M. Salomon, in Physical Chemistry of Organic Solvent Systems, ed. A. K. l9 J. Strong and T. R. Tuttle, J. Phys. Chem., 1973, 77, 533. 2o R. H. Boyd, J. Chem. Phys., 1969,51,1470. '' M . F. C. Ladd, 2. phys. Chem. (N.F.), 1970,72,91. 22 D. A. Johnson and J. F. Martin, J.C.S. DuZton, 1973, 1583. 23 D. R. Stull, E. F. Westrum and G. C. Sinke, Chemical Thermodyriumics oforgatzic Compounds 24 W. M. Latimer, K. S. Pitzer and C. M. Slansky, J. Chem. Phys., 1939,7, 108. 25 Y.-C. Wu and H. L. Friedman, J. Phys. Chem., 1966,70,501. 26 C . M. Criss and E. Luksha, J. Phys. Chem., 1968, 72,2966. * An interesting prediction (see table 10) is that entropies of solvation pass through minimum Covington and T. Dickinson (Plenum Press, London, 1973). (John Wiley, New York, 1969). values ( i z , more negative values) in solvents with E~ N 4. (PAPER 8/280)

ISSN:0300-9599    

DOI:10.1039/F19787402858

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Meats of Hydrogenation of Large Molecules Part 3.-Five Simple Unsaturated Triglycerides (Triacylglycerols) BY DONALD w. ROGERS* AND DEBENDRA N. CHOUDHURY Department of Chemistry, The Brooklyn Center, Long Island University, Brooklyn, New York 11201, U.S.A. Received 27th February, 1978 We have determined the enthalpies of hydrogenation of the simple triglycerides of oleic, elaidic, petroselinic, linoleic and linolenic acids. These compounds include the largest molecules ever studied by hydrogen calorimetry and one of them (trilinolenin) produces the largest heat of hydrogena- tion yet observed. Heats of hydrogenation per double bond for the five triglycerides reported here are remarkably consistent with each other and with those of the free fatty acids and methyl esters previously studied.These results suggest a molecular conformation in hexane solution in which the unsaturated acid residues are well separated and the double bonds are thermochemically independent. In Parts 1 and 2 of this series,'. we found that the heats of hydrogenation per double bond of several unsaturated and polyunsaturated fatty acids and their methyl esters are remarkably self consistent. Cis acids and esters have a heat of hydrogena- tion of z - 125 kJ mol-l regardless of whether they are mono-, di- or tri-unsaturated while the trans acid and trans ester included in the study were x4 kJ mol-l more stable. The natural question to ask at this point and the motivation of this paper is, " does this energetic consistency remain when three fatty acids are linked by a glyceride backbone to form one of the biologically important triglycerides ? " EXPERIMENTAL Briefly, the sample is injected in the form of a hexane solution into a stirred slurry of Pd catalyst and hexaiie by means of a mm3 syringe. Sample injections are alternated with injections of a standard (methyl oleate) for which the heat of hydrogenation is assumed to be known (-125.1 kJmol-i).ii The ratio of heat output per mole of the standard to that of the sample leads to the heat of hydrogenation of the latter.Completeness of reaction was tested by transesterifying the reaction product at 80°C in methanol containing sodium rnetho~ide.~ The resulting methyl ester was extracted into benzene and subjected to g.1.c. analysis. In no case was any unsaturated methyl ester detected in the transesterification product.Tests with saturated esters intentionally doped with small amounts of unsaturated esters showed that unsaturated ester present in an amount of 0.4 % of the saturated ester gave a clearly discernible chromatographic peak. Detailed discussion of experimental design and method has been RESULTS AND DISCUSSION Heats of hydrogenation for five simple unsaturated triglycerides (triacylglycerols) are shown in table 1. Each entry in table 1 gives the mean of 16 thermochemical runs on aliquot portions of one solution. The 95 % confidence limits include errors, presumably random, resulting from the method but do not include weighing errors, dilution errors or variations in purity from one sample to the next. Thus, agreement between two sets of experimental runs may be slightly outside the confidence limits 2868D .W . ROGERS AND D . N . CHOUDHURY 2869 of either one because more sources of error are included in the deviation between sets than are reflected in the deviations among members of the same set. Systematic nomenclature for the compounds listed in table 1 is : glycerol cis-9- octadecenoate, glycerol trans-9-octadecenoateY glycerol cis-6-octadecenoate, glycerol cis-9-cis- 1 2-octadecadienoate and glycerol cis-94s- 12-cis- 1 5-octadecatrienoate in the order shown. Determinations were carried out on 5-10 mg of triglyceride and reaction times were x 12 s, permitting repetitive determinations every 2 to 3 min. The high sensitivity of hydrogen microcalorimetry makes it the method of choice for thermo- chemical studies of the unsaturated triglycerides because they are commercially available only in small quantities.Each datum in table 1 is the arithmetic mean of aliquot portions of a solution made up from 500 mg of triglyceride in 5 cmS of hexane. We are presently refining the method so that an entire set of thermochemical runs can be made on 100-200 mg of sample with the intention of determining the heats of hydrogenation of monoglycerides, diglycerides and mixed triglycerides. The latter compounds are more difficult to prepare in high purity than simple triglycerides, consequently they are available in even smaller quantities. TABLE 1 .-HEATS OF HYDROGENATION* OF SOME SIMPLE UNSATURATED TRIGLYCERIDES compound formula wt -AHh/kJ mol-1 95 % C.L./kJ mol-1 triolein triolein trielaidin trielaidin tripetroselinin tripetroselinin trilinolein trilinolein trilinolenin t ri lin olenin 885.4 380.0 381.1 885.4 378.1 375.7 885.4 377.0 379.1 879.4 755.6 758.5 873.3 1132 1133 3.5 3.2 1.9 2.1 0.5 0.8 1.4 2.1 1.3 3.3 * Relative to a value of - 125.1 for methyl oleate.While this work was developing, the principal problem was poor accuracy. We feel that this problem has been largely overcome for large molecules, partly by technical refinements which have reduced the standard deviation and partly by exploiting the rapidity of the method so as to perform many replicate measurements of AH,.,. At comparable levels of the sample standard deviation, s, many replicates give more reliable results than few as shown by the curve of the 90 and 95 % confidence limits as a function of the number of measurements, N, in fig.1. Accordingly, each entry in table 1 represents the arithmetic mean of 16 hydrogenation runs. The confidence limits, C.L., t s C.L. = - JNY where t is Student’s t parameter, are comparable with the best in the thermochemical literature. None of this is to suggest that we have refined hydrogen calorimetry to the point that it gives relative errors as small as those obtained in combustion calorimetry where measurements can be made to kO.01 %. The advantage of hydrogen calori- metry comes about because its absolute error is independent of molecular size, as discussed in Part 2 of this series. Large molecules and small molecules pose quite different challenges to the thermochemist.Because pure samples of large molecules2870 HEATS OF HYDROGENATION OF LARGE MOLECULES are usually expensive, frequently of limited solubility, and because the interesting part of the molecule is a double bond, a functional group, or a strain site, located at one specific point in a molecule which is otherwise inert, sensitivity is at a premium. For example, when we want to look at the isomerization energies of cis and trans isomers of alkenes, the double bond is the only thermochemically interesting part of the molecule. The saturated parts of the molecule, while they influence its energy through steric interference, are, from the point of view of hydrogen thermochemistry, inert. The larger the unreactive part of the molecule, the more sensitive must be any thermochemical method ; a hydrogenation method sensitive enough to determine AHh using 5 mg samples of pentene would require 20 mg of eicosene.I a I2 16 20 N FIG. 1.-Variation of t / z / N with the number of measurements contributing to a sample standard deviation, s. Upper curve, 95 %; lower curve, 90 %. Conversely, the exacting criteria of accuracy demanded of thermochemistry done on small molecules are not realistic criteria for large ones. The triglycerides studied here are not available in >99 % purity. While avoidance of cumulative error encourages one to reduce thermochemical error as much as possible, modifications of the method which increase accuracy much beyond the present level at the expense of sensitivity are not justified at the level of sample purity currently attainable.Taking precision to be a good estimate of acc~racy,~ the results of table 1 are probably reliable to better than 1 %. Since these compounds were described as " 99+ % pure ", the limit of accuracy imposed by the method is about the same as that imposed by sample purity. In answer to the question posed iii the introduction, heats of hydrogenation of the triglycerides of cis unsaturated acids (the naturally occurring form) do indeed show aremarkable regularity, being almost exactly - 12672 kJ mol-l(- 30n kcal mol-l) where n is the number of double bonds in the compound. Each heat of hydrogenation is slightly larger in absolute value than would be expected by taking three times the heat of hydrogenation of the corresponding fatty acid residue.2 For Irregularities are small and close to the limits of experimental error.D.W. ROGERS AND D. N. CHOUDHURY 287 1 example, in the case of triolein, 380 > 3(125.1), indicating a slightly greater structural relaxation on going from the cis unsaturated triglyceride to the saturated triglyceride than we observe on going from the unsaturated free acid to its saturated product. These observations support the model of fatty acid residues which suffer little constraint imposed upon them by the glycerol backbone in triglycerides. Crowding effects, which might be expected to appear when nine double bonds occupy the same molecule, are not observed. There is neither increased strain in the reactant molecule brought about by introducing presumably planar elements into its structure, nor is there a stabilizing effect brought about by cooperative interaction between two or three double bonds in an acid residue, as contrasted to one.On the contrary, the double bonds in an acid residue are energetically entirely independent of one another, which is consistent with the picture of alternant (as contrasted to conjugated) double bonds existing in acid residues which are well separated. In an inert solvent like hexane, the acid residues probabJy radiate away from a central glycerol axis at mutual angles of 120". The evidence is also consistent with a strained, unsaturated trigly- ceride going to an equally strained saturated compound, but we think this is unlikely. The triglyceride containing all trans acid residues, trielaidin, has a heat of hydro- genation which is smaller in magnitude than the cis triglycerides but the effect is not as pronounced as one would expect from the usual stabilization of a trans carbon chain relative to the cis configuration.Trans isomers are usually 4-5 kJ mol-1 more stable than cis isomers 1p 2* 6 s unless very severe steric interference are involved as they are not in the present case. This leads to an expected lowering of 12-15 kJ mol-1 of the magnitude of the heat of hydrogenation of the trans triglyceride relative to the cis. The actual lowering is only 3.6 kJ mol-l. Evidently steric interactions which make the cis free acid and ester less stable than the trans form are replaced by inter-chain crowding in the triglyceride which removes most of the stabilizing effect of the trans configuration.None of these effects is large, however, and we feel that they should be regarded as second order perturbations on a fundamentally open structure which suffers little in the way of steric constraint. or ring strain CONCLUSION This paper represents a significant step in our attempt to apply hydrogen calori- metry to large molecules. The largest molecules examined by this method prior to the present work have a formula weight of less than 30OY1s lo* l1 while all data in table 1 are for molecules about 3 times as large. The largest molar heat of hydrogen- ation previously observed is -584.5 (1, 7 octadiyne)12 while the value for trilinolenin reported here is - 1132 kJ mol-l. We hope that the innate advantages of hydrogen calorimetry over combustion calorimetry for large hydrocarbons and related un- saturated molecules will lead to better understanding of their thermochemistry and of the forces involved in determing their structure.The authors acknowledge the support of the US. National Institutes of Health during this work. D. W. Rogers and N. A. Siddiqui, J. Phys. Chem., 1975,79, 574. D. W. Rogers, 0. P. A. Hoyte and R. K. C. Ho, J.C.S. Farahy I, 1978, 74,46. D. W. Rogers and P. A. Papadimetriou, Mikrochemica Actu, 1974, 937. Chromatography/Lipids, Bulletin 721B-EsterificationY Supelco Inc., Bellefonte, Pa., 16823, (1975). J. L. Jensen, in Progpzss in Physical Organic Chemistry, ed. R. W. Taft (John WiIey and Sons, New York, 1976), vol. 12, p. 189. 1-912872 HEATS OF HYDROGENATION OF LARGE MOLECULES G. B. Kistiakowsky, J. R. Ruhoff, H. A. Smith and W. E. Vaughan, J. Amer. Chem. Soc., 1936, 58, 137. R. B. Turner, D. E. Nettleton, Jr. and M. Perelman, J. Amer. Chem. SOC., 1958, 80, 1430. D. W. Rogers, H. von Voithenberg and N. L. Allinger, J. Org. Chem., 1978, 43, 380. lo R. B. Turner, W. R. Meador and R. E. Winkler, J. Amer. Chem. SOC., 1957,79,4122. D. W. Rogers and S. Skanupong, J. Phys. Chem., 1974,78,2569. lZ T. L. Flitcroft, H. A. Skinner and M. C. Whiting, Trans. Furuday Soc., 1957, 53, 784. ' J. L. Franklin, Ind. and Eng. Chem., 1949, 41, 1070. (PAPER 8/357)

ISSN:0300-9599    

DOI:10.1039/F19787402868

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Kinetics, Thermochemistry and Mechanism of Hydrogenolysis of Aliphatic Aldehydes on Ni-SiOz BY KEVIN F. SCOTT,"? MOHAMMAD A. ABDULLA AND FALAH H. HUSSEIN Department of Chemistry, University of Sulaimaniyah, Iraq Received 21st March, 1978 An investigation of the kinetics of the hydrogenolysis reaction of aliphatic aldehydes on NiSiO' by stopped-flow chromatography is described and the rate plots fitted to theoretical equations which reveal that the reaction proceeds in two kinetically discernible stages, one apparently by a LangmUir- Hinshelwood mechanism and the other by a Rideal-Eley mechanism. These conclusions are supported by the dependences of the rates on the hydrogen pressure. Thermochemical measurements using a microcalorimetric device, have enabled the determination of the thermochmistry of a reaction leading to a surface species.A discussion of the prevailing mechanism is presented. The hydrogenolysis reaction on Ni-SiO, and other supported metal catalysts has received wide attention in the literature, particularly in the case of the reaction of hydrocarbons,l* where on nickel, the reaction has shown considerable selectivity toward cleavage of terminal C-C bonds of hydrocarbon chains and in some cases almost exclusive terminal cleavage.2 Investigation of this reaction while operating the catalyst in a " chromatographic " regime provided new mechanistic information and this paper describes an extension of this work wherein the functionality of the terminal carbon atom is altered to a carbonyl group. Molecules with functionality other than simple alkanes have received scant attention with respect to the hydrogenolysis reaction; in particular the reaction of aldehydes on metal catalysts does not appear to have been widely reported, although investigations of aldehyde and alcohol adsorptions on Ni have been studied using i.r.spectroph~tometry.~ The results of this present work indicate that a study of the hydrogenolysis reaction using molecules of different functionality may throw important new light on the chemical mechanism involved. Investigations reported here were of two kinds : a kinetic investigation in which the reaction was observed by stopped- flow chromatography in a chromatographic system, using Ni-SiO, as a catalyst and as a stationary phase to effect the separation of the products, and thermochemical investigations where a microcalorimetric method was employed to measure the enthalpies of reactions occurring on the catalyst.EXPERIMENTAL Kinetic studies were carried out using a Varian model 920 gas chromatograph fitted with an integral katharometer detector, and an external flame ionization detector, constructed in the laboratory. The chromatograph was fitted with 1 and 3 m stainless steel columns 4.5 mm i.d., 6.7 mm o.d., checked for catalytic inactivity by blank experiments, and packed with catalyst material. A stop-tap was inserted in the carrier gas supply directly before the injection port, to allow stopped-flow chromatography to be carried out. The catalyst was prepared as described previously,' from a chromatographic grade of silica gel (Silica Gel for Adsorption Chromatography 60-120 mesh, B.D.H.) and nickel nitrate (Analytical Grade, Riedel De Haan AG, Seelze, Hannover).The resulting catalyst was a uniform t Present address : Inorganic Chemistry Laboratory, Oxford University, South Parks Road, Oxford. 28732874 black colour and contained 2.3 % w/w nickel metal. After packing, the catalyst was activated at 350°C using a hydrogen carrier gas (30 cm3 min-I) for 1 h ; this procedure was repeated before each new kinetic experiment. Hydrogen carrier gas was purified catalytically to remove traces of oxygen by passage through Ni-Si02 catalyst at 250°C, and dried using liquid nitrogen traps and freshly zctivated silica gel. Reagents were introduced by means of syringe injections ; in most experiments aliquots of 4-5 mg were employed.The Aux of product hydrocarbon issuing from the catalyst column was measured by interrupting the carrier-gas flow for periods of 1 min at intervals during the development of the reaction chromatogram, a procedure which generated a chromatogram of the products from which the molar flux of each component could be ~alculated.~ The detectors were calibrated by introduction of measured amounts of alkanes, and the catalyst was calibrated for its retention of alkanes so that retention times could be used for product identification. Thermochemical measurements were carried out using a microcalorimeter system which was a modification of the design of Jones et a1.' Principally, the modification involved replacing the bed of the calorimeter with a packing of the supported metal catalyst, around a filament, being chosen of the same metal as on the catalyst support.The calorimeter consisted of a Pyrex tube nominally 6mm o.d., and 4mm i.d. which contained a coil of nickel wire of total length 40 cm, of diameter 0.33 mm, wound in a helix 1.9 mm in diameter. Copper connecting wires were brazed to each end, and the coil surrounded by a packing of the supported Ni-SO,. The packing material was retained by plugs of glass wool, and glass capillary tubing was used to connect the calorimeter to the chromatographic column which followed it. A chromatographic grade of silica gel was used as a stationary phase in the latter, and the column effluent fed directly to a flame ionization detector.The temperature of the filament was controlled by a self-balancing bridge circuit used by Jones et a1.' and by Wolstenholme,6 which achieved the balance of a bridge (of which one arm was the filament) by varying the bridge current. The nickel filament and other bridge resistors were calibrated using a Wayne Kerr autobalancing bridge (Type B642). A coil of resistance 0.363 SZ at 20"C, and of temperature coefficient O.O064"C-', could be heated to 6OO0C, using currents of up to 2.4 A, while dissipating up to 10 W. A certain degree of lagging (using glass wool) around the calorimeter was advantageous in reducing noise caused by draughts, although excessive lagging was undesirable since for quantitative measurements of exothermic reactions, a reasonable standing power dissipation was required.Under our conditions, a stream of 30 cm3 min-l of dry hydrogen was used as a carrier gas, and reagents in quantities of 0.2-10 mm3 were introduced into the carrier stream through a septum cap above the calorimeter. The injected material was allowed to fall first on to a plug of glass wool maintained at 200°C by an external heater, before entering the calorimeter, since the material then entered as a sharp plug, which led to greater precision in the thermal measurements. The amount of heat abstracted from or supplied to the calorimeter as a result of any reaction occurring was determined by recording continuously the voltage across the filament which, together with a knowledge of the resistance of the filament at its operating temperature, could be used to compute the power dissipation. The integral with respect to time of the change in power dissipation during a reaction was then a measure of the enthalpy of the reaction under the experimental conditions.The microcalorimeter system was checked for quantitaiive behaviour by measuring the enthalpies of reactions which could be reacted to completion to gas phase products. Such reactions included : the hydrogenation of benzene, the complete hydrogenolysis of n-hexane to methane, and the complete hydrogenolysis of propanal to methane and water. In all these cases measured enthalpies corresponded within 10 % to the accepted values. HYDROGENOLYSIS OF ALDEHYDES ON Ni-Si02 CHROMATOGRAPHIC PROPERTIES OF THE Ni-Si02 CATALYST Since stopped-flow reaction chromatography has been the principal method for deter- mining the kinetics described in this paper, the chromatographic properties of the Ni-Si02 catalyst are important and deserve brief description.In cases like those described here, where investigations are carried out without a separate analytical chromatographic column,KO Po SCOTT, M. A. ABDULLA AND F. H. HUSSBIN 2875 the general requirements are that the catalyst should be able to resolve reactants and products under conditions of reaction, and the chromatographic capacity of the catalyst should be such that the concentrations of products generated by flow-stops under reaction conditions are those which result in linear chromatography (Le., the relevant range of the adsorption isotherm should be linear).The catalyst used in these studies fulfilled both of these requirements . Injection of any of the reactant aldehydes onto a three metre catalyst column at tempera- tures up to 200°C resulted in reaction products only being eluted from the column, the retention of the aldehydes being extremely long. Even after following a reaction for several hours, there was no indication that the reactant had moved any significant distance along the column. This behaviour of the reactants was satisfactory for three reasons : first, the time for which the reaction could be studied was not limited by the retention time of the reactant ; second, during the progress of the reaction, the reactant remained effectively in the first section of the column leaving a maximum column length for separation of the products; and third, the number of catalytic sites occupied by the reactant was fixed which results in a simpler theoretical treatment (see below).t I I I I I I 2.15 2.20 2.25 2.30 2.35 2.40 2.45 lo3 KIT FIG. 1 .-Chromatographic retention volumes (in column dead volumes) for n-alkanes on Ni-Si02 catalyst : (a) ethane, (b) propane and (c) n-butane. Hydrocarbon products were, however, eluted more rapidly from the catalyst column and could be conveniently separated in periods up to 10min under reaction conditions. The chromatographic characteristics were determined before investigation of reactions by injecting n-alkanes on to the column at different temperatures; the results are shown in fig. 1 where the corrected retention volumes (expressed in column dead volumes) are plotted against the reciprocal of the absolute temperature.At all temperatures investigated methane was eluted at the column dead volume as an unretained peak, and other hydrocarbons retained up to ten column dead volumes were eluted symmetrically using injection amounts of up to mol. RESULTS KINETIC RESULTS AT 2 atm HYDROGEN PRESSURE The results presented here indicate that the reactions undergone by aliphatic aldehydes C2-C4 with hydrogen in the presence of a Ni-Si02 catalyst to yield volatile products can be represented by RCHZCHO + Hz + RCHS + CH4 +H20 (1)2876 HYDROGENOLYSIS OF ALDEHYDES ON Ni-SiQ, RCHZCHO + Hz + RCHZCH3 + H20 (11) RCH3 + H2 + RH + CH4 etc. (111) Over the whole range of experimental conditions (130-200°C and 0.01-2 atm hydrogen pressure) reaction (I) has been found to be the predominant reaction, with reaction (11) accounting for <4 % of the total stoichiometry, and reaction (111) insignificant below 180°C.2 Both reactions (11) and (111) increase in significance with increasing temperature; at 300°C at 2 atm hydrogen pressure methane and water are the only products.However, it is reaction (I), the principal reaction in the experimental temperature range, which is the subject of detailed study in this paper. Almost exclusive cleavage of the carbon-carbon bond at the alpha position has been observed, cleavage of bonds further along the carbon chain amounting to 2-3 %, as a proportion of the volatile products in the case of both straight- and branched-chain aldehydes at temperatures up to 180°C.I I I I I I I 10 20 30 40 50 60 70 80 time/min FIG. 2.-Rates of production of ethane and methane against time for the hydrogenolysis of propanal at 177°C : (0) ethane production ; (e) methane production. When an aliphatic aldehyde is introduced into a previously activated Ni-Si02 catalyst column, under the above reaction conditions, the two hydrocarbons from reaction (I) are generated as a result of two separate rate processes. Under the conditions studied, the methane production was invariably slower than the production of the higher hydrocarbon. Fig. 2 illustrates the process in the case of propanal, where the rates of methane and ethane production are plotted against time for an experiment in which 4.8 mg of propanal was syringe-injected into a previously activated 3 m Ni-SiO, column at 177°C.The principal features of the plots are that both rates pass through a maximum and decay away with time, that the methane production is zero at zero time, and the rate of production of ethane has a non-zero starting value. The areas under the profiles give a measure of the total stoichiometry of the reaction; under the conditions of fig. 2, approximately 2.3 x mol ofK. F. SCOTT, M. A . ABDULLA AND F. H. HUSSEIN 2877 ethane, 1.3 x mol of methane and <lo-’ mol of propane were generated from an injection of 8.3 x mol of aldehyde. The balance of carbonaceous material which remains on the catalyst, and which reacts only slowly at the reaction temperature was substantially recovered as methane when the column was heated to 350°C under hydrogen and the catalyst restored to its original condition. It appears from these observations, and particularly from the zero methane produc- tion rate at zero time, that the reaction generates products as a result of consecutive reactions, which could be represented by the following scheme : 1 ( 2 ) 2 ( 7 ) 3 RCHZCHO + A, + Ca -j CH,+H,O + RCHj where A and C represent adsorbed species.The stoichiometric results (e.g., those above) indicate that, at the experimental temperature, only a proportion of these species, labelled A, and C,, is active, and the balance is inactive (labelled Ai and CJ, only reacting when the temperature is raised. We now turn attention to the investi- gation of the kinetics of the reaction occurring in the experimental temperature range indicated by the horizontal processes, 1, 2 and 3 in the above scheme.From the observation that the rate of ethane production in fig. 2 has an initial non-zero value it is concluded that the adsorption process 1 is rapid compared with further steps. In the development of the kinetic equations for product formation process 1 can, therefore, be eliminated. The possible factors which, at a constant temperature, could control the rate of generation of the two hydrocarbon products from the catalyst surface are : surface concentrations of A and C, surface concentra- tion of hydrogen, and hydrogen gas pressure. Under conditions of constant hydrogen pressure, the last of these will be absorbed into an experimental rate constant and further, if the hydrogen pressure is such as to render the available catalyst sites effectively fully covered with adsorbed hydrogen (ix., the situation prevailing at the top of the hydrogen adsorption isotherm), the surface hydrogen concentration will be given by the total number of available sites less those occupied by other surface species.By setting up differential equations corresponding to the various dependences of the rate of generation of the hydrocarbons on these factors, solving these equations numerically and comparing with the experimental data, it was possible to identify those rate equations which generated functions most closely followed by the experi- mental data. The simplest of these rate equations are as follows : - k,A,(M - A, - Cln), d A a - = dt and dCa -- - Bk2 A,(M- A, - Cln) - k,C, dt where A, represents the surface concentration of that species, C represents tLt total concentration of the C, and Ci species, and M is a parameter which is a measure of the total number of available sites on the catalyst surface which are involved in the reaction.n is a parameter which takes into account the different numbers of active sites occupied by C relative to A. k2 and k3 are rate constants and B is a parameter which represents the proportion of species A, which reacts to give the reactive species2878 HYDROGENOLYSIS OF ALDEHYDES ON Ni-Si02 C,. Thus (M-A,- C/n) represents the concentration of hydrogen on the catalyst surface, and the equations can be solved by writing C in terms of A, and C, from the total stoichiometry, and inserting a boundary condition that the initial concentration of A occupies a fractionfof the total surface sites.It is important to observe that these rate equations which gave optimum agreement with the experimental data, indicate that species A, reacts at a rate controlled by surface hydrogen, and species C , reacts with a rate independent of surface hydrogen. TABLE COMPARISON OF EXPERIMENTAL AND THEORETICAL VALUES FOR THE RATES OF PRODUCTION OF HYDROCARBONS IN THE HYDROGENOLYSIS OF PROPANAL ON Ni-SiOz AT 177°C rate of production of hydrocarbon/mol min-1 X lo7 time Imin 2.3 6.4 9.4 12.4- 16.4 19.4 22.4 25.5 28.5 31.4 34.4 37.4 40.4 43.4 ethane 8.31 (8.69) 8.45 (8.56) 7.94 (8.10) 7.33 (7.37) 6.40 (6.12) 4.94 (5.11) 4.06 (4.14) 2.98 (3.26) 2.44 (2.51) 1.78 (1.91) 1.48 (1.42) 1.30 (1.05) 1.15 (0.77) 1.03 (0.56) methane 0.83 (0.64) 1.71 (1.69) 2.07 (2.26) 2.50 (2.66) 2.64 (2.94) 3.00 (2.98) 3.00 (2.90) 2.86 (2.75) 2.71 (2.53) 2.34 (2.29) 2.11 (2.03) 1.97 (1.78) 1.71 (1.54) 1.57 (1.32) Theoretical values shown bracketed. TABLE 2.-vALUES OF PARAMETERS WHICH GAVE OPTIMUM CORRESPONDENCE BETWEEN THEORETICAL AND EXPERIMENTAL RESULTS IN THE HYDROGENOLYSIS OF ALDEHYDES reactant pr opanal propanal propanal propanal n-butanal n-but anal n- bu t anal 2-methyl propanal 2-methyl propanal 2-methyl propanal M/mol kz/mol-l tempPC x 105 min-1 kjlmin-1 n B 155 161 174 177 151 162 172 152.5 160 173 4.2 3.45 3.4 3.1 11.0 6.6 6.9 11.0 8.8 8.7 860 2200 3100 3900 1 70 420 920 200 480 1000 0.022 0.030 0.057 0.066 0.019 0.026 0.048 0.019 0.026 0.050 3 0.27 0.71 3 0.29 0.65 3 0.43 0.74 3 0.58 0.64 4 0.28 0.48 4 0.61 0.48 4 0.55 0.48 4 0.41 0.50 4 0.39 0.68 4 0.47 0.65 The results of curve fitting the experimental data to the filnctions of eqn (1) and (2) are shown in table 1, for the reaction of propanal at 177°C.A close correspondence is observed between the experimental values and the predicted values for the rates of production of the two hydrocarbon products. The values of the various para- meters found for optimum correspondence, for different temperatures and different reactant aldehydes, are shown in table 2. As would be expected, the value of M, which represents the number of sites involved in the reaction, increases with increasing size of reactant molecule, and this trend is also reflected in the fraction f of these sites which are initially occupiedK.F . SCOTT, M. A . ABDULLA AND F . H. HUSSEIN 2879 by reactant, which decreases with increasing reactant molecular size. The value of M also decreases with rising temperature, and this may be due to the increase in velocity of the adsorption process. As expected the rate constants increase with temperature; k3 at a fixed temperature has a value which is approximately invariant with reactant molecule, while k2 decreases by a factor of =4 on going from propanal to the butanals. This suggests that the step 3 (characterised by k3) is essentially the same in all cases while step 2 changes in character with reactant molecule. lo3 KIT FIG. 3.-Arrhenius plots for the production of hydrocarbons in the hydrogenolysis of aldehydes : e, propane production from n-butanal ; A, propane production from 2-methyl propanal ; a, ethane production from propanal ; 0, methane production from n-butanal ; A , methane production from 2-methyl propanal ; 0, methane production from propanal.The parameter n has no great influence on the curve fitting in the range 3-6, and so values were chosen in this range to correspond to the expected carbon numbers of the species A relative to C . The factor B increases with rise in temperature; this is expected since it reflects a crude species activity distribution, and a greater propor- tion will react at higher temperatures. TABLE 3 .-ARRHENIUS DATA FOR THE HYDROGENOLYSIS OF ALDEHYDES ON Ni-SiOz AT 2 atm HYDROGEN PRESSURE RCH3 production rate methane production rate reactant min-1 mol-1) /kJ mol-1 min-1) /kJ mol-1 b i o ( A / E b i o ( A / E propanal 14.6k0.5 94+10 S.lkl.0 80+5 n-bu tanal 17.8k0.5 130+5 6.7k0.5 70+5 2-methyl propanal 17.1k0.5 12055 7.5k0.5 75f5 The Arrhenius plots for the production of methane and the higher hydrocarbon RCH3 from propanal, butanal and 2-methyl propanal, are shown in fig.3 and the Arrhenius parameters are listed in table 3. This table shows that both the pre- exponential factor, and the activation energy of the reaction step which leads to the2880 HYDROGENOLYSIS OF ALDEHYDES ON Ni-Si02 higher hydrocarbon, increase with increasing reactant molecular weight, but the parameters associated with the formation of methane are unaffected within experi- mental error.This is expected if the species which is hydrogenated to give methane is the same in each case. The effect of water on the reaction rate was investigated by performing the experiments using a carrier gas saturated with water vapour at 0°C. The reaction followed a similar pattern as under dry conditions with similar rates, but the specific. activity of the catalyst was reduced, and the reactant occupied a longer section of the reactor column. This resulted in approximately the same molar flux of products from the column as under dry conditions, but the width of the stop-flow peaks indicated that reaction was occurring in a section of catalyst approximately twice as long. INFLUENCE OF HYDROGEN PRESSURE ON REACTION The differential equations which were found to correspond to the experimental rate plots in the forepoing section indicate that the reaction of the surface species A, (to yield the higher hydrocarbon) depended upon the concentration of surface hydrogen, while the reaction of species C, (to yield methane) was independent of the hydrogen surface concentration. The influence of the hydrogen gas pressure was, therefore, of importance. - 1 5 - 1.0 -05 0 log10 (P&tm) FIG.4.-Effect of hydrogen pressure on the rate of formation of (0) methane and (e) ethane in the hydrogenolysis of propanal. A series of experiments were carried out in which propanal was reacted in a catalyst column using a nitrogen+hydrogen mixture as a carrier gas. The composi- tion of the mixture was adjusted to give hydrogen partial pressures ranging from 0.01 to 2 atm.The rates of production of methane and ethane (expressed as pseudo- first-order rate constants) are plotted against the hydrogen partial pressure using logarithmic axes in fig. 4.K . F . SCOTT, M . A . ABDULLA AND F. H . HUSSEIN 288 1 Fig. 4 shows that the rate of production of the higher hydrocarbon is substantially invariant with respect to the hydrogen pressure. The slight decrease in rate as the hydrogen pressure is reduced can be attributed to the lowering of the surface hydrogen concentration which would accompany the reduction in pressure. In contrast to this the methane production varies markedly with the hydrogen pressure, showing an experimental order with respect to hydrogen of = + 1.4.Whereas methane produc- tion is dependent upon hydrogen pressure but independent of the surface hydrogen concentration, ethane production is independent of hydrogen pressure, but dependent upon the surface hydrogen concentration. Thus the dependence of the reactions on hydrogen is in accordance with the rate laws determined from curve fitting. I! b 1.6 1.4 1.2 1 .o 0.8 D.6 5. 0.4 4 9 0.2 2 .c) 0 - _ _ 100 230 300 400 500 600 700 filament temperature/"C FIG. 5.-Hydrocarbon yield and heat change for the hydrogenolysis of propanal : (0) ethane yeilds, (0) methane yield, (A) heat change. THERMOCHEMICAL RESULTS Samples of aldehydes were introduced into the microcalorimeter while the filament was maintained at different temperatures. The enthalpy of the reaction occurring was determined from the change in power supplied to the calorimeter as detailed previously.The stoichiometry of the reaction was determined from the integration of the peaks eluted from the chromatographic column which followed the calorimeter. By varying the calorimeter temperature it was possible to vary the proportion of injected aldehyde which underwent the various reactions (1)-(III) above, and also the partial reaction yielding the higher hydrocarbon RCH3 and leaving a certain2882 HYDROGENOLYSIS OF ALDEHYDES ON Ni-Si02 amount of surface species C remaining. The results are illustrated by fig. 5 where the total enthalpy and the hydrocarbon yields are plotted against calorimeter filament temperature in an experiment where 0.15 mg aliquots of propanal were introduced by syringe.Below 150°C there is little reaction, but as the temperature is raised first the higher hydrocarbon ethane is generated, followed by an increase in the yield of methane at higher temperatures. Above a filament temperature of 4OO0C, the ethane itself begins to hydrogenolyse, and finally as the temperature is raised to 6OO0C, complete reaction occurs to yield methane and water. The reactions undergone by the aldehyde can be expressed by the three following equations : (IV) 09 (VI) where C* represents the surface species which may or may not contain oxygen. CH3CH2CHO + CH3CH3 + C* (+ H2O) CHSCH2CHO + CH3CH3 + CH4 + H20 CH3CHsCH0 + 3CH4 + HzO 1.6 1.4 c, 1 7 1.2 a 1.0 c? s. $ + & 0.8 Q.6 + W 0" 0.4 I @I Q.2 0 0 1 2 3 4 5 6 7 8 9 10 (C, +2C2)/mol x lo6 FIG.6.-Plot of eqn (3) for propanal. Reactions (V) and (VI) are both complete reactions with readily calculatable enthalpies while reaction (IV) is a partial reaction with a surface product. By writing down the carbon atom mass balance equations, and a thermal balance equation and solving these, it is possible to generate an expression from which the enthalpy of reaction (IV) is readily determined. (3) where C1 and Cz are the molar yields of methane and ethane respectively, AH,,, AH, and AH,, are the enthalpies of reactions (IV), (V) and (VI), respectively, and Q is the heat change measured. Fig. 6 shows a plot of eqn (3) using data in the range 1 8O-48O0C, using enthalpies of formation estimated at the mean catalyst temperature. From the slope, the heat of reaction AH,, was calculated to be - 170 +5 kJ mol-l. Similar calculations in the case of the reaction of 2-methyl propanal gave a value of - 194+ 5 kJ mol-1 for the heat of reaction (IV).For propanal this equation is : Q-(Cl +3C2)AHv+ C2AHv, = -AHIv(Cl +2C2) +constantK. F. SCOTT, M . A . ABDULLA AND F. H. HUSSEIN 2883 DISCUSSION The hydrogenolysis of aliphatic aldehydes has been found to proceed with essentially exclusive terminal C-C bond cleavage under the experimental conditions outlined in this paper. This cleavage yields two fragments : an alkyl group, which is hydrogenated rapidly to the corresponding alkane in the gas phase, and a C1 species which is hydrogenated at a measurable rate to methane in a separate rate process. From the product distribution with different reactants, it is likely that the methane originates from the carbonyl group and the higher hydrocarbon from the alkyl group since in no case in our experiments was extensive fragmentation of the alkyl group observed.The C-C bond cleavage reaction which yields the higher hydrocarbon appears from the kinetic dependences of its formation and by the order of the reaction with respect to hydrogen pressure, to follow the Langmuir-Hinshelwood mechanism and exhibits an apparent activation energy of 90-1130 kJ mol-l, which is a value very much smaller than that reported for the hydrogenolysis of alkanes at 1 atm hydrogen pressure where values of between 360-550 kJ mol-1 were observed. Essentially to achieve C-C bond cleavage catalytically, the catalyst must provide an environment in which the C-C bond energy is reduced. It has recently been shown ’ that the probable mechanism by which this occurs in alkane hydrogenolysis on Ni-Si02, is through the formation of multiple bonds from the carbon atoms adjacent to the C-C bond to the metal surface. As each of these are formed (endothermically in a series of surface equilibrium reactions) hydrogen is released into the gas phase which gives a large positive entropy effect and consequently a large pre-exponential factor in the experimental Arrhenius equation.At the same time the endothermicity of the equilibrium steps accumulates in the apparent activation energy, which is correspondingly very large. The reaction appears to exhibit a non- catalytic activation energy but does in fact proceed at very low temperatures as a result of the large value of the pre-exponential factor.Here the catalytic environment produces reduction of the C-C bond energy and, therefore, a low real activation energy, but this process involves a series of endothermic reactions the enthalpies of which make the activation energy appear larger. In aldehyde hydrogenolysis, the requirement that the C-C bond energy be reduced remains, if the reaction proceed catalytically. However, the Arrhenius parameters observed in this case are quite “ normal ”, and so the reduction in effective C-C bond energy must be caused by some other mechanism, for which the presence of the carbonyl group must ultimately be responsible. The thermochemical measurements indicate that the C-C bond cleavage reaction in the case of aldehydes is considerably exothermic, with a heat of reaction of some -(130-150) kJ mol-l, after allowing for the heat of adsorption of the reactant molecule.If this heat of reaction were due to the large heat of adsorption of species C compared to that of A, some of this energy could be utilized in the C-C bond cleavage in a concerted reaction, in which the C-C bond is broken as a new bond is formed between C and the surface. The cleavage reaction could then proceed with only perhaps one or two C-metal bonds to the surface in species A, thus giving the normal Arrhenius parameters observed. From the kinetic rate equations, and the dependence of the reaction on hydrogen pressure, the methanation of species C remaining after the C-C bond cleavage has taken place, appears to proceed via a Rideal-Eley mechanism, with the surface species C reacting with hydrogen in the gas phase. Infrared studies have indicated the presence of acyl groups and what appears to be adsorbed CO on a Ni surface after exposure to acetaldehyde.The acyl groups may well correspond to the species2884 HYDROGENOLYSIS OF ALDEHYDES ON Ni-Si02 A, while the adsorbed CO would correspond to the species C . If this is correct, the kinetics of the methanation of C should be close to those of the CO methanation reaction. The order of the reaction of C with respect to the hydrogen pressure is very close to that found for the CO methanation (+ 1.4 in both cases),8 and the activation energies are also similar (75 and 80-120 kJ mol-l, respectively). Van Herwijnan et aL9 report that a Langmuir type model can be used to describe the methanation of CO and suggest that an enolic species on the metal surface reacting with hydrogen is the rate limiting step, although they were not able to distinguish between Langmuir-Hinshelwood and Rideal-Eley behaviour since the variation in hydrogen pressure was small in their experiments. It seems that the intermediates may well be the same in the two cases : either adsorbed CO or the enol form. The reaction could then be represented by : RCHzCHO RCHs 0 OH / " 7 R-CH2-C ----+ RCHz xx + H2O. I X I X I X The authors acknowledge the assistance of Mr, Shirwan 0. Baban in the drawing of uniform wires from a pure nickel sample, for the construction of the microcalori- meter used in this work. J. H. Sinfelt, Adu. Catalysis, 1973, 23, 91. K. F. Scott and C. S. G. Phillips, J. Chromutogr., 1976, 112, 61. G. Blyholder and L. D. Neff, J. Phys. Chem., 1966,70,1738. R. M. Lane, B. C. Lane and C. S. G. Phillips, J. Catalysis, 1970, 18,281. A. Jones, J. G. Firth and T. A. Jones, J. Phys. E, 1975,8, 37. J. Wolstenholme, personal communication. ' K. F. Scott and C. S. G. Phillips, J. Catalysis, 1978, submitted for publication. * R. A. Dalla Betta, A. G. Piken and M. Shelef, J. Catalysis, 1974, 35, 54. T. Van Herwijnan, H. Van Doesburg and W. A. De Jong, J. Catalysis, 1973,28,391. (PAPER 8/538)

ISSN:0300-9599    

DOI:10.1039/F19787402873

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Transport in Aqueous Solutions of Group IIB Metal Salts at 298.15 K Part 5.-Irreversible Thermodynamic Parameters for Zinc Perchlorate and Verification of Onsager's Reciprocal Relationships BY ANDREW AGNEW AND RUSSELL PATERSON* Department of Chemistry, University of Glasgow, Glasgow G12 8QQ Received 13th April, 1978 The isothermal vectorial transport properties of zinc perchlorate have been measured over the concentration range 0.1-3.0 mol dm-3 at 298.15 K. Hittorf transport numbers were shown to agree with those measured earlier using concentration cells. Conductance and salt diffusion co- efficients were also measured. These data were used to calculate mobility (&) and resistance (&) coefficients using irreversible thermodynamic theory. Analysis showed that the Onsager reciprocal relationships (O.R.R.) were obeyed.Previous papers in this series have dealt with the anomalous transport properties of cadmium iodide solutions, caused by self-c~mplexing.~-~ The anomalies in the concentration dependence of the mobility coefficients of this salt are shared by other complexed halides in the group IIB metal series in particular by cadmium and zinc chlorides.6 Studies were initiated on aqueous solutions of zinc perchlorate to allow com- parisons between complexed and uncomplexed salts of this series. It is certain that complexing, of the type observed for the halide salts, is not possible in zinc per- chlorate. Raman studies on zinc halides '-lo which have provided independent evidence of complexing have given negative results for both zinc nitrate and zinc perchlorate.ll9 l2 The activity coefficients for both salts, and transport numbers for zinc perchlorate l3 show normal behaviour and certainly no evidence for com- plexation.There is, however some doubt as to the possibility of ion-pair formation. Using conductance measurements Davies and Thomas l4 found no evidence for ion association and although Dye, Faber and Karl considered the conductance anomalous, they were unable to assign the effect to the presence of the ion pair ZnClOt. Frei and Podlahova l6 assigned to this ion pair an association constant of 4.5k0.01. It seems certain, therefore, that zinc perchlorate is completely devoid of higher complexes and that ion association, if indeed it occurs at all, is of a minor nature.Transport properties of this salt may, therefore, be used in a comparison with those of complexed zinc chloride to observe directly the effect of complexing upon the irreversible thermodynamic parameters which define the transport properties of these solutions. Excellent transport number data are available from the work of Stokes and Levien l3 on concentration cells, using zinc amalgam electrodes. In this work directly measured Hittorf transport numbers were obtained in order to verify the 28852886 TRANSPORT I N AQUEOUS SOLUTIONS Onsager reciprocal relationships, (Lik = Lki). The remaining two transport para- meters, electrical conductance and volume-fixed salt diffusion coefficients were also measured to a precision of 0.05 and 0.3 %, respectively. EXPERIMENTAL Zinc perchlorate was prepared by adding a slight excess of spectroscopically pure zine oxide (Koch Eight Laboratories) to AnalaR perchloric acid (B.D.H.) and filtering off thc unreacted oxide.This solution was made slightly more acidic by dropwise addition of perchloric acid until the pH was fir 2. On cooling to 0°C crystals of zinc perchlorate were obtained. On dissolution, these crystals gave clear solutions at all dilutions. Conductance measurements of solutions prepared from different batches of crystals prepared in this way gave reproducible conductivities. The solutions were analysed for zinc using EDTA titrations, by the methods described earlier for zinc chloride.17 The precision of the method was kO.05 %. To allow conversion between molal (mol kg-l) and molar (mol concentrations, accurate density measurements were made at 298.15+ 0.005 K.Pyknometers of capacity 30 cm3 were used and the densities obtained were reproducible to k0.02 %. Results are given in table 1, together with molality to molarity conversion equations. 11 TABLE 1.-POLYNOMIAL EXPRESSIONS OF FORM y = C ajx' X a0 C C C+ In C C C In m In m 0.997 401 1.002 753 12 1.115 7 4.833 594 1.061 359 5.650 279 0.285 227 321 1.323 017 88 X a3 C 0.00007893 C 1.728 243 84 x In C -0.006 089 6 C 8.903 89 In m 9.860 176 12 x lnm 0.171 591 501 C* -1.511 311 C -1.02225 i = O concentration range a1 a2 /mol dm-3 0.198 024 - 0.002 427 4 0.1-3.0 - 0.705 702 1.831 215 0.03-3.15 0.774 400 - 0.048 270 2 0.1-0.35 6.604 568 x 7.556 885 14 x 0.1-3.0 168.880 5 - 59.803 46 0.3-1 .O 115.267 4 -37.353 2 0.9-3.0 0.1 -2 .o 0.589 824 368 0.407 860 767 2.0-4.0 0.453 586 887 0.176 694 894 concentration range 11.4 a5 /mol dm-3 - - 0.1-3.0 1.256 318 41 x - 0.1-3.0 0.882 843 - 0.260 420 0.03-3.1 5 -0.002 091 0 - 0.1 -0.35 - - 0.3-1 .O 0.9-3.0 0.506 203 - - 6.111 581 28 x - 0.1-2.0 3.066 428 84 x - 2.0-4.0 Where p is the density of the solution at 298.15 (g CM-~), C and m are concentrations, mol dm-3 and mol kg-', respectively.Dv is the volume fixed salt diffusion coefficient (cm2 s-') and K, the conductance (In-' cm-I). Electrical conductances were measured by the methods described previously to a precision of k0.02 % and were reproducible to k0.05 %, due largely to the uncertainty in chemical analysis of k0.05 %. Results are given in table 4 where conductances at rounded concentrations in the range 0.1-3.0 mol dm-3 are given.Diffusion coefficients were measured using Rayleigh interferometry on a closed column of diffusing solution in a regular parallelopiped optical cell. The methods used were identical to those described ear1ier.l Diffusion coefficients were reproducible to k0.3 %, table 2. Data at rounded concentrations, shown in table 4, were obtained by interpolation, using computer curve fitting procedures, table 1.A . AGNEW A N D R . PATERSON 2887 Transport numbers were measured using the Hittorf method. The Hittorf cell was of a similar design to that introduced by MacInnes and Dole and used recently by Pikal and Miller.Ig* 2o The cell was constructed with 10 mm bore Pyrex glass tubing made in two sections and coupled together by a B14 " Quickfit " joint.The assembled cell could be separated into three compartments using two 10 mm bore stop-cocks. These compart- ments were the cathodic, anodic and middle portions of the cell. For the analysis of a Hittorf experiment it is essential that concentration changes are confined to the two electrode compartments. For this reason every precaution was taken to prevent thermal and gravitational convection currents within the cell. The cell was bent in several places to minimise such currents and the barrels of the stop-cocks were drilled to allow free circulation of thermostat water internally, within the stop-cock barrel. The water thermostat was maintained at 25_+0.005"C and all sources of vibration from thermostat stirrers were minimised.As with earlier designs, the electrode polarities were arranged such that the more dense electrode solution was that of the lower electrode, usually the anode. Both electrodes consisted of short lengths of spectroscopically pure zinc rod sealed into a B14/20 Quickfit cone with epoxy resin. Pure metallic electrodes were found to be unsuitable for electrolysis experiments due to side reactions, particularly hydrogen evolution at the cathode and oxide formation at the anode. These side reactions were greatly reduced by using amalgamated electrodes. These electrodes were prepared by dipping the cleaned zinc metal electrodes into dilute mercuric chloride solution for a few minutes, rinsing with distilled water and polishing gently with a clean tissue.A Solartron constant current supply (P.S.U. AS1413) was used to deliver currents of 5.8 or 15.8 mAkO.1 %. Both currents are safely below the maximum current values recommended by Pikal and Mi11er.l The current was monitored when entering and leaving the cell and no leakage of current from the cell to the thermostat was ever found. TABLE 2.-EXPERIMENTAL SALT DIFFUSION COEFFICIENTS, D, FOR ZINC PERCHLORATE Dvx 10J/cm2 s-l 1.046 1.054 1.030 1.034 1.085 1.119 1.127 1.214 C/mol ~ I r n - ~ 0.0255 0.0416 0.0823 0.0915 0.1908 0.3525 0.3632 0.6198 Dvx 105/cm2 s-' 1.222 1.493 1.507 1.574 1.564 1.539 1.347 C/mol 0.6426 1.5136 1.5677 2.1701 2.2175 2.5495 3.1705 The experimental method of determination was identical to that of Pikal and Mi1ler.l Experiments lasted from 5-15 h, dependent upon the current used and the amount of electrolysis possible before side reactions at the electrodes became troublesome.In all cases the concentration changes in the electrode compartments were restricted to a maximum of 10 % to minimise risk of concentration changes outwith the electrode compartments. The weights of solution in both electrode compartments were determined and weighed solution samples withdrawn for analysis. Two samples were also taken from the middle compartment, one close to the anodic compartment the other close to the cathodic one. Analysis of these latter solutions showed that concentration changes due to electrode reaction were confined to the electrode compartments, otherwise the experiment was discarded. Triplicate analyses of the solution samples agreed within k0.07 %, using the EDTA titration analysis method described earlier." RESULTS AND DISCUSSION CONDUCTANCE A N D DIFFUSION The electrical conductances of solutions were measured in the range 0.1-3.0 mol dm-3.Salt diffusion coefficients on a volume-fixed frame of reference, D,/cm2 s-I were obtained from light interferometry, using Rayleigh optics. Experimental details of the method have been given ear1ier.l The fringe shifts, Am, were measured between two levels at & and 3 of the total height of the diffusing column of solution, as in Equivalent conductances, given in table 4, are precise to k0.05 %.2888 TRANSPORT I N AQUEOUS SOLUTIONS Harned's conductometric method.21 Since Am is proportional to AC (the cor- responding difference in salt concentration at these two levels) only if the refractive index of the solution is a linear function of concentration, Clmol d ~ l l - ~ , independent measurements of the refractive indices of these solutions were made.The relationship was not precisely linear over the full concentration range (0.1-3.0 mol dm-3), but for small concentration differences, such as encountered in the diffusion experiments, linearity was excellent. As before diffusion coefficients were obtained from eqn (1) where I is the height of the enclosed column of diffusing solution (cni) and 8 is the time in seconds. Excellent linearity was obtained in all cases and the resulting dif- fusion coefficients were reproducible to a precision of k0.3 %.This error is due largely to uncertainty in measurement of fractional fringe shifts. Experimental data, shown in table 2, was curve-fitted as a function of ,/C to obtain diffusion coefficients at rounded concentrations, table 4. In Am = - D,(n/E)28 + constant (1) HITTORF TRANSPORT NUMBERS FOR ZINC, th+ From the definition of transport number, th,i8/2F103 moles of zinc are trans- ported by a current of i/mA, during a time, 8/s. Allowing for electrode reactions at the zinc amalgam electrodes, the change in the number of moles of zinc, An, in anodic and cathodic compartments are Ananode = (1 - th,)i8/2m03 mOl Ancathode = (th, - l ) i 8 / 2 m 3 m01. (2) (3) If the weight of solution in an electrode compartment is W/g at the end of electrolysis and the initial and final zinc concentrations are mi and mf mol kg-l (solution) respectively, then on a solvent-fixed frame of reference, where M, is the molecular weight of the salt.From eqn (2)-(4) Experimental results are summarised in table 3. The first column represents zinc perchlorate molarity, C/mol dm-3. The next five columns are concentrations in units, mol kg-1 of solution. These are m i m ~ r n ~ m ~ and ma, representing the original solution concentration, final concentrations of the middle compartment, sampled at the cathodic and anodic sides and final concentrations in cathodic and anodic compartments, respectively. Wc and Wa are the weights of solution in these electrode compartments at the end of the electrolysis. Agreement between mi, mz and m;, within experimental error shows that con- centration changes were confined to the electrode compartments.Using eqn (5) zinc transport numbers (t:) were obtained from analyses of anodic and cathodic solutions, with uncertainty in the average tt of 50.01 units. These results are less accurate than might have been expected. Although amalgamation of the zinc electrodes reduced the tendency for side reactions to occur at the electrode, formation of a film of white zinc oxide or zinc hydroxide at the anode remained a problem, particularly on prolonged electrolysis. For this reasonA . AGNEW AND R . PATERSON 2889 experiments of limited duration were made and consequently concentration changes at the electrodes were much smaller than optimal ( M 10 % for dilute, reducing to 3 % for more concentrated solutions).For the most dilute solution, hydrogen gas evolution occurred at the cathode and so transport number was estimated from analysis of the anodic compartment alone. Stokes and Levien l3 have reported transport numbers from concentration cell measurements, t t , which they consider accurate to k0.002 units. These data have been interpolated to compare with the Hittorf transport numbers obtained here, table 6. The two sets, t: and th+, agree within the uncertainty of measurement. This agreement implies the validity of the Onsager reciprocal relationships (O.R.R.) which are verified by a more stringent method applied in the next section. TABLE 3 .-HITTORF TRANSPORT NUMBER DETERMINATIONS C/mol drn-3 0.097 30 0.322 25 0.460 84 0.574 31 1.015 39 1.155 71 1.463 6 1.965 6 2.424 7 C/rnol dm-3 0.097 30 0.322 25 0.460 84 0.574 31 1.015 39 1.155 71 1.463 6 1.965 6 2.424 7 mi 0.095 70 0.303 76 0.423 31 0.517 35 0.849 05 0.945 32 1.142 00 1.427 40 1.658 23 WClg - 28.430 28.912 30.074 31.701 32.788 34.190 37.146 39.095 ma "F "Z rnc m /mol kgl (ofiolution) - 0.303 43 0.423 32 0.517 39 0.848 57 0.945 36 1.142 69 1.429 00 1.658 03 wvg 26.499 27.835 28.838 29.161 - - 33.645 36.279 - 0.095 73 0.303 91 0.423 72 0.516 46 0.849 96 0.944 78 1.142 20 1.427 70 - 2Flie /mol-l x lo3 1.422 07 0.815 40 0.216 93 0.569 54 0.222 47 0.380 21 0.435 26 0.505 43 0.501 73 0.111 07 0.279 46 0.329 85 0.332 80 0.514 49 0.485 11 0.550 73 0.776 14 - 0.909 13 I 1.110 28 1.175 17 1.404 80 1.451 99 1.637 54 - ta anodic 0.406 0.3 60 0.362 0.358 - - 0.308 0.287 - zinc transport number tr cathodic - 0.384 0.359 0.360 0.336 0.325 0.320 0.307 0.279 fa (av.1 0.406 0.372 0.361 0.359 0.336 0.325 0.314 0.297 0.279 mi, m&, mt, mf and m! represent the original solution concentration, the final concentrations of the middle compartment on cathodic and anodic sides, respectively and the concentrations in the cathodic and anodic compartments at the end of the experiment.These concentrations are expressed as mol kg-I of solution. IRREVERSIBLE THERMODYNAMICS As in earlier papers in this series on group I13 metal salts zinc perchlorate will be designated as a, b, where a and b represent zinc and perchlorate ions, respectively and ra and Pb are the stoichiometric coefficients of the salt (ra = 1 and rb = 2 in this case).The theory of irreversible thermodynamics requires that, for systems close to equilibrium, there will be linear relationships between the observed flows (Ja and Jb) and their conjugate forces (Xa and Xb). In accord with earlier practice the flows J, and Jb are defined relative to stationary solvent (mol cm-2 s-l) and the forces, are defined by the local negative gradients of electrochemical potential (-grad PI) : dimensions 3 mol-' cm-l.2890 TRANSPORT I N AQUEOUS SOLUTIONS Two equivalent representations of the phenomenological equations are possible, eqn (6) and (7). In the first flows are expressed as linear functions of the forces, defining mobility coefficients, &k Ja = Laaxa+Labxb (6) Jb = Lbaxa+Lbbxb where the mobility coefficients L i k have dimensions mo12 J-1 cm-l s-l.resistance coefficients, Rik, eqn (7), In the second or inverse manner forces are expressed in terms of flows defining where resistance coefficient Rik have dimensions J cm s mol-2 and may be obtained from the mobility coefficients of eqn (6) by matrix inversion. The three isothermal transport properties, electrical conductance, transport number and salt diffusion coefficient may be expressed as functions of mobility coefficients,22 eqn (8)-( 10) where A and h: are the equivalent conductance (cm2 f2-l equiv-I) and " specific " conductance ohm-l cm-l respectively. N is the equivalent concentration (equiv dm-9 and A = k - 1 0 3 1 ~ = a ~ 2 1 0 3 p (8) is the function, Z z L a a + Z z & b + z a z b ( L a b +&a). t$ = t," = (Z:&-,afZaZ&ab)/a t: = ti = (ZzLaa+ZaZ&ba)/a.( 9 4 (9b) Superscripts h and c distinguish transport numbers determined from Hittorf experi- ments and concentration cell e.m.f. measurements, respectively. Finally, where y is the stoichiometric mean molal activity coefficient and r = ra+r,, the sum of the stoichiometric coefficients for the salt. Mobility coefficients were obtained from these three equations, assuming, on the basis of the equality of tk and ti (table 6) that L a b = Lba and so the Onsager reciprocal relationships are obeyed, table 5. The activity term of eqn (10) was obtained from osmotic coefficients, #, tabulated by Robinson and Stokes,23 using the identity (1 1) Polynomials expressing In 6 and 6 as power series of In m were used to obtain the differential in eqn (11).Coefficients for these polynomials are given in table 1, together with the concentration ranges over which they were used. The estimated uncertainties in the activity terms calculated from these data are 0.5 %. Resistance coefficients R,,, R b b and Rab were obtained from the corresponding set of mobility coefficients by matrix inversion. The additional resistance coefficients RaO and RbO relating to the interaction between the ions and water (symbol 8) were obtained from the identity 24 (I+%) = d ( 4 4 -&-. CJ,, = 0, k = 0, a, b. i = 0,a.bA . AGNEW AND R. PATBRSON 289 1 The three mobility coefficients expressed as LiR/N are shown as functions of J N in fig. 1 and 2. There are few analyses of dissociated 2 : 1 electrolytes for comparison.Only Miller’s data 22 for barium chloride has been made over a sufficient range ta make comparisons meaningful. The anomalous concentration dependences of mobility coefficients associated with self-complexing behaviour are absent.2 * Zinc perchlorate is normally considered to be a highly solvated salt. This view is sup- ported by such evidence as the abnormally large mean molal activity coefficients in concentrated solution (at 4.0 mol kg-l the activity coefficient is 37.9).23 There are however, few indications of such effects in the analysis of mobility coefficients, table 5 . The only analogy which might be made is between the concentration dependence of Lab/N for zinc perchlorate and lithium chloride.22 The latter, which is the most hydrated member of the alkali chloride series, shares with zinc per- chlorate, low values of Lab/N (designated LI2/N in Miller’s paper)22 and the fact that for both Lab/N passes through a maximum and decreases markedly thereafter with increasing concentration. 1 2 3 1/N FIG.1 .-(a) Intrinsic mobilities (Laa/N) for zinc (e) and barium ions (-) in zinc perchlorate and barium chloride solutions, respectively. (b) Coupling coefficients (&b/N) for perchlorate (e) and In an earlier paper 2 5 it was shown that solvation effects in the alkali chloride series were more apparent when the ion-to-water frictional coefficients h o / [ Z i 1 were examined. This function, defined by eqn (12), measures the friction between one equivalent of charge of the ion i and those moles of water around it per unit volume chloride (-) ions in zinc perchlorate and barium chloride solutions. (Dimensions : see table 5).(12) A 0 CiRio lzil Izit - = -- Eqn (12) allows in principle the comparison of water friction for ions of different valencies. Using the data of table 5fa0/lza\ andfbo/lZbl are compared with those of barium chloride 22 and lithium chlorideYz2 fig. 3. It is obvious that the water2892 TRANSPORT I N AQUEOUs SOLUTIOXS '. I 1 2 FIG. 2.-Intrinsic mobilities (Lbb/N) for perchlorate (0) and chloride (-) ions in zinc perchlorate and barium chloride. (Dimensions : see table 5). 0 1 2 3 - z/N FIG. 3.-Ion-to-water frictional coefficients, fio/lZil for : zinc perchlorate, @, 2:'; +, ClO? : barium chloride, -, Ba2+ ; 8, Cl- : lithium chloride, - -, L{ ; + + , C1-.A .AGNEW AND R. PATERSON 2893 friction of zinc ion is strongly concentration-dependent and rises steeply with in- creasing concentration and is larger than that of barium or even lithium ions in their respective chlorides (when compared at equal equivalent concentrations). Per- chlorate ion has a relatively low ion-to-water frictional coefficient at infinite dilution, as reflected by its high equivalent conductance, j1&,, eqn (13)26 Such anions might well be expected to continue to have low water friction even in concentrated solution. In barium chloride and lithium chloride the chloride-to- water frictional coefficient actually decreases initially from its value at infinite dilution, before rising in more concentrated solutions due to the effects of the solvent order producing cation.In zinc perchlorate, perchloratelwater friction increases from infinite dilution, with no obvious minimum and rises steeply with concentration in a manner more marked than for chloride in either lithium or barium chlorides. It appears that the dominant solvent effect is that of zinc ion which increasingly affects the bulk solvent as salt concentration is increased and consequently makes the movement of perchlorate through that solvent more dif5cult. ap = i03~2~(fi0~~zi~); i = cio,. (13) TABLE 4.-ISOTHERMAL TRANSPORT DATA FOR ZINC PERCHLORATE AT ROUNDED CONCENTRA- TIONS C m A (*) /mol dm-3 /mol kg-1 /cm2 8 - 1 equiv-1 ts Dv/cm2 s- 1 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 .o 1.5 2.0 2.5 3.0 O.OO0 0.1009 0.2032 0.3070 0.4122 0.5189 0.6271 0.7369 0.8485 0.9617 1.0766 1.6801 2.3368 3.0569 3.8512 120.16 87.57 81.65 77.65 74.52 71.66 68.99 66.45 64.03 61.73 59.52 49.21 39.67 30.83 22.78 0.439 0.409 0.389 0.377 0.368 0.360 0.354 0.348 0.342 0.337 0.332 0.331 0.294 0.280 0.271 1.1818 1.0357 1.061 7 1.0966 1.1342 1.1752 1.2104 1.2476 1.2838 1.3187 1.3521 1.4903 1.5635 1.5480 1.4198 1 ,m 0.9943 1.0825 1.1942 1.3179 1 A496 1.5878 1.7321 1.8820 2.0373 2.8999 3.9861 5.0937 6.3083 C and m axe, respectively, molar and mold concentrations.A, ta and D, are equivalent conduc- tance, zinc transport number and volume-fixed salt diffusion coefficient, respectively. The activity term (1 +d In y/d In m) was obtained from the osmotic coefficient data of Robinson and Stokes 23 using eqn (1 1) and the polynomials of table 1.ONSAGER RECIPROCAL RELATIONSHIPS The equality, within experimental error, of the transport numbers for zinc, t: and t,", has been used above as sufficient reason to assume the validity of the Onsager reciprocal relationships, using eqn (9a,b). The experimental difficulties related to electrode performance in the Hittorf experiments made the uncertainty in zinc transport number ( & 0.01) and accordingly the concentration cell data Stokes and Levien were used for the calculation of irreversible thermodynamic coefficients (table 4). Zinc transport numbers from the Hittorf experiments, tk summarised in table 3, are compared with those of Stokes and Levien,13 t:, in table 6.2894 TRANSPORT I N AQUEOUS SOLUTIONS From eqn (8) and (9) and ~ 1 0 - 3 From table 6, it is obvious that the difference (Lab-Lba)/N is much smaller than the value of L,b/N, interpolated from the data of table 5.The ratio Lab/& is shown to be unity within the uncertainty of the experimental data. It is noteworthy (not- withstanding the inherent inaccuracies of the Hittorf determination of transport number) that zinc perchlorate is a most unfavourable system for a precise proof of the reciprocal relationships. The probable deviation of the ratio .&/Lab from unity TABLE 5.IRREVERSIBLE THERMODYNAMIC COEFFICIENTS FOR ZINC PERCHLORATE C L s a l N LablN LablN NR3a - N R m NRbb -CoRa -cORbO CO lmol dm-3 x 1012 x 1012 x 1012 x 10-12 x 10-11 X 10-11 X lO-1p x 10-11 /mol dm-3 0.0 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 .o 1.5 2.0 2.5 3.0 1.41 8 1.050 0.977 0.910 0.849 0.793 0.744 0.699 0.659 0.622 0.490 0.357 0.269 0.196 0.000 0.394 0.383 0.346 0.312 0.276 0.247 0.223 0.202 0.184 0.106 0.088 0.075 0.061 7.235 6.145 5.961 5.751 5.551 5.343 5.147 4.971 4.799 4.637 3.748 3.183 2.534 1.906 0.705 0.976 1.050 1.125 1.203 1.284 1.365 1.451 1.536 1.626 2.052 2.820 3.745 5.147 O.OO0 0.625 0.674 0.678 0.677 0.664 0.655 0.650 0.645 0.644 0.580 0.778 1.107 1.647 1.382 1.667 1.721 1.780 1.840 1.906 1.974 2.041 2.111 2.182 2.685 3.167 3.979 5.301 3.527 4.255 4.574 4.949 5.337 5.758 6.171 6.602 7.037 7.485 9.681 13.319 17.619 24.088 1.382 1.355 1.384 1 .a 1 1 SO1 1.574 1.647 1.716 1.788 1.860 2.335 2.773 3.426 4.477 55.35 54.62 54.25 53.87 53.49 53.1 1 52.73 52.34 51.95 51.56 49.56 47.51 45.40 43.24 N = 2C is the equivalent concentration, (equiv dm-3) and Co is the molar concentration of water (mol dm-9.The dimensions of Lik and Rik coefficients are molZ J-' cm-' s-' and J cm s mo1-2, respectively. TABLE 6.-TEST OF ONSAGER'S RECIPROCAL RELATIONSHIPS FOR ZINC PERCHLORATE concentration transport numbers for zinc (Lat-Lba)/N Lab X IN error limit Clmol dm-3 rt a tp b x 1012 1012 LbafLab in ratio 8LbalLb 0.0973 0.3223 0.4608 0.5743 1.01 54 1.1557 1.4636 1.9656 2.4247 0.406 0.372 0.361 0.359 0.336 0.325 0.314 0.297 0.279 0.409 0.374 0.363 0.355 0.330 0.324 0.312 0.295 0.282 0.014 0.008 0.008 - 0.015 -0.019 - 0.003 - 0.005 - 0.004 0.005 0.254 0.374 0.325 0.283 0.180 0.155 0.110 0.086 0.075 0.95 0.98 0.98 0.95 1.11 1.02 1.05 1.05 0.93 0.16 0.11 0.12 0.13 0.17 0.19 0.24 0.25 0.23 a t," are Hittorf data from table 3.b tf are interpolated data from concentration cell measure- m e n t ~ . ~ ~ = 8Lba/Lab are based on the probable uncertainty of 0.01 units in t," where G(Lba/Lab) 5% (at: + at:)K/LabZaZbFZ and K is the electrical conductance.A . AGNEW AND R. PATERSBN 2895 is almost entirely due to errors in transport number measurements, as Miller 27 has shown, eqn (16) where IC is the electrical conductance cm-l). This equation shows that the sensitivity of the test is increased by both low conductances and large coupling coefficients. For zinc perchlorate the reverse is true, conductances are large (table 4) and coupling coefficients are abnormally small, particularly in concentrated solutions, table 5. Zinc perchlorate is, therefore, a most unfavourable system for a precise proof of the Onsager reciprocal relationships.Self complexed salts, such as cadmium and zinc chlorides,6 which are typified by low electrical conductances and large coupling coefficients provide much more sensitive tests of the reciprocal relationships ; the latter is discussed in the succeeding paper. a(L,a/Lab) 53 (dt;+dt,") "/LabZaZbF2 (16) R. Paterson, J. Anderson and S. S. Anderson, J.C.S. Furaduy I, 1977,73, 1763. R. Paterson, J. Anderson, S. S. Anderson and Lutfullah, J.C.S. Faraday I, 1977,73, 1773. R. Paterson and Lutfullah, J.C.S. Faruday I, 1978, '74,93. R. Paterson and Lutfullah, J.C.S. Faraday I, 1978, 74, 103. A. J. McQuillan, J.C.S. Faruduy I, 1974, 70, 1558. A. Agnew and R. Paterson, J.C.S. Furuday I, 1978, 74, Part 6, 2896. C. 0. Quicksall and T. G. Spiro, Inorg. Chem., 1966,5, 2232. B. Gilbert, Bull. SOC. chim. belges, 1967, 76,493. 35. ' D. E. Irish, B. McCarrol and T. F. Young, J. Chem. Phys., 1963, 39, 3436. l o J. Beer, D. R. Crow, R. Grzeskew and I. D. M. Turner, Inorg. Nuclear Chem. Letters, 1973, 9, l 1 D. E. Irish, A. R. Davies and R. A. Plane, J. Chem. Phys., 1969,§0,2262. l2 M. M. Jones, E. A. Jones, D. F. Harmon and R. T. Seames, J. Amer. Chem. SOC., 1961, 83 l 3 R. H. Stokes and B. J. Levien, J. Amer. Chem. SOC., 1946, 68, 333. l4 C. W. Davies and G. 0. Thomas, J. Chem. SOC., 1958, 3660. l5 J. L. Dye, M. P. Faber and D. J. Karl, J. Amer. Chem. SOC., 1960, 82,314. l6 V. Frei and J. Podlahova, Chem. Z., 1963,67,47. l7 Lutfullah, H. S. Dunsmore and R. Paterson, J.C.S. Faraday I, 1976, 72,495. l 8 D. MacInnes and M. Dole, J. Amer. Chem. SOC., 1931,53, 1357. l9 M. J. Pika1 and D. G. Miller, J. Phys. Chem., 1970,74, 1337. 2o M. J. Pikal and D. G. Miller, J. Chem. Eng. Data, 1971, 16,226. 21 H. S. Harned and D. M. French, Ann. N. Y. Acad. Sci., 1945,46,267. 22 D. G. Miller, J. Phys. Chem., 1966, 70, 2639. 23 R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworth, London, 2nd edn, 1968). 24 L. Onsager, Ann. N. Y. Acad. Sci., 1945,46, 241. 2 5 H. S. Dunsmore, S. K. Jdota and R. Paterson, J. Chem. SOC. A, 1969, 1061. 26 S. K. Jalota and R. Paterson, J.C.S. Faraday I, 1973, 69, 1510. 27 D. G. Miller and M. J. Pikal, J. Solution Chem., 1972, 1, 11 1. 203 8. (PAPER 8 /703)

ISSN:0300-9599    

DOI:10.1039/F19787402885

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Transport in Aqueous Solutions of Group IIB Metal Salts at 298.15 K Part 6.-Irreversible Thermodynamic Parameters for Zinc Chloride and Verification of Onsager’s Reciprocal Relationships BY ANDREW AGNEW AND RUSSELL PATERSON* Department of Chemistry, University of Glasgow, Glasgow G12 SQQ Received 13th April, 1978 The isothermal vectorial transport properties of zinc chloride (conductance, transport number and salt diffusion coefficients) have been measured over the concentration range 0.1-3.5 mol dm-3 at 298.15 K. These data have been used to calculate mobility (&) and resistance (&) coefficients, using irreversible thermodynamic theory. Hittorf and concentration cell estimates of transport number have been used to verify the Onsager reciprocal relationships. The effects of self complexing in zinc chloride are illustrated by comparison with zinc perchlorate and other group IIB metal halides for which data are available. There is a scarcity of irreversible thermodynamic analyses on the transport properties of 2 : 1 salts, consequently there are few comparisons which can be made between pairs of salts, to illustrate the effects of self-complexing. Previous papers on cadmium iodide 1-5 showed unusual concentration dependence of transport coefficients.Because this salt is very strongly complexed, these coefficients could be predicted, at least in dilute solution, from its stability constants and by application of Onsager’s limiting laws expressed in irreversible thermodynamic form.6 Zinc chloride is a much less complexed salt ’ and there is no possibility of predict- ing its transport properties by such theories. It is, therefore, instructive to compare its transport properties with those of zinc perchlorate, which is completely dissociated at concentrations up to 3.0 mol dm-3.Although conclusions must be qualitative, complexing enhances all three mobility coefficients (La,, Lbb and Lab = &,a) in con- centrated solutions, while in dilute solutions up to 0.2 mol dm-3 when zinc chloride is almost completely dissociated the two salts are very similar, as might be expected. Certain features of the concentration dependence of self complexed salts of the group IIB halides, most dramatically illustrated by cadmium iodide, still remain but to a reduced degree for zinc chloride and also cadmium chloride which was studied previously by McQuillan.Experimental measurements of electrical conductance, salt diffusion coefficients and transport numbers have been made in the range 0.1-3.5 mol dm-3 as 25°C. It is shown that transport numbers estimated by Hittorf and concentration cell measure- ments agree within experimental error and consequently that the Onsager reciprocal relationships (Lab = Lba) hold in this system. EXPERIMENTAL Zinc chloride solutions were prepared from spectroscopically pure zinc oxide and AnalaR hydrochloric acid by the method described previously. ’ The stoichiometry of the salt solution was verified by analysis for zinc and chloride to a precision of kO.05 % for each 2896A . AGNEW AND R. PATERSON 2897 element. Density measurements were made and relationships between density, molarity (mol dm-3) and molality (mol kg-I) expressed as polynomials, table 1.Electrical conductance, salt diffusion coefficients and transport numbers were measured in the range 0.1-4.0 mol d ~ r t - ~ by the methods described earlier for cadmium iodide and zinc per~hlorate.~ Conductances were precise to k0.07 % due largely to uncertainties in chemical analysis, Transport numbers were determined by both Hit torf and concentration cell measurements : the latter by the technique described in the previous paper,g the former using the cell devised by Pika1 and Miller lo and used for cadmium iodide studies.l Since both silver-silver chloride and zinc amalgam electrodes are availableY7 a pair of each of these electrodes was used in each half cell.The combinations of measured cell potentials thus available allowed transport numbers for both zinc and chloride to be obtained inde- pendently and thus the internal self consistency of calculation methods could be verified. Y X P C mlC C ~ " ~ 1 0 5 c+ 11~x105 c+ K ~ 1 0 3 c K ~ 1 0 3 c K ~ 1 0 3 c Ta EB Ta Tb Tb t+ C f4 In m f4 In m n i - 0 TABLE PO POLYNOMIAL EXPRESSIONS OF THE FORM Y = aiXi Y X P C m l c C ~ " ~ 1 0 5 c+ ~ " ~ 1 0 5 c+ K ~ 1 0 3 c Kx103 c K ~ 1 0 3 c Ta EB Ta EB # In m d In m ao 0.998 163 1.001 665 97 1.210 34 2.391 04 0.186 215 2.315 346 0.400 527 394 39.739 47 -0.0619 570 07 -0.591 500 235 1.067 168 72 0.368 596 453 0.804 233 821 0.888 220 402 a3 O.OOO781 1 - - 6.177 83 -0.468 191 146.196 3 32.095 9 4.688 44 206.950 41 1 152.171 747 - 124.410 775 -209.520 195 0.060 243 047 2 0.012 894 265 5 -0.675 845 636 a1 0.117 757 0.020 885 006 0 - 1.327 79 -3.164 110 196.497 2 171.946 4 70.617 46 11.373 037 3 -1.360 161 87 1.751 358 27 - 11.517 490 1 - 0.068 31 5 774 4 - 0.054 496 535 9 -0.554 994 617 - 4.657 07 - -0.332 509 - - 0,007 024 276 45 0.004 447 122 71 0.210 629 967 concentration a2 lmol dm-3 -0.007 632 4 0.005 566 622 84 3.916 32 2.240 799 -212.075 2 -119.111 1 - 27.374 62 19.414 685 4 - 94.555 407 7 -25.409 334 9 95.658 985 5 -0.155 357 189 -0.009 445 978 77 0.967 938 651 0.1-4.0 0.1-4.0 0.1-1.5 1.2-3.9 0.01-0.3 0.25-1.1 1 .O-4.7 0.025-0.096 (V) 0.096-0.115 (V) 0.025-0.096 (V) 0.096-0.1 15 (V) 0.09-3.65 1.6-4.5-(111) 0.3-1.6 (m) concentration range as /mol dm-3 - 0.1-4.0 0.1-4.0 - 1.294 86 0.1-1.5 - 1.2-3.9 - 0.01-0.3 - 0.25-1.1 - - 1.0-4.7 0.025-0.096 (V) 0.096-0.1 15 (V) - - - 0.025-0.096 (V) - 0.096-0.1 15 (V) 0.09-3.65 - 0.3-1.6 (m) - 1.6-4.5 (m) - The preparation of zinc amalgam electrodes has been described earlier.' Silver-silver chloride electrodes, of the thermoelectrolytic type, were prepared as recommended by Ives and Janz.ll Both types of electrode had bias potentials of < 0.02 mV when placed in zinc chloride solutions. Salt diffusion coefficients were determined by the Rayleigh interferometric method, described by Chapman and Newman.12 Details of experiments are similar to those described for zinc perchlorate studies and had experimental uncertainty of k0.3 %.2898 TRANSPORT IN AQUEOUS SOLUTlONS RESULTS AND DISCUSSION TRANSPORT NUMBERS Transport numbers determined by the Hittorf method, had a precision of kO.01 which reflected difficulties encountered using zinc amalgam electrodes in these solutions.g In several instances light deposits of zinc oxide or oxychloride were formed on the anode and in such cases only cathodic half-cell results are given, table 2.TABLE ~.-HITTORF TRANSPORT NUMBER DETERMINATIONS mi %I mb m: mr C/mol dm-3 lmol kg-1 (of solution) 0.157 69 0.423 03 0.624 59 0.879 73 1.360 69 2.020 64 2.494 53 2.745 07 3.367 80 WC 28.593 28.218 25.116 30.256 31.522 33.720 34.054 34.495 36.524 0.155 14 0.403 95 0.583 78 0.802 05 1.187 48 1.669 33 1.986 20 2.144 68 2.516 60 W" I g 26.507 28.242 25.640 29.41 9 31.985 33.121 35.681 - - 0.155 28 0.403 86 0.583 71 0.801 75 1.187 34 1.669 36 1.985 57 2.146 12 2.516 78 2PliO /(mol x lO-3)-1 1.371 1 0.231 89 0.273 40 0.354 76 0.518 82 0.465 57 0.370 69 0.372 85 0.406 14 0.155 07 0.403 99 0.583 95 0.802 24 1.668 53 1.986 67 2.516 57 - - ti 0.347 0.316 0.278 0.225 anodic - - 0.015 - 0.091 - 0.235 - 0.135 77 0.502 61 0.487 42 0.738 84 1.187 41 1.618 80 1.922 06 2.082 76 2.462 29 tl: 0.336 0.309 0.281 0.239 0.118 - 0.022 -0.109 -0.132 - 0.227 cathodic 0.176 35 0.304 17 0.678 60 0.867 45 1.721 70 2.050 98 2.572 65 - - 1: (av.) 0.342 0.313 0.279 0.232 0.118 - 0.01 8 - 0.100 -0.132 - 0.231 G 0.354 0.316 0.279 0.225 "3.116 - 0.024 - 0.105 --0.142 - 8.226 mi, nz&, m:, mc and rnt represent the original solution concentration, the final concentrations of middle compartment on cathodic and anodic sides, respectively and concentration in the cathodic and anodic compartments at the end of the experiment.These concentrations are expressed as mol kg-l of solution. t t are transport numbers obtained by interpolation from cell data using the poly- nomials of table 1. For comparison, and to verify Onsager's reciprocal relationships, transport numbers were determined by potentiometric measurements using concentration ce1ls.l In this instance the four electrode positions in each half-cell were occupied by pairs of zinc amalgam and silverlsilver chloride electrodes. Bias potentials between pairs of like electrodes, which were never > 0.02mVY could be monitored during con- centration cell measurements. Four independent cell measurements could be made in any experiment, listed below as cells A, B, C, D and E Ag+AgCl IZnC1, aql [ZnCl, aql AgCl+Ag Ag + AgCl JZnCl, aql Zn + Zn,Hg IZnC1, aql AgCl+ Ag Zn + Zn,Hg IZnCl, aql AgCl + Ag (A> (B) (C) c2 c1 c2 c1 CA. AGNEW AND R.PATERSON 2899 Zn+Zn,Hg IZnC1, aql IZnC1, aql Zn+Zn,Hg Zn + Zn,Hg IZnC1, aql Ag + Cl IZnCl, aql Zn + Zn,Hg. (D) (El C1 c2 C1 c2 In principle the potentials of all cell combinations might be obtained with con- centration C1 kept constant at 0.044 143 mol dm-3, while C, was varied from 0.098 to 4.4 mol dm-3. In practice, since large concentration differences are to be avoided, due to heats of mixing, the same net effect was obtained by covering the concentration range in stepwise fashion to obtain the data of table 3. Integral transport numbers Ta and Tb (for zinc and chloride, respectively) were obtained from eqn (1) and (2) ?'ABLE 3.-POTENTIALS OF ZINC CHLORIDE CONCENTRATION CELLS REFERRED TO A DILUTE SOLUTION OF 0.048 143 mol dmV3 0.041 44 0.098 36 0.215 70 0.252 08 0.434 59 0.694 70 0.742 33 0.863 91 1.031 31 1.356 25 1.361 68 1.821 16 2.009 26 2.020 21 2.524 92 2.981 51 2.992 8 3.402 39 3.558 94 3.631 69 4.447 36 0 0.016 08 0.032 80 0.035 95 0.048 24 0.058 29 0.059 84 0.063 23 0.067 68 0.074 35 0.074 60 0.083 63 0.087 15 0.087 35 0.097 53 0.107 26 0.107 03 0.116 54 0.119 27 0.122 03 0.142 95 0 0.009 72 0.018 76 0.020 63 0.026 63 0.030 73 0.031 50 0.032 58 0.033 90 0.035 03 0.035 26 0.036 20 0.036 40 0.035 96 0.034 77 0.033 62 0.033 54 0.032 04 0.032 78 0.030 99 0.026 61 0 0.025 93 0.051 62 0.056 71 0.075 01 0.089 33 0.091 56 0.096 11 0.101 69 0.109 77 0.110 13 0.120 09 0.123 79 0.123 57 0,132 52 0.141 10 0.140 79 0.148 86 0.152 04 0.153 26 0.169 61 - 0.3751 0.3635 0.3638 0.3550 0.3440 0.3440 0.3390 0.3334 0.3191 0.3202 0.3015 0.2940 0.2910 0.2624 0.2383 0.2382 0.2152 0.2156 0.2022 0.1569 I 0.6202 0.6355 0.6340 0.6431 0.6526 0.6535 0.6579 0.6656 0.6773 0.6774 0.6964 0.7040 0.7069 0.7359 0.7602 0.7602 0.7829 0.7844 0.7962 0.8428 - - - 0.995 0.359 0.360 0.999 0.345 0.347 0.998 0.343 0.347 0.998 0.314 0.319 0.997 0.266 0.271 0.998 0.269 0.262 0.997 0.236 0.238 0.999 0.187 0.188 0.996 0.110 0.113 0.998 0.108 0.110 0.998 0.015 0.016 0.998 -0.018 - 0.018 0.998 - 0.020 -0.019 0.998 -0.107 -0.107 0.999 -0.178 -0.178 0.998 -0.177 -0.177 0.998 -0.234 -0.234 1.000 -0.245 -0.247 0.998 -0.262 -0.262 1.000 -0.325 -0.326 For the data given, experimental values of T,+Tb average 0.998, with standard deviation 0.001, rather than unity.Differential transport numbers, t,C and tg were calculated using eqn (3) and(4)lO t," = Ta+ E B dTa/dEB tg = Tb+ E B dTb/dEB. (3) (4) For this purpose T, and Tb were expressed as polynomials of E B (table 1) and differ- entiated. Agreement between these two calculations is excellent, as shown by the comparison of t: from eqn (3) and (1 - ti) from eqn (4), table 3. Hittorf and con- centration cell transport numbers are equal, within the experimental uncertainty of kO.01 for Hittorf and k0.005 for concentration cell measurements, table 3, fig. 1. These data may be compared with those of Harris and Parton l3 obtained from2900 TRANSPORT I N AQUEOUS SOLUTIONS concentration cells (0.5-4.0 mol kg-l) and additional points in more dilute solution reported by Robinson and Stokes,14 fig.1. There is good agreement between our data and those literature sources, especially at concentrations above 1 .O mol kg-'. Only two of these data points are significantly in disagreement with our data. Those at 0.5 mol kg-l (the lowes'i data point of Harris and Parton) and at 1.0 mol kg-l. The transport number for zinc in zinc chloride becomes zero at 2.0 mol kg-l (1.89 mol dm-3) and thereafter becomes negative. This inversion occurs at - 2.5 rnol kg-I for zinc bromide and - 3.5 mol kg-l for zinc iodide, reflecting the more com- plexed nature of the chloride salt.Our own earlier work on complexation showed that zinc chloride was not strongly complexed in dilute solution in contrast with aqueous cadmium iodide l5 for which cationic transport number (ta) becomes zero at 0.3 mol dm-3.1 Electrical conductance and salt diffusion coefficients (0,) were obtained by the methods described above to precisions of k0.03 and kO.5 %, respectively, given in tables 1-4 and at rounded concentrations, table 5. IRREVERSIBLE THERMODYNAMICS As shown previously,2* l6 the incidence of self-complexing or ion association does not affect the formal representation of the linear phenomenological equations of the salt a,,b,,, eqn (5)-(8). Ja = LaaXa+LabXb ( 5 ) Jb = LbaXa+LbbXb xa = RaaJa + RabJb or in inverse form where Ja and Jb are the total flows of zinc and chloride (mol cm-2 s-l) and Xa and X, their conjugate forces defined as the negative gradients of electrochemical potential of the free ions (J cm-l s-l).The dimensions of mobility coefficients Llk are mo12 J-1 cm-l s-l and for resistance coefficient, Rik, J cm s mo1-2. For cadmium iodide, in which self complexing is extensive in dilute solution, the component coupling coefficients between complexes and free ions have been identified and esti- mated, using classical theories limited to dilute solutions.2* For zinc chloride these methods of estimation are no longer valid since the major features caused by self-complexing occur at concentrations far above the limits of usefulness of current predictive theories. The comparison of zinc chloride and zinc perchlorate and zinc chloride and the complexed halides of other group IIB metals serve, however, to illustrate the effects of complexing upon transport coefficients and may be coin- pared with cadmium iodide, for which predictive analysis is available.2 The transport data measured are presented at rounded concentrations in table 5.There is agreement between Hittorf and concentration cell determinations of transport numbers as shown above, table 2 and fig. 1. In consequence the Onsager reciprocal relationships are obeyed, L a b = (The uncertainty in this assumption is examined below.) Since concentration cell measurements provided more accurate data, these trans- port numbers (tz) are used in table 5. As Miller l7 has shown, the mobility coefficients Llk may be obtained from [eqn ( 5 ) and (6)].eqn (9) -' r sk - tft,"A rirkDv , i = a, b N 103F2ZiZk 103RT rr,z,(d In y/d In m); . k = a, b. + - (9)A . AGNEW AND R . PATERSON 290 1 To evaluate the activity term of eqn (9) the osmotic coefficients (4) of Robinson and Stokes l4 were used, eqn (10) 0.6 0.3 0.2 0.1 2 0.0 - 0.1 i I '. FIG. 1.- aqueous 0.5 1.0 2.0 l/c -Zinc transport numbers (fa) as a function of the square root of the molar concentration for zinc chloride: 0, data derived from concentration cells, table 3 ; A, data from Hittorf measurements, table 2 ; 0, literature data from ref. (13) and (14). To evaluate the activity term, 4 was expressed as a polynomial in lnm over two ranges 0.1-1.6 and 1.6-4.4 mol kg-l. The activity term was also evaluated in dilute solution from the mean molal activity coefficients obtained previously (up to 1.0 mol kg-l), table 1.Agreement between these two calculations was excellent for all but 0.1 and 0.2 mol kg-I. Since these were at the lower limit of osmotic data where curve fitting procedures might be expected to be suspect, activity coefficient estimates TABLE 4.-EXPERIMENTAL SALT DIFFUSION COEFFICIENTS, Dv, FOR ZINC CHLORIDE D,X 105/cm2 s-1 1.048 1.01 8 1.005 1.002 0.995 0.979 C/mol dm-3 0.0499 0.1748 0.2477 0.2796 0.3254 0.6137 D, x 105/cm2 s-1 0.967 0.975 0.971 1.007 1.012 1.042 C/mol dm-3 0.7442 0.7651 0.9096 1.283 1.312 1.811 D,X 105/cm2 S-1 1.050 1.127 1.206 1.204 1.248 1.271 C/mol dm-3 1.831 2.357 2.995 3.095 3.442 3.9072902 TRANSPORT I N AQUEOUS SOLUTIONS were considered the more accurate for these concentrations.At 1.5 mol dm-3 osmotic coefficients pass through a minimum and so at this concentration and at 2.0 mol dm-3 the activity term was estimated graphically from a plot of C$ against In m, rather than from the polynomials of table 1. Irreversible thermodynamic transport coefficients are given in table 6 and the additional frictional coefficients, -CoRao, -CoRbo and the resistance coefficient for water Roo, were obtained from the identity, eqn (1 1)18 CiRil, = 0 k = a,b,O. (1 1) i=a,b,O TABLE s.-ISOTHERMAL TRANSPORT DATA FOR ZINC CHLORIDE AT ROUNDED CONCENTRATIONS C /mol dm-3 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 .o 1.5 2.0 2.5 3 .O 3.5 rn /mol kg-1 0 0.1004 0.2012 0.3024 0.4042 0.5066 0.6095 0.713 0.817 0.922 1.028 1.569 2.135 2.728 3.352 4.008 A /cm2 R-1 equiv-1 129.15 89.31 80.43 73.33 67.77 62.56 57.89 53.72 49.98 46.64 43.65 32.74 25.91 21.09 17.48 14.67 ta* 0.409 0.360 0.350 0.336 0.320 0.303 0.284 0.264 0.243 0.221 0.198 0.084 - 0.020 -0.106 -0.177 - 0.245 Dv/crnz s-1 1.2090 1.031 1.005 0.993 0.986 0.982 0.980 0.979 0.980 0.982 0.985 1.01 6 1.070 1.140 1.201 1.248 (I*) 1 .ow 0.8500 0.8414 0.8359 0.8204 0.8045 0.7898 8.7768 0.7655 0.7558 0.7477 0.7800 0.8872 1.0339 1.1991 1.4363 * Transport numbers from concentration cell measurements were used, table 3.TABLE 6.-IRREVERSIBLE THERMODYNAMIC COEFFICIENTS FOR ZINC CHLORIDE C /mol dm-3 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 .o 1.5 2.0 2.5 3 .O 3.5 1.41 8 1.127 1.067 1.021 0.995 0.975 0.960 0.948 0.940 0.935 0.932 0.882 0.81 1 0.748 0.688 0.608 0.000 0.526 0.625 0.719 0.824 0.933 1.037 1.135 1.229 1.317 1.399 1.616 1.650 1.616 1.543 1.408 Lbb/N x 1012 8.200 7.187 6.871 6.670 6.598 6.550 6.527 6.520 6.523 6.537 6.558 6.451 6.138 5.738 5.295 4.778 NRaa x 10-12 0.705 0.919 0.990 1.060 1.122 1.188 1.258 1.333 1.412 1.494 1.579 2.095 2.720 3.417 4.187 5.192 - NRab x 10-11 0.000 0.672 0.900 1.143 1.401 1.691 1.999 2.322 2.659 3.009 3.369 5.247 7.312 9.623 12.200 15.302 NRbb x 10-11 1.220 1.441 1.537 1.623 1.691 1.767 1.850 1.938 2.034 2.136 2.244 2.864 3.595 4.453 5.443 6.603 -CoRao x 10-11 3.527 3.922 4.052 4.159 4.206 4.247 4.292 4.344 4.400 4.461 4.527 5.229 6.288 7.461 8.737 10.656 - CoRbo x 10-11 1.220 1.105 1.087 1.051 0.990 0.922 0.850 0.777 0.704 0.631 0.559 0.241 - 0.061 -0.358 - 0.657 - 1.048 CO /mol dm-3 55.35 55.29 55.18 55.06 54.93 54.79 54.65 54.49 54.34 54.17 54.00 53.06 52.00 50.87 49.68 48.48 N = 2C is the equivalent concentration, (equiv dm-3) and Co is the molar concentration of wate, (mol dm-3). The dimensions of Ljk and Rik coefficients are mo12 J-l cm-' s-l and J c m ~ r n o l - ~ r respectively.A .AGNEW AND R . PATERSON 2903 Mobility coefficients L,,/N, Lbb/N and Lab/N are shown in fig. 2, 3 and 4. For zinc chloride, self complexing becomes significant only above 0.1 mol dm-3. Com- parison of intrinsic mobility coefficients for zinc (L,,/N) in chloride and perchlorate salts shows that there is no significant difference between the two salts up to and including 0.2 mol dm-3, fig. 2. Above this concentration L,,/N decreases less steeply in the chloride salt.In cadmium iodide and cadmium chloride * La,/N passes through a minimum. For zinc chloride this minimum is absent, but an 1.5 1.0 0.5 0.0 0.0 0.5 1.0 1.5 2.0 2.5 3.0 1/N FIG. 2.-Comparison of the intrinsic mobilities of zinc ions (Laa/N) in zinc chloride, 0 ; andlzinc perchlorate solutions, 0. The dimensions of Laa are given in table 6. (33 05 1.0 1.5 2.0 2.5 dN FIG. 3.-Intrinsic mobilities for chloride ions (Lbb/N) in zinc chloride, ; compared with those ob- served for chloride in barium chloride," - - - - - , * in cadmium chloride8 (insert, 0 ) ; and with perchlorate ion in zinc perchlorate, 0 .' The dimensions of Lbb are given in table 6. 1-922904 TRANSPORT I N AQUEOUS SOLUTIONS inflexion remains. The minimum in L b b / N , common to these self-complexed salts is still observed with chloride in zinc chloride, fig.3, although it is once more less pronounced than for the cadmium salts cited. Zinc perchlorate behaves as a fully dissociated salt and variation of its transport parameters with concentration is unremarkable and similar to barium chloride for which analysis is available,17 fig. 3. For the coupling coefficient L a b / N , comparison with barium chloride and zinc perchlorate (fig. 4) shows that in dilute solution its concentration dependence in zinc chloride is similar to a dissociated salt. As coinplexing becomes important, con- tributions to the coupling coefficient from complexes give additional increases, which are clearly obvious when comparison is made with zinc perchlorate and barium chloride, fig.4. It is this additional increase in Lab, due to contributions from the direct mobility coefficients of the complexes,2* which raises L a b until at above 2.0 mol dm-3 [zaZbLa,I > ZZL,, and so the zinc transport number becomes negative, eqn (12)17 where z,”Laa + ZaZbLab 01 t: = dN FIG. 4.--Coupling coefficients Lab/N for zinc chloride, @ ; zinc per~hlorate,~ 0 ; and barium chloride,” - . - - -. The dimensions of Lab are given in table 6. In part 5 increasing ion-to-water frictional coefficients were observed for both zinc and perchlorate ions as concentration was increased. In zinc chloride the zinc-to-water frictional coefficient (fa* / 12, I) increases less markedly than in zinc perchlorate while the chloride-to-water frictional coefficient, in contrast to perchlorate, decreases steadily and ultimately passes through zero and becomes negative, following the sign inversion for ta, eqn (13)16 The remaining independent friction, that between cation and anion, &,/[&I = fba/lZbl = - N & / I z a z b I , is much larger in zinc chloride than in the perchlorate salt.A .AGNEW AND R . PATERSON 2905 Again as with Lab/N the magnitude of this parameter is greatly enhanced by com- plexation. That these effects are due to increasing complexation of the salt can be seen by comparison of interionic frictional coefficients in cadmium chloride * and cadmium iodide.2 TABLE 7.-TEST OF ONSAGER'S RECIPROCAL RELATIONSHIP FOR ZINC CHLORIDE Lab error limit (+It N N in ratio LblLbB - Lab - Lsb- Lba x 1012 x 10'2 C/mol dm-3 t,h* t: 0.1577 0.4230 0.6246 0.8797 1.3607 2.0206 2.4945 2.745 1 3.3678 0.342 0.313 0.279 0.232 0.118 - 0.01 8 - 0.100 -0.132 - 0.23 1 0.354 0.316 0.279 0.225 0.116 - 0.024 - 0.105 - 0.143 - 0.226 -0.108 -0.021 0.000 0.036 0.007 0.017 0.01 1 0.021 -0.008 0.590 0.856 1.065 1.300 1.575 1.651 1.616 1.582 1.440 0.94 0.97 1 .oo 1.03 1.01 1.01 1.01 1.01 0.99 0.22 0.12 0.09 0.06 0.04 0.03 0.02 0.02 0.02 * t,h are Hittorf data from table 2.f 8(Lab/&) defined in Part 5 are based on the probable experimental uncertainty of 0.015 units in ( t t - ti). 1 .o 2 .o 3 .O z/N FIG. 5.-Ion-to-water frictional coefficientsho/JZil :fao/lZal zinc-to-water friction in zinc chloride, 0 ; and zinc perchlorate, 0. fbo/zb, chloride-to-water friction in zinc chloride: ; and perchlorate-to- water friction in zinc perchlorate, 0.ONSAGER RECIPROCAL RELATIONSHIPS The equality, within experimental error, of transport numbers t t and t:, has been used as sufficient reason to assume the validity of the Onsager reciprocal relationships Lab = &a. It is, however, instructive to assess the degree to which this equality may be justified. The experimental uncertainties in tk and t: were kO.01 and k0.005, respectively.2906 TRANSPORT I N AQUEOUS SOLUTIONS Using eqn (14)-(17) of the previous paper,9 the data of table 7 were obtained. Onsager's reciprocal relationships are verified to much closer limits than for zinc perchlorate. The probable error limit on the ratio LablLba is now only a few percent even although the probable experimental error (st:+&;), 0.015, is similar to that assumed for zinc perchlorate (0.01). Even although this parameter is rather higher than might be expected due to difficulties with the zinc amalgam electrodes in Hittorf determinations the probable error limit d(Lab/Lba) is only 2 % in the most concentrated solutions. As noted previously, in zinc perchlorate where similar problems bear,9 this degree of uncertainty could only be achieved if 6(tg+ ti) were +O.OOl due to low coupling coefficients L,b/N and large conductivities. R. Paterson, J. Anderson and S. S. Anderson, J.C.S. Faraday I, 1977, 73, 1763. R. Paterson and Lutfullah, J.C.S. Furuduy I, 1978,74,93. R. Paterson and Lutfullah, J.C.S. Faraduy I, 1978, 74, 103. R. Paterson, Faraday Disc. Chem. Soc., 1977, 64, 304. M. J. Pikal, J. Phys. Chem., 1971, 75, 3124. A. J. McQuillan, J.C.S. Furaduy I, 1974, 70, 1558. A. Agnew and R. Paterson, J.C.S. Faraduy I, 1978, 74, 2885. lo M. J. Pikal and D. G. Miller, J. Phys. Chem., 1970, 74, 1337. Reference Electrodes ed., D. J. G. Ives and G. J. Jam (Academic Press, N.Y., 1961). T. W. Chapman, Ph.D. Thesis (University of California at Berkeley, 1967). ' R. Paterson, J. Anderson, S. S. Anderson and Lutfullah, J.C.S. Faruduy I, 1977, 73, 1773. ' Lutfullah, H. S. Dunsmore and R. Paterson, J.C.S. Furuduy I, 1976,72,495. l3 A. C. Harris and H. M. Parton, Trans. Faraduy Suc., 1940, 36, 1139. l4 R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworth, London, 1968). Lutfullah and R. Paterson, J.C.S. Faruduy I, 1978,74,484. l6 S. K. Jalota and R. Paterson, J.C.S. Faruduy I, 1973, 69, 1510. D. G. Miller, J. Phys. Chem., 1966, 70, 2639. L. Onsager, Ann. N.Y. Acad. Sci., 1945, 46, 241. (PAPER 8/704)

ISSN:0300-9599    

DOI:10.1039/F19787402896

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Ionic Oxides : Distinction Between Mechanisms and Surface Roughening Effects in the Dissolution of Magnesium Oxide BY ROBERT L. SEGALL, ROGER ST. C . SMART* AND PETER S. TURNER School of Science, Griffith University, Nathan, Queensland 41 1 1, Australia Received 27th April, 1978 The mechanism of the rate determining step (r.d.s.) for dissolution of well characterised, very perfect MgO “smoke” crystals, has been re-evaluated by studies of the dependence of log (rate) on pH. Surface potential barrier modification is the most likely r.d.s. when solution diffusion is not limiting. Electron microscopic studies of changes in surface structure of partly dissolved crystals have been related to rate changes during dissolution. There is a doubling of surface area per unit mass in the first 10 % of dissolution, due to initial attack at defect sites, after which the increase in surface area per unit mass (to more than ten times the initial area) is due to both decreasing particle size and surface roughening.The complex set of factors that affect oxide dissolution, particularly those con- cerned with the metal oxide and the electrolyte solution, have been reviewed by Diggle and recently, by Valverde and Wagner.2 A theoretical treatment of the kinetics of dissolution has been proposed by Engell.3 Whilst this treatment has been accepted for ionic oxides, the extension of its application to semiconducting oxides has been ~riticised.~ In fact, much of the previous experimental work has been carried out with oxides that may be regarded as semiconducting, e.g., Fe0,39 Fe304,3* CUO,~? COO,^ Cr203,5 Ni0.6p In these oxides the electronic properties of the metal oxide, in particular, charge carrier concentration and mobilities, and the extent of charge depletion layers at the surfaces, may be of prime importance in the dissolution rate.The simplest case in which to study dissolution kinetics is that of a predominantly ionic oxide where the charge characteristics are not rate determining. The major factors of importance are then, for the solid, atomic surface detail, surface morphology, bond strength and, for the solution, diffusion, pH and electrolyte concentrations. In this paper, we present new results for MgO which lead to a critical re-examina- tion of the reliability of experimental tests of the theoretical predictions. We have used well characterised, very perfect MgO, the surface structure of which is examined before and during dissolution.The effect of diffusion of protons in solution on the dissolution rates is confirmed. The importance of the solid surface, in addition to the factors identified in the previous theoretical treatment, is established and quantita- tive estimates of rate changes due to surface roughening are given. EXPERIMENTAL MATERIALS Magnesium oxide was made from pure magnesium turnings. These were ignited in a slow stream of air and the MgO “ smoke ” collected on clean glass slides E 2 cm from the point of ignition. The oxide was immediately transferred to a small closed glass phial and kept in a dessicator. A.R. grade nitric acid and deionised, distilled water were used in dissolution experiments.29072908 DISSOLUTION OF IONIC OXIDES DISSOLUTION STUDIES The dissolution rates were measured by introducing precisely known masses of MgO to a stirred solution of known pH between 2 and 7. The mass was adjusted (e.g., from 0.5-20 mg in 50 cm3 of solution) to give accurately measurable rates over the pH range. The pH was recorded during dissolution with a Townson digital pH meter and the output recorded on a chat recorder. The solution and powder were contained in a 200 cm3 Pyrex beaker con- taining a Teflon-covered magnetic stirrer. Vermilyea s has shown that there is no difference in dissolution rates with air or inert gas atmosphere above the solution. The initial pH was successively set at 2.0, 2.5, 3.0, 3.5 and 4.0, and rose as the oxide dissolved.The dis- solution rate for both the initial well-chaacterised MgQ, and the partially dissolved oxide, could then be calculated from the rate of change of pH. A single run produced rates for the initial and higher pH values. Experiments were conducted at 25°C. We encountered the same problems with sluggish response in the pH range 6 to 8, as those discussed by Vermilyea.8 This was not a major concern as our discussion relates to the pH range below 5 in which theoretical predictions may be tested. Above pH 5, solution diffusion obscures the mechanism of dissolution of the MgO. Dissolution rates were calculated by a method similar to that used by Vermilyea.8 TECHNIQUES Specimens of partially dissolved oxide were prepared for electron microscopy by dipping an uncoated gold specimen support grid into the solution, removing excess fluid on filter paper and immediately inserting the grid into the microscope.The initial, unattacked oxide was examined by both coating a grid during burning of the magnesium and ultrasonically dispersing some of the stored powder in A.R. petroleum ether and depositing onto a grid. Specimens were examined on a JEM lOOC microscope at 100 keV. Surface areas were measured by adsorption of nitrogen (a = 16.2 A2) at - 196°C using the continuous flow method of the Perkin Elmer Model 212D sorptometer. More than 5 adsorption/desorption cycles were measured. The surface area of initial oxide was found to be between 9.8 and 10.9 m2 8-l. RESULTS AND DISCUSSION CHARACTERISATION OF MAGNESIUM OXIDE MgO prepared by burning magnesium in air produces small cubic crystals with a very high degree of perfection (fig.1). The edge length of the crystals ranges from about 50 nm to over 200 nm. No dislocations are evident in any of the crystals ; they are clearly very well annealed compared with MgO crystals prepared by de- composition of the hydroxide, carbonate etc. This is borne out by the relatively low surface area (about 10 m2 g-l) compared with that formed from the hydroxide by decomposition at 600°C in air (67 m2 g-l). The atomic surface detail of this material has been extensively examined by Moodie and Warble,9 who heated small crystals to desorb gas and examined the surface by high resolution electron microscopy using phase contrast.Their micro- graphs show growth steps of one or two unit cell height, as isolated cubic projections with a volume of a few unit cells. The steps occur roughly within 50 A of each other. On the smoother surfaces, Moodie and Warble have estimated, in some cases, a surface roughness of only 0.05 using the definition of Burton et aZ.lo In contrast, their work shows that MgO formed by decomposition is considerably less perfect, with high index mean faces composed of closely spaced {loo} steps, dislocations and sintered intergrowth showing both perfect and imperfect alignment. The obvious sites for attack in the smoke crystals are the sharp edges, corners and the kink sites (i.e., ions with fewer nearest neighbour ions than the ions in a perfect { 100) surface) associated with the projections.FIG. 1.-Transmission electron micrograph of MgO smoke crystals.The edge length of the cubic crystals in the powder is in the range 50-200 nm. [To face page 2908FIG. 3. F~G. 4. FIG. 3.-Transmission electron micrograph of MgO smoke crystals after initial attack to < 5 total dissolution from initial pH 2.5. FIG. 4.-Transmission electron micrograph of MgO smoke crystals after 95 % dissolution from initial pH 2.5. ofPLATES A AND B.-Electron micrographs of collapsed micelles isolated by the spreading-drop technique from micelle solutions of the polystyrene-polyisoprene block copolymer in DMA. Specimens shown on both plates were stained in solution with Os04 and the specimen shown on Plate A was To firrepage 3355.1 in addition lightly shadowed with C/Pt.The scale marks are ( A ) 200 and ( B ) 400 nm.R. L. SEGALL, R. ST. C . SMART AND P. S . TURNER 2909 RATE CONTROL I N MECHANISM OF DISSOLUTION The theoretical treatment of oxide dissolution has been developed by EngelL3 Vermilyea and Digg1e.l Their predictions, for the experimental value of the slope of log (rate) against pH dependence, for different mechanisms in MgO dissolution are summarised in table 1. In a previous experimental study of the dissolution of MgO and Mg(OH)2, Vermilyea considers that, at low proton concentration between pH 5 and 7, proton diffusion in solution is the r.d.s. Our work has confirmed the importance of solution diffusion in rate measurements even down to pH 2. Stirring rates directly affected the measured rate in all experiments.Rates discussed below were always measured with stirring rates above the level at which no further increase in dissolution rate could be measured simply by increasing the stirring rate. At pH > 5, this could not be achieved; proton diffusion is clearly limiting. TABLE 1.-THEORETICAL PREDICTIONS OF THE SLOPE OF LOG (RATE) AGAINST pH DEPENDENCE rate potential barrier determining reaction of anions with protons4 modification by [H+].1 step OH-+ H+ -+ H20 02-+H+ + OH- 0 2 - +2H+ + H2O anion removal cation removal slope of log (rate) - 0.67 - 0.5 - 1.0 -0.5 - 0.5 against pH Vermilyea has proposed that, at low pH, MgO first reacts to form Mg(OH)2 and that dissolution of the hydroxide is then rate limiting. Below pH 5, for Mg(OH)2 he has found a slope of approximately -0.47 for a log (rate) against pH dependence.However, his curves were reproducible to w 50 % and results below pH 5 were only considered correct as to order of magnitude. Part of the reason for this un- certainty arises from surface area estimates. Vermilyea's MgO powders were obtained by sieving (10-30 pm fraction) ground, fused optical grade MgO crystals, sedimentation in alcohol, dried at 400°C and stored in a dessicator. The surface area was then calculated from the particle size assuming smooth surfaces, and presumably perfect cubes. The surface area was corrected for the amount dissolved whereas, as discussed below, the area per unit mass actually increases markedly during dissolution. TABLE 2.-DIsso~uTIoN RATES FOR MgO IN NITRIC ACID AT 25°C PH ratelmol cm-2 s-1 2.0 9.5 x 10-l0 2.5 6.3 x 10-lo 3.0 3 .0 ~ 10-lo 3.5 1.8 x 10-lo 4.0 1.05 x 10-lo We have used highly perfect MgO smoke crystals of known surface area and surface morphology. Table 2 lists initial rates at different pH from a number of runs between pH 2 and 4. These values are at least one order of magnitude lower than those found by Vermilyea. However, replotting these results as log (rate) against pH as in fig. 2, it can be seen that the error in the measurements, even with well characterised, relatively perfect oxide, is sufficient to give a slope between -0.4 and - 0.6 with a mean value of - 0.5. It is clear that dissolution is controlled by a surface mechanism, not solution2910 DISSOLUTION OF IONIC OXIDES diffusion, in this pH range.Vermilyea has suggested that it is surface reaction, i.e., a slow reaction of surface hydroxide with a second proton and, hence, expects a slope of -3 for the log (rate) against pH plot. He explains the lower value in his work (i.e., -0.47) as being due to a low value of the transfer coefficient a+ for the cation (i.e.y - 0.35 instead of 0.5). In support of Vermilyea, infrared studies 11* l2 have shown that the MgO surface is covered with both free, noninteracting hydroxyl and hydrogen bonded hydroxyl groups immediately after exposure to air but this does not constitute a brucite Mg(OH)2 structure. It is not possible to avoid surface hydroxylation without preparation and transfer in vacuum. It seems reasonable to assume that surface hydroxylation to form OH groups takes place virtually instantaneously in solution.If surface reaction (Le., reaction of OH- to form H20) is rate limiting, a slope of - 3 would, indeed, be expected on this assumption. -9.c -9.2 - 9.4 n c1 .s % -9.6 - -9.8 -10.0 I I I I - 2-0 2.5 3.0 3.5 4-0 P H is -0.49. FIG. 2.-Log (rate) against pH for MgO smoke crystals in nitric acid at 25°C. The slope of the graph Infrared studies ll. l 2 also show very rapid formation of water molecules on the surface. Electron microscopy, from our work and that of Moodie and Warbley9 shows rapid surface etching by water vapour. It seems more likely, since both protonation reactions are very fast, that the rate is limited in solution by modification of the surface potential barrier by the protons. This would suggest a slope of -0.5, in reasonable agreement with Vermilyea’s and our results.This result does not distinguish between rate control by cation or anion removal. The ion responsible for rate control would, however, have a transfer coefficient very close to 0.5. Again, in view of the very fast hydroxylation, it seems more likely that cation removal is rate limiting.R. L. SEGALL, R. ST. C. SMART AND P. S. TURNER 291 1 In reassessing distinction between different r.d.s. on the basis of log (rate) against pH data, it does not appear likely that, even for well characterised material, a definite prediction of reaction mechanism can be made. Part of the reason for this uncertainty lies in differences in surface properties between different preparations, and in initial surface attack.SURFACE CHANGES DURING DISSOLUTION Fig. 3 shows the initial stages of attack on the small cubes of MgO at < 5 % of total dissolution. The edges and corners appear to be shaved off leaving an average (110) facet. Steps are observed across these facets (e.g., at A) suggesting that material has been removed through dissolution at steps which move along the facets in the (100) directions. The { 100) faces of the original cubes show considerable roughening suggestive of the development of pits or channels (e.g., at B). Further attack leads to far less regularly shaped crystals (fig. 4 at 95 % dissolution) which show the development of distinct steps and ledges as dissolution proceeds. The surface of these particles is obviously very rough with a high concentration of kink sites at steps, ledges, corners etc.The surface detail from all micrographs obtained between 10 and 95 % dissolution is consistent with the model, demonstrated by Moodie and Warble for growth of MgO sinters, in which the various (hkl) crystal faces are made up of (100) steps down to the unit cell level, e.g., an average (110) face is actually a series of closely packed ( 100) steps. % dissolution FIG. 5.-Area ratio (expressed as part dissolved area per unit mass/initial area per unit mass) against percent of total dissolution. Full line from experimental results, broken line calculated from area ratio change with decreasing particle size during dissolution. The mechanism for dissolution is dependent on the surface structure in that protons diffusing to the surface attack preferentially at kink sites.It is possible, from the pH changes, to follow the change in surface area per unit mass during dissolution. For example, from an initial pH of 2 with initial oxide of known area, the rate at pH 2.5, 3.0, 3.5 and 4.0 may be calculated in mol g-l s-l,2912 DISSOLUTION OF IONIC OXIDES since the mass remaining at each pH can be calculated. These rates are then com- pared with initial rates, for the same initial oxide, at the respective pH value, i.e., in mol cm-2 s-l. This allows the ratio of the area of the partially dissolved oxide to the initial area to be computed as a function of percent of total dissolution. Results are shown in fig. 5. There is a doubling of the area per unit mass within the first 10 % of dissolution.This would appear to be associated with the expected, and observed (fig. 3), initial attack at edges, corners and kink sites associated with the surface projections. The area gradually increases until, at more than 70 % dissolution, the very irregular crystals give an increase in area to more than five times the initial area. At this stage the irregular crystals (fig. 4) provide a great variety of kink sites for preferential attack. In considering fig. 5, it is clear that, in addition to the effects of changing kink site density, there is an increase in surface area per unit mass associated with decreasing partick ,yize. For cubic particles, of initial mass M,, the area ratio (expressed as part-dissolved area per unit mass/initial area per unit mass) is given by (Mx/Mo)-* where M, is the mass remaining at x % dissolution. Assuming that the increase in surface area due to this decreasing particle size effect is additional to the doubling of area due to initial attack (within the first 10 % of dissolution) we have plotted the particle size area increase from 10 % onwards as the broken line in fig.5. Our results suggest, that after 10 % dissolution, the two factors causing an increase in surface area per unit mass, namely (i) particle size decrease and (ii) surface roughening (fig. 4) associated with production of an increasing density of kink sites, may be roughly equal in magnitude. Each effect separately appears to give a factor of roughly 1.5 after about 60 %. It is clear, from these results, that correction during dissolution for loss of surface area due to mass loss alone is not adequate. In determining dissolution rates for ionic oxides, the importance of atomic surface detail and particle size is at least as great as that of theoretical considerations of the mechanism of the solid/solution interface. Support from the Australian Research Grants Committee and the Australian Institute for Nuclear Science and Engineering is gratefully acknowledged. J. W. Diggle, Dissolution of Oxide Phases in Oxides and Oxide Films, ed. J. W. Diggle (Marcel Dekker, N.Y., 1972), vol. 2. N. Valverde and C . Wagner, Ber. Bunsenges. phys. Chem., 1976, 80, 330. H. J. Engell, 2. phys. Chem., 1956, 7, 158. D. A. Vermilyea, J. Electrochem. SOC., 1966, 113, 1067. N. Valverde, Bey. Bunsenges. phys. Chem., 1976, 80, 333. C . F. Jones, R. L. Segall, R. St. C . Smart and P. S. Turner, J.C.S. Faraday I, 1977, 73, 1710. C . F. Jones, R. L. Segall, R. St. C . Smart and P. S. Turner, J.C.S. Faraday I, 1978, 74, 1615. A. F. Moodie and C . E. Warble, J. Cryst. Growth, 1971, 10, 26. W. K. Burton, N. Cabrera and F. C. Frank, Phil. Trans. A, 1951, 243,299. R. S t . C. Smart, T. L. Slager, L. H. Little and R. G. Greenler, J. Phys. Chem., 1973,77, 1019. * D. A. Vermilyea, J. Electvochem. SOC., 1969, 116, 1179. l 2 J. V. Evans and T. L. Whateley, Trans. Faraday SOC., 1967, 63, 2739. (PAPER 8/786)

ISSN:0300-9599    

DOI:10.1039/F19787402907

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Photoluminescent Spectra of Surface States in Alkaline Earth Oxides BY SALVATORE COLUCCIA,? A. MICHAEL DEANE AND ANTHONY J. TENCH* Chemistry Division, AERE Harwell, Oxfordshire Received 15th May, 1978 Photoluminescence has been observed from high surface area alkaline earth oxides with exciting light of a much lower frequency than that expected from the band gaps of the bulk oxides. This Iuminescence can be quenched by oxygen and hydrogen and the quenching is reversible under some conditions. It seems probable that both the excitation and the luminescence spectra observed arise from excitons in the surface region. The different behaviour towards oxygen and hydrogen has been used to develop a model in which the excitation site, where absorption of light occurs, is thought to be associated with anions on the surface in states of unusually low coordination such as might be found at step, edge or corner sites on the surface. The luminescence is thought to be associated with a cation on the surface in a similar state of coordination, Pure MgO in single crystal form shows no absorption in the near ultraviolet but high surface area powders have been shown to give a luminescent spectrum when excited by U.V.light with an energy much less than the band gap of the bulk solid. The absorption and emission were attributed to the presence of local surface states, intrinsic to the oxide, which correspond to ions in sites of unusually low coordination or to the presence of high index planes on the surface. Molecular species such as hydroxyl groups on the surface can also give rise to absorption and luminescent spectra which are significantly different from those associated with the intrinsic ions.Similar surface states should be present on other oxides and it is likely that ions in unusually low states of coordination play a major role in determining the reactivity of the solid surface. With this in mind, we have studied the photoluminescent behaviour of the alkaline earth oxides, both in vacuu and in the presence of oxygen and hydrogen, to try to establish the chemical properties of the sites involved in the absorption and luminescent processes. The results from the excitation data are compared with reflectance meas~rements.~ EXPERIMENTAL The apparatus used was developed during this work from that described previously.' A 250 W xenon lamp provided the excitation source and excitation wavelengths were selected using two Spex f4 monochromators coupled to form a double monochromator.The studies were limited to 2 > 230nm (5.40eV) because of the very low intensity of the exciting light at shorter wavelengths. Emission spectra were analysed using a similar f4 monochromator. The excitation and emission band pass wits 5 nm for all the results reported here. When the emission wavelengths permitted, a sharp cut filter (Corning 3-74) was used before the emission monochromator to eliminate higher orders of the exciting wavelength. A small fraction of the monochromatic exciting light was detected with it photo- multiplier, type 95263 (silica window, bialkali cathode) and emission spectra were detected with a photomultiplier, type 9798B (glass window, S20 cathode). An Ortec photon counting t On leave from the Institute of Physical Chemistry, University of Turin, Italy.291 32914 SURFACE STATES I N OXIDES system was used to compare the photon pulse rates from the two photomultipliers and to present spectra corrected for variations of excitation intensity with wavelength or time. No corrections were made for the variation of photomultiplier sensitivity with wavelength but both multipliers were used in wavelength ranges where their photon sensitivity with wavelength was fairly uniform. Variable temperature measurements below ambient were carried out with the specimen cell mounted in a flow of cooled nitrogen gas, while for those above ambient the specimen was mounted in a small furnace equipped with two windows at right angles. Preliminary measurements of luminescent lifetimes were made on an Applied Photo- physics nanosecond spectrometer, model SP2.Lifetimes were measured using the mono- chromator to select an excitation wavelength and the integral emission over all luminescent wavelengths was collected on the photomultiplier. Magnesium oxide was prepared by controlled thermal decomposition of high purity hydroxide or carbonate by standard techniques '9 and calcium and strontium oxides from the Specpure carbonates to give surface areas of 100-200 m2 g-l for MgO, 50-70 m2 g-' for CaO and 10-20 m2 g-l for SrO. Most preparations were carried out on a vacuum line with a base pressure about Torr), but one preparation of magnesium oxide was carried out on a grease free system with a final pressure about N m-2 (lo-' Torr).A low area specimen of barium oxide was prepared by outgassing a specimen of Hopkin and Williams " fine chemical " barium oxide to 1200 K in vacuum. A further specimen was prepared by decomposing barium nitrate on magnesium hydroxide in the manner described by Zecchina and Stone.4 Each preparation was carried out in a silica ampoule attached to a rectangular Spectrasil cell 2mm thick and the powder shaken into the cell after preparation. The effect of hydrogen and oxygen on the luminescence was studied using Specpure gases with the specimen attached to the vacuum line while mounted in the spectrometer ; gas pressures quoted were measured after equilibration with the surface.N m--2 RESULTS LUMINESCENCE OF CLEAN OXIDE SURFACES All the alkaline earth oxides could be excited with ultraviolet light to show a broad luminescence band which in general moved to lower energy as the cation size increased (table 1, fig. 1 to 4). The excitation spectra show a doublet for all the oxides, however the maxima of the peak at higher energies could not be observed clearly for MgO and CaO because of the high noise levels below 230nm. The position and shape of the luminescence band varied significantly with the excitation TABLE 1 .-ABSORPTION,3 ' EXCITATION AND EMISSION SPECTRA OF ALKALINE-EARTH OXIDE POWDERS absorption excitation emission oxide lev lev lev MgO 5.70; 4.58 >5.40; 4.52 3.18 CaO 5.52; 4.40 >5.40; 4.40 3.06 SrO 4.64; 3.96 4.43 ; 3.94 2.64 BaO 3.60; 3.22 3.70 2.67 wavelength indicating that more than one component was present.The intensity of the luminescence and the ratio of the excitation peaks varied to some extent from one preparation to another and considerable care was needed to obtain reproducible results. The luminescent species were present after < 1 s irradiation in the excitation waveband and this was taken as evidence that the luminescence was not the result of a photochemical reaction. At the illumination intensities used, ultraviolet induced changes were limited to intensity variation of a few per cent ; spectra were normally recorded with the specimen in equilibrium with the excitation source.S. COLUCCIA, A. M . DEANE AND A .J . TENCH 2915 Some samples of MgO were prepared with Fe, Cr or Mn added (20 to 100 p.p.m. of impurity) to examine whether the luminescence described above was associated with impurity transition metal ions. All these samples had the same broad lumines- cence but in addition for the CrlMgO sample both the surface and bulk luminescence of Cr3+ ions was observed around 710 nm (1.75 eV). Surface contamination by carbon compounds does not seem to cause the luminescence since a specimen of magnesium oxide prepared under completely grease-free conditions showed a similar spectrum to that of other preparations. Deliberate attempts to contaminate the surface with carbon compounds resulted in a loss of the luminescence spectrum. eV 60 5.0 4.0 3.0 2 .o ___ -- , ’ IT---- 7 I 200 300 400 500 600 wavelength /nm cence, (6) emission excited at 230 nm, (c) emission excited at 274 nm.eV FIG. 1 .--Photoluminescence spectra of MgO at 300 K. (a) Excitation spectrum of 390 nmllumines- 6.0 5.0 4 .O 3 .O 2 .o r----’i” ’ 1 I 200 300 400 500 600 wavelength/nm FIG. 2.-Photoluminescence spectra of CaO at 300 K. (a) Excitation spectrum of 405 nm lumines- cence. Emission excited at (b) 282, (c) 310 and (d) 330nm.2916 SURFACE STATES IN OXIDES Several specimens of MgO and SrO were prepared with the surface area reduced by sintering in oxygen or water vapour at various temperatures and then evacuating at 1200 K. Some scatter in intensity was observed between the different preparations, but sintering the same sample progressively in air or O2 leads to a reduction in the intensity of the emission and after long sintering periods the reduction was very marked.The luminescence of the oxides showed strong temperature dependence ; that of magnesium oxide showed an order of magnitude decrease with temperature from 150 to 400 K. Since the emission is not a single peak the data do not justify further analysis at present. eV 60 5.0 4 0 30 2 .o wavelength/nm FIG. 3.-Photoluminescence spectra of SrO at 300 K. (a) Excitation spectrum of 470 nm lumines- cence. Emission excited at (6) 280, (c) 315, ( d ) 330 and (e) 350 nm. eV 6.0 5.0 4 0 30 2.0 i 200 I_-- i I 300 LOO 500 I 6 00 wavelength /nm FIG. 4.-Photoluminescence~spectra at 300 K of BaO supported on MgO. (a) Excitation spectrum of 465 nm luminescence. Emission excited at (b) 270, (c) 330 and (d) 360 nm.S .COLUCCIA, A . M. DEANE AND A . J . TENCH 2917 Luminescent lifetimes were calculated for MgO and SrO on the assumption that the observed decay curves were the sum of several first order decay curves (table 2). The apparatus was not suitable for the determination of lifetimes $= 100 p, and a specimen of polycrystalline barium hydroxide, known to have a 2 s lifetime showed only a background of uniform luminescence over the 1 0 0 ~ s time span of the apparatus. The excitation spectrum of BaO coated MgO [fig. 4(43 showed a band at 270 nm (4.59 eV) very similar to that of the supporting material together with a second band at 335 nm (3.70 eV) characteristic of BaO. The band at 270 nm (4.59 eV) coincides with the excitation band of MgO though comparison with the BaO powder indicates that the BaO could be responsible for some contribution.However, the emission is very different from that of MgO (fig. 1) and is attributed to the BaO on the surface. TABLE 2.-LUMINESCENCE LIFETIMES OF SURFACE STATES ON MAGNESIUM AND STRONTIUM OXIDE excitation lifetime oxide energy/eV IPS MgO 4.52 1, 7, 25 SrO 4.40 40, 100 SrO 3.94 22, 50 EFFECT OF ADSORBED GASES ON LUMINESCENCE All adsorbate gases tested (02, H2, C02, H20 and organic vapours) quenched the luminescence in the oxides in the pressure range 1 to 100 N m-2 to 1 Torr), but quenching by oxygen and hydrogen was observed at much lower pressures without significant modification of the surface. In contrast, the other gases react chemically with the surface and in addition to quenching the intrinsic surface luminescence, several adsorbates formed new luminescent species which will be considered separately.6 The reversible quenching of luminescence from MgO by oxygen has already been described 1’ but further measurements have shown that there is a small irreversible quenching at room temperature; the intensity was restored to only about 80 % of its original value after evacuation.The initial quenching by oxygen produced similar reductions of intensity in each of the excitation bands, but after exposure to high pressures of oxygen (400 N m-2, 3 Torr) at 295 K under U.V. irradiation the higher energy band was irreversibly reduced to half its original value. At 295 K quenching to half the original intensity required 0.7Nm-’ (5x Torr) of oxygen but at 550 K the same proportional reduction required 700 N m-2 (5 Ton) of oxygen while at 575 K oxygen quenching was not detected under 700 N m-2 of oxygen.The luminescent spectra from CaO and SrO were quenched by oxygen, the amount of irreversible quenching increasing along the series MgO, CaO and SrO. The effect of H2 on the luminescence of clean MgO is dependent on whether the sample is left in contact with the gas in the dark or under U.V. irradiation. In the former condition, contact with 10 Torr H2 for 10 min produces only minor changes in the original spectra. However, if H2 is admitted on MgO under U.V. irradiation the luminescence is quenched reversibly only if the gas is pumped off immediately, but irreversibly if left to stand for a few hours.During the quenching experiments small variations (about 5 nm) were noticed in the position of the excitation maximum. These changes give some indication that the excitation peak contains at least two unresolved components which are present in varying proportions. The spectrum of strontium oxide showed the variations unambiguously; a more detailed study of that oxide has been made.291 8 SURFACE STATES I N OXIDES The excitation band of strontium oxide shows considerable changes in shape when hydrogen is adsorbed (fig. 5). A more quantitative picture of this was obtained by resolving the envelope into five component gaussian bands. These bands remained unchanged in position and width during the quenching experiments ; however, their relative contribution to the spectrum changed markedly after addition of hydrogen.Those bands on the low frequency side of the envelope spectrum are preferentially eroded by hydrogen and this is irreversible at 295 K. There was no change of the luminescence band shape except a slight narrowing (105 to 90 nm) on the first hydrogen addition but a progressive shift of the band to shorter wavelengths was found as quenching proceeded. The addition of oxygen after the excess hydrogen had been pumped off caused a further reduction of intensity, and an increase of luminescence band width to its original value, but no frequency shift. After reactivation at 1200 K the luminescence band was restored close to its original value and the differences between the original clean surface and the reactivated surface were typical of the differences observed between samples prepared separately.eV 50 L O I 3 wavelength/nm FIG. 5.-Excitation spectra of SrO at 300 K. (a) SrO in vacuum, (b) SrO in contact with 1 N m-2 (8 x Torr) of Hz. In contrast to the effect of hydrogen, quenching of the luminescence from strontium oxide by oxygen at low pressure caused no changes of shape in the excitation band as the intensity decreased. The luminescent band position and shape were also unaltered, The luminescence spectrum was asymmetric along the linear wavelength abscissa and replotted on a linear energy abscissa, the luminescence excited at 3 15 nm (3.94 eV) was a single symmetrical gaussian with a small long wavelength tail which never exceeded 5 % of the height of the main band. The behaviour of this part of the spectrum when quenched showed that it was caused by a separate weak lumines- cence centred about 625 nm (1.98 eV).Because of its low intensity this second luminescence band has not been studied further. DISCUSSION All the alkaline earth oxides studied show a characteristic luminescent spectrum when excited by light in the near U.V. This is not found for the pure single crystal material but is characteristic of the relatively high surface area (20-150 m2 g-l) materials studied. Similarly, apart from BaO the first excitation levels of the bulkS . COLUCCIA, A . M. DEANE A N D A . J . TENCH 2919 crystals (table 3) are at much higher energy than the observed excitation bands of the luminescent spectra.These factors lead us to suppose that the observed luminescence and corresponding excitation spectra are associated with energy levels differing from those present in the bulk material and of major significance in the small crystallites because of the large ratio of surface to bulk ions. It has been previously argued for MgO that these new levels do not result from impurities but are associated with the reduced coordination of ions at the surface; the similarity in behaviour of the other alkaline earth oxides supports this picture. This general model is confirmed by the sensitivity of the luminescence to a wide range of gases. The quenching induced by these gases is reversible in some cases but in others a reaction takes place with the surface and the luminescence is destroyed irreversibly at that temperature, indicating the enhanced reactivity of those ions with a reduced coordination.The luminescence process consists of two parts, the initial energy absorption and the subsequent emission of light ; both of these need to be considered in some detail before going on to develop a model. The excitation spectra observed for the crystallites in vacuo are essentially identical with the surface absorption bands reported by Zecchina and Stone 3 9 (table 1) although the latter measurements were made under conditions where the luminescence had been quenched by an oxygen overpressure of 133 N m-2 (1 Torr). The absence of any significant difference between the absorption and excitation bands indicates that the adsorption of sufficient oxygen to quench the luminescence does not modify the levels responsible for the absorption of light.It is clear from the reflectance data that at higher pressures, particularly in the case of BaO, modification of the surface energy levels does occur but this appears to be associated with higher surface coverages where strong adsorp- tion occurs. The evidence indicates that absorption and emission of light may take place at different sites on the surface and it is particularly interesting to look at the effect of hydrogen on the luminescence. Quenching of the emission was found for several of the oxides; the detailed study of SrO shows clearly that the excitation band was changed in shape (fig. 5) for a sample which was partially quenched.Hydrogen reacts at different rates with species absorbing in different parts of the excitation band causing the observed change of shape. No corresponding change of shape was found in the luminescence band as it decreased in intensity. Quenching by oxygen caused no change of shape in either the excitation or the emission bands, only a decrease in intensity. Since we might expect oxygen and hydrogen to adsorb on different sites it seems likely that the absorption and emission processes occur at different, but probably closely related, sites on the surface and that energy can be transferred along the surface. NATURE OF EXCITATION A N D EMISSION SITES It is convenient to use a chemical approach to try to define the sites on the surface. In general we associate oxygen adsorption with a part of the surface which is oxygen deficient or has a local excess of cations.Although adsorption studies do not give a picture of the surface on the atomic scale, studies using other techniques, such as electron spin resonance,' show that adsorbed forms of oxygen such as 0; are associated with a nearby cation on the surface and a hyperfine interaction is frequently observed if the cation has a non-zero nuclear spin. The position is less clear for hydrogen, in some situations it is known to react with the cationY8* e.g., ZnO to form a hydride but in general we may expect hydrogen to react with that part of the surface which is cation deficient or has a local excess of oxygen ions. For SrO most of the loss of intensity of the emission is not easily reversible and the hydrogen must2920 SURFACE STATES I N OXIDES be strongly held on the surface possibly as a hydroxyl.Based on the arguments above we suggest that the sites on the surface responsible for emission are those that are oxygen deficient or have a local excess of cations, whereas the sites responsible for the absorption of light are cation deficient or have a local excess of oxygen ions. Such sites will represent situations where the coordination of the ion is markedly lower than found in the bulk crystal. These ideas enable us to examine a range of possibilities in more detail. From a comparison of the excitation spectra of the oxides it is clear that the excitation energy decreases with increasing cation radius from Mg2+ to Ba2+ ; these spectra have been shown to arise from energy absorption at the surface (where the surface is defined as those atomic layers able to react rapidly with an adsorbate). An intrinsic bulk absorption in the U.V.can arise at energies just less than the band gap from a charge transfer process to give excitons (electron-hole pairs) which can be considered as M+O-. In the surface region, it seems likely that less energy would be required for the charge transfer process because of the reduced Madelung energy at the surface. It is possible to make a qualitative calculation of the intrinsic surface state energies using the approach of Levine and Mark lo where the surface ions are considered to be equivalent to bulk ions except for their reduced Madelung constant. The band gap ratio 8 of surface (&) to bulk (&) given by where y = c,/cb is the ratio of surface to bulk Madelung constant, and where I is the ionization potential (in eV) of the cations, A the electron affinity (in eV) of the anion, c b the bulk Madelung constant, r the lattice parameter and 2 the valence of the ions.Values for y and p are available lo and since the alkaline earth oxides will not be 100 % ionic the calculation has been carried out for both 2 = 1 and 2. Calculated values for the surface band gaps on the surface (100) plane and on higher index planes are shown in table 3. = (Y -PM1 - P h p = 0.0347 r(1-A) CbZ TABLE 3.-ALKALINE EARTH OXIDE SURFACE BAND GAPS (ev) CALCULATED FROM THE MADELUNG POTENTIAL lo CaO SrO BaO surface Eb = 8.7 11, 12 Eb = 7.712, 13 Eb = 6.7 12 E b = 4.4 1 3 MgO planes coordination Z = 1 Z = 2 z-1 z=2 z = l 2-2 z = l z=2 5 8.25 8.00 7.32 7.08 6.37 6.15 4.19 4.03 4 7.11 6.26 6.37 5.54 5.57 4.79 3.66 3.11 4 6.11 4.70 5.51 4.15 4.84 3.56 3.19 2.29 3 4.18 1.74 3.90 1.54 3.47 1.23 2.30 0.73 (100) (110) (210) (211) Eb denotes the bulk band gap This theoretical approach can be considered only as an approximation ; however, it does reproduce the main features of the results and it predicts a decrease in excitation energy through the series MgO to BaO for the various index planes.The energies for the (100) surfaces are only slightly shifted from the bulk while those for the higher index planes are much closer to the observed data. However, the energies of the surface excitons would be expected to be slightly lower than the calculated band gaps.On an atomic scale this means that ions whose coordination number is t 5 must be involved in the photoluminescence process, i.e., edge and corner sites on the surface. This appears to be an appropriate model since there is no evidence to suggest anyS. COLUCCIA, A. M. DEANE AND A. J . TENCH 292 1 non-stoichiometry of the oxide and samples prepared in vacuo will adsorb both hydrogen and oxygen. In the charge transfer model it seems most plausible that an electron is transferred from an anion in a position of low coordination to a cation. Independent evidence for the existence of anions on the surface which can easily lose an electron comes from the studies on the interaction of electron acceptors with the surface where it has been shown that nitrobenzene or related molecules 14* l5 can abstract electrons from -0.1 % of the surface ions.It appears probable that some oxygen anions in sites of very low coordination such as corner or edge positions are responsible for both phenomena in oxides that have been evacuated at high temperatures. This is in agreement with the ideas proposed by Cordischi et all6 suggesting that on MgO outgassed at 1200 K the predominant surface centres are coordinatively unsaturated 02- ions. Other sites appear to make only a minor contribution to the luminescence spectrum. The excitation spectra are characterised by two main bands; this is taken as strong evidence that there are two sites on the surface which absorb at different energies and which correspond to sites of different coordination number.The lower energies will correspond to the lower coordination numbers (table 3). In addition to the main doublet structure there is evidence that both bands are complex. For example the more detailed analysis of the excitation spectrum for SrO shows several components. These components are most likely to result from minor differences in the environment of the ions such as changes in the next nearest neighbour ions. It has already been argued that luminescence occurs at a different site, which is reactive towards oxygen, and hence at an exposed cation. This luminescence arises from a process such as, M+O- + M2+02-+hv at an exposed metal ion site, and the luminescence energy decreases in the expected way through the oxide series (table 1).We can understand the processes occurring at the surface in terms of the formation of an exciton by adsorption of light at a cation-anion pair involving an oxygen ion of low coordination. This exciton, after moving over the surface in a random fashion, can decay at a cation-anion pair which involves a cation site of low coordination, with the emission of light. The excitation and emission sites may be close together but the lifetime is sufficient for the exciton to have moved some distance over the surface. This picture is consistent with the way in which hydrogen and oxygen affect the two sites independently. We only observe the radiative pathways but many of the excitons will decay via non-radiative processes. The measured lifetimes of the luminescent states to s, table 2) are relatively long and probably indicate triplet state formation but the data so far recorded do not justify detailed analysis.The curves have been interpreted as several different first order lifetimes arising from different luminescent sites on each oxide, consistent with the other experimental evidence. However it is possible that the lifetime is much longer than s and that the decay curve reflects the non- radiative decay through a higher order process. The large temperature dependence shows that non-radiative processes are dominant at room temperature. The sites at which absorption and luminescence occur cannot be defined un- ambiguously, although the earlier arguments lead to the conclusion that they must be associated with anions and cations respectively, in positions of low coordination on the surface.For the absorption process such sites could be associated with oxygen ions on edge or corner sites of the crystal, cation vacancies or possibly where cations are missing from an edge or corner site and oxygen ions are exposed. E.s.r. data2 922 SURFACE STATES IN OXIDES indicate that holes trapped at cation vacancies can be produced by U.V. irradiation and these holes are destroyed by hydrogen. However, the trapped holes absorb at much longer wavelengths.l* The sites linked to the emission could be surface anion vacancies, similar point defects or cations at step and corner sites where the coordina- tion is less than 5. The bulk anion vacancy with an electron (Ff centre) in MgO absorbs at 250 nm (5.0 eV) and emits at 396 nm (3.13 eV) and for CaO it is 340 nm (3.65 eV) and 370 nm (3.35 eV) respectively.lg* 2o Even after allowing for modified values at the surface, these do not agree well with the observed values for the powders and there is no indication of any significant formation of trapped electron centres from the e.s.r.data. There is, however, an interesting correspondence between the broad luminescence observed at slip planes in the bulk and that from the surface. Luminescence from deformed crystals 21 9 22 of magnesium, calcium and strontium oxides gave excitation peaks of about 0.2 eV above those reported here (table 1) and emission about 0.4 eV lower than the values for the surface luminescence. Impurity ions did not contribute and the luminescence was thought to be associated with vacancy clusters; such a cluster can be regarded as a small hole in the bulk crystal where surface luminescence occurs.We would like to thank Miss K. Mahmood and Miss J. Abbott, who carried out some of the measurements described ; S. Coluccia and A. J. Tench acknowledge financial support from NATO. A. J. Tench and G. T. Pott, Chern. Phys. Letters, 1974,26,590. S. Coluccia, A. M. Deane and A. J. Tench, Proc. 6th Int. Congress on Catalysis, London, 1976, 1, 171. A. Zecchina, M. G. Lofthouse and F. S. Stone, J.C.S. Furaday I, 1975,71, 1476. A. Zecchina and F. S. Stone, J.C.S. Faraduy I, 1976,72,2364. P. B. Merkel and W. H. Hamill, J. Chem. Phys., 1971,55,2174. S. Coluccia, A. M. Deane and A. J. Tench, to be published. J. H. Lunsford, Catalysis Rev., 1973, 8, 135. A. L. Dent and R. J. Kokes, J. Phys. Chern., 1969,73,3772. lo J. D. Levine and P. Mark, Phys. Rev., 1966,144,751. l 1 G. H. Reiling and E. B. Hensley, Phys. Rev., 1958,112, 1106. l 2 V. I. Neeley, Ph.D. Tliesis (University of Oregon, Eugene, 1963). l3 H. H. Glascock and E. B. Hensley, Phys. Rev., 1963, 131, 649. l4 A. J. Tench and R. L. Nelson, Trans. Faraday Soc., 1967, 63,2254. l 5 M. Che, C. Naccache and B. Imelik, J. Catalysis, 1972, 24, 328. l6 D. Cordischi, V. Indovina and M. Occhiuzzi, J.C.S. Furaday I, 1978, 74,456. l 7 A. J. Tench and J. F. J. Kibblewhite, J.C.S. Chem. Comm., 1973, 955. l9 G. P. Pells and A. E. Hughes, AERE R8686,1977. 2o A. E. Hughes and B. Henderson, Point Defects in Solids, ed. J. H. Crawford and L. M. Slifkin 21 T. J. Turner, N. N. Isenhower and P. K. Tse, Solid State Comm., 1969, 7, 1661. 22 Y. Chen, M. M. Abraham, T. J. Turner and C. M. Nelson, Phil. Mag., 1975, 32,99. * R. P. Eischens, W. A. Pliskin and M. J. D. Low, J. Catalysis, 1962, 1, 180. Y. Chen and W. A. Sibley, Phys. Rev., 1967, 154,842. (Plenum Press, N.Y., 1972), vol. 1, chap. 7. (PAPER 8/896)

ISSN:0300-9599    

DOI:10.1039/F19787402913

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Enthalpy of Interaction Between Some Cationic Polypeptides and n-Alkyl Sulphates in Aqueous Solution B Y MARIA I. PAZ-ANDUDE,? MALCOLM N. JONES AND HENRY A. SKINNER* Department of Biochemistry and Chemistry, University of Manchester, Manchester M13 9PL Received 1 5 th May, 1978 The enthalpies of interaction of a homologous series of n-alkyl sulphates with poly(L4ysine)- hydrobromide, poly(L-arginine)hydrochloride and poly(L-histidine)hydrochloride have been measured at 25°C. A linear relationship between the enthalpy of interaction and carbon chain length has been found for alkyl chain lengths above C8. The data support a model based on a stoichiometric interaction between the anionic head group of the n-alkyl sulphates and the cationic side chains of the polypeptides. The results are discussed in relation to the interaction between surfactants and proteins, and lead to the view that a major contribution to the enthalpy of interaction arises from the surfactant-cationic residue interactions, but that there remains an additional contribution from surfactant-apolar aminoacid residue interactions.The interactions between surfactants and globular proteins have been extensively studied inaiiily because surfactants at low concentrations unfold proteins and also because of the general use of surfactants as solubilizing agents.l’ The enthaipy of interaction between a globular protein and a surfactant is generally a complex function of surfactant concentrati~n.~-~ There are at least two contributions to the net enthalpy of interaction arising from (1) the binding of surfactant ions, (generally an exothermic process) and (2) the unfolding of the native conformation (an endothermic process).It is believed that the initial stage in the interaction involves the binding of the surfactant ion to oppositely charged sites on the surface of the native protein molecule, although the ionic interaction is modulated by hydrophobic effects3 The complexity of protein-surfactant interactions has prompted us to investigate simpler systems consisting of cationic polypeptide plus anionic surfactants. Micro- calorimetric measurements are reported here of the enthalpies of interaction of three synthetic polypeptides, [poly(L-lysine)hydrobromide, poly(L-arginine)hydrochloride and poly(~-histidine)] with a homologous series of n-alkyl sulphates in aqueous solution.Previous work on such systems has been limited mainly to spectroscopic and binding studies on poly(L-lysine) plus surfactants. 6* EXPERIMENTAL MATERIALS Poly(L-1ysine)hydrobroniide (lot LY 190A, mol. wt . 40 300), poly(L-arginine)hydro- chloride (lot AR 55, mol. wt. 13 900) and poly(L-histidine) (lot HS 36, mol. wt. 11 100) were purchased from Miles-Yeda, Israel. The materials were used as supplied. The molecular weights were determined by the manufacturers by ultracentrifugation. or supplied by t Present address : Departmento de Fisica, Universidad de Santiago, Santiago de Compostela, Sodium n-dodecyl sulphate was prepared as previously described Spain. 29232924 INTERACTION BETWEEN POLYPEPTIDES AND N-ALKYL SULPHATES Cambrian Chemicals, Croyden.No difference was found between results obtained with the two samples. All the other n-alkyl sulphates were supplied by Cambrian Chemicals. All solutions were made up in doubly distilled water. MICROCALORIMETRY The Beckman 19OB twin-cell conduction calorimeter has been described in detail else- where ’ and was calibrated electrically.1° Matched drop well annular glass cells were used. The procedure adopted was to charge one cell with 0.30(+0.01)x kg of polypeptide solution in the drop well and lO.O(+O.l)x kg of surfactant solution of known con- centration in the annular space. The reference twin cell was charged similarly with the same amount of polypeptide solution and lO.O(+O.l)x kg of water. On mixing, the dilution of polypeptide is identical in both cells so that the dilution enthalpies cancel.The enthalpies of surfactant dilution are negligible under these conditions. The majority of the measurements were made with surfactant concentrations below their respective critical micelle concentrations in water at 25°C. The initial polypeptide concentrations were : poly(L-1ysine)HBr (1.5 % w/v), poly(L-arginine)HCl (1.4 % w/v) and poly(L-histidine) (0.9 % w/v). On dilution the final amino acid residue concentrations were approximately 2.1,2.1 and 2.0 mmol dm-3 calculated on the basis of molecular weights of 209.09, 192.65 and 137.14, respectively. Because poly(L-1ysine)HBr is very hygroscopic the solutions were made up in a glove box purged with dry nitrogen. Poly(L-histidine) required acidification with hydrochloric acid to facilitate solution.The enthalpy data presented below are expressed in kJ mol-1 of poly- peptide residues [kJ (mol res.)-l] and were calculated from the equation. AH = 99’7.1($) where Q is the measured enthalpy in kJ, w is the weight of polypeptide solution of residue concentration c mol dm-3. The equation assumes a negligible difference in density between polypeptide solution and water (density 0.9971 g ~ m - ~ ) at 25°C. RESULTS Fig. 1 shows the enthalpy of interaction of poly(L-lysine) as a function of surfactant concentration for sodium n-alkyl sulphates covering carbon chain lengths from C to Cq. In every case the enthalpies come to limiting values at a surfactant concentration indicative of saturation of the polypeptide chain with surfactant ions.The data for n-hexyl sulphate are slightly anomalous at the higher surfactant concentrations. The most exothermic interaction is found with the longest carbon chain (C,,), and it is noteworthy that in this case the two linear portions of the curve extrapolate to intersect at a concentration of approximately 2.4mn1oldrn-~, which is close to the lysine residue concentration of 2.1 mmol dm-3. This conforms with the view of Satake and Yang that there is a stoichiometric (1 : 1) interaction between the cationic amino acid side chain and the anionic n-dodecyl sulphate ion. For the other surfactants saturation occurs in the region of 2-4 mmol dm-3. Although interaction of the surfactant with the poly(L-lysine) hydrobromide resulted in a turbid solution or a flocculant precipitate for the Cl0, CI1 and CI2 alkyl sulphates at concentrations > 1 mmol dm-3, all the other systems remained visually clear.In the former cases there may be a thermal contribution from precipitate forma- tion, but the consistency between the results for the same system above and below the precipitation threshold, and between these systems which gave precipitates and those that did not, makes it unlikely that precipitation contributes significantly to the interaction enthalpies.M. I. PAZ-ANDRADE, M . N. JONES AND H. A. SKINNER 2925 The enthalpy data for poly(L-arginine)hydrochloride and poly(L-histidine)hydro- chloride are plotted in fig. 2 and 3 respwtively. These plots show similar patterns to that given by poly(L-lysine)hydrochloride, and likewise show that for n-dodecyl sulphate, saturation occurs at surfactant concentrations close to the aminoacid residue concentrations of 2.1 and 2.0 mmol dm-3 for poly(L-arginine) and poly(L- histidine), respectively.mmol dm-3 -10 -20 5 " - O--C 12 - 4 3 2 1 0 - 1 -2 -3 - 4 -5 -6 -7 -8 I 1 FIG. 1 .-Enthalpy change [AN/kJ (mol amino acid residue)-l] on interaction of sodium n-alkyl sulphates (chain length Ca with poly(L-1ysine)hydrobromide in water at 25°C. mmol dm-3 4 2 d o g - 4 - - 6 & - 2 3 -10 w 8 -8 -12 4 -11 - 16 t , o - o , , , , , 0-c 4, 6 7 8 9 1 0 1 1 8-CB O ClO r FIG. 2.-Enthalpy change [AH/kJ (mol amino acid residue)-l] on interaction of sodium n-alkyl sulphates (chain length C,) with poly(L-arginine)hydrochloride in water at 25°C.For the three polypeptides the limiting enthalpy values for each surfactant are plotted as a function of the number of carbon atoms in fig. 4. These plots are linear for surfactants with carbon chain lengths in the range C8 to CI2 but markedly deviate from linearity below C8. The equations for the linear regions based on a regression analysis of all the values of the limiting enthalpies are as follows (+ represents the standard deviation of slope and intercept) poly(L-1ysine)HBr AH/kJ (mo! res.)-l = -2.374(+0.036)n,+22.61(+0.12)2926 INTERACTION BETWEEN POLYPEPTIDES AND N-ALKYL SULPHATES pol y (L-arginine)HCl pol y( L-histidi ne) HCl AH/kJ (mol res.)-l = - 1.445(+0.120)n,+0.180($_0.380) AH,,kJ (mol res.)-l = - 1.863( +0.063)n,+9.190(+0.201). (3) (4) In water, the side chains of all three polypeptides should be fully ionised.The pH of the solutions were - 5, except for the poly(L-histidine) solutions which were acidified to pH N 3. Some measurements were carried out on the system poly(L-1ysine)f sodium n-dodecyl sulphate in glycine buffers of ionic strength 0.0167. At both pH 2.95 (glycine-HC1) and pH 8.50 (glycine-NaOH) the limiting enthalpies did not differ significantly from those measured with pure water. mmol dm-3 FIG. 3.-Enthalpy change [AH/kJ (mol amino acid residue)-'] on interaction of sodium n-alkyl sulphates (chain length C,) with poly(L-histidine)hydrochloride in water at 25°C. nC 6 4 2 M O -2 2 - 4 Q - 6 g -e 2 -10 $ -12 - --- Q -11, -1 6 -1 8 -29 -22 FIG. 4.-Enthalpy change [AH/kJ (mol amino acid residue)-'] on interaction of sodium n-alkyl sulphates as a function of carbon chain lengths in water at 25°C 0, poly(L-1ysine)hydrobromide : A, poly(L-histidine)hydrochloride ; 0, poly(L-arginine)hydrochloride.M. I .PAZ-ANDRADE, M . N. JONES AND H. A. SKINNER 2927 DISCUSSION The linear relationships between the enthalpy of interaction and the carbon chain lengths of the surfactants from c8 to C12 as represented by eqn (2)-(4) correspond to methylene increments of -2.37, - 1.45 and - 1.86 kJ (mol res.)-' for poly(L-lysine), poly(L-arginine) and poly(L-histidine), respectively. In comparison with methylene increments calculated for other processes these values are very large, e.g., the enthalpies of micellization of the Clo and C12 alkyl sulphates at 25°C give a methylene increment of -0.42 kJ mol-1 ; the enthalpies of transfer of normal aliphatic hydro- carbons (C,C,) from aqueous solution to the liquid state (data quoted by Nemethy and Scheraga)12 give -0.82 kJ mol-', whereas the extensive assessment of hydro- phobic interactions in n-alkanes (C2-C10) reported by Gill and Wadso l3 gives -0.318 kJ mol-1 for the latter process.A higher value of the methylene increment is obtained for the transfer of aliphatic alcohols from water to the liquid state, e.g., for n-butanol and n-pentanol l4 we calculate - 1.59 kJ mol-'. The present study is of the interaction of surfactant with a polyelectrolyte of random conformation, in which it seems probable that the surfactant head group interacts ionically with the cationic charge on an amino acid side chain.The main question is that of the resultant orientation of the surfactant hydrocarbon chain in the polyelectrolyte + surfactant complex. Two extremes can be envisaged ; (a) the surfactant chains pack intimately into the polypeptide to maximise contact with the side chain methylene groups and possibly with part of the polypeptide backbone, or (b), the surfactant chains orient themselves away from the polypeptide chain, into the aqueous phase, possibly then to interact with similarly bound surfactant on adjacent amino acid side chains, leading to incipient formation of micelle~.~~ To form roughly spherical micelles in this way would require the polypeptide chain to adopt a coiled conformation. Of these two extremes, the first seems to be more likely in that, (1) the methylene increments are much larger than for micellization in aqueous solution at 25"C, (2) the methylene increments are specific for a given polypeptide and (3) circular dichroism spectroscopy has indicated that sodium n-octyl sulphate in neutral solution induces a partially helical conformation with poly(L-lysine), whereas the higher homologues (Clo-C1 6 ) induce a P-pleated sheet structure, the percentage of p-structure increasing with surfactant chain lengtha6 Such rigidity precludes globular micellar structures of the type postulated for the complex formed between polyethylene oxide and sodium-n-dodecyl sulphate, in which the polymer is wrapped around a surfactant micelle.'6 On the other hand, favouring the alternative possibility, is the fact that n-alkyl sulphates with carbon chain lengths below c8 do not form micellar s ~ h t i o n s , ' ~ and our results (fig.4) show that the breakdown of linearity occurs below c8. We consider now to what extent the present results relate to the measured enthalpies of interaction of surfactants with globular p r ~ t e i n s . ~ ' ~ Native proteins have relatively rigid structures, so that for comparison with a simple polypeptide in solution the unfolded protein is a more strict starting point. The final conformation of a protein + surfactant complex will resemble that of a polypeptide + surfactant complex, apart from conformational restriction imposed by disulphide bonds, and a less uniform distribution of binding sites. At saturation, globular proteins bind 2 1.4 g of sodium n-dodecyl sulphate per gram of protein,19-21 which, for an average molecular weight of -130 per amino acid residue, corresponds to the binding of 0.6 surfactant molecules per residue, i.e., rougly half that in a polypeptide + surfactant complex.Table 1 records thermochemical data for several protein-surfactant systems.2928 INTERACTION BETWEEN POLYPEPTIDES AND N-ALKYL SULPHATES The numbers of ionised lysine, arginine and histidine residues in the proteins 2 2 are shown together with the experimentally measured enthalpies of interaction (AHexp) at a surfactant concentration of 4 mmol dm-3. This concentration is below that required for saturation, but high enough to have induced denaturation, (except in case of the lysozyme + Clo surfactant system).If we now assume ; (i) that surfactant molecules interact preferentially with cationic sites, the enthalpy of interaction, AHc, TABLE 1 .-CONTRIBUTIONS TO ENTHALPIES OF INTERACTION OF GLOBULAR PROTEINS WITH SURFACTANTS IN AQUEOUS SOLUTION AT 25°C no. aminoacid system residues (ionised) AHb Lys Arg His Y O /kJmol-l /kJmol-1 /kJmol-l /kJmol-I SDS ref. AH& AHc AHa AHb /kJmol-1 lysozyme + CI2Hz50SO;Na+, pH 3.6,IO.W 6 11 1 25 lysozyme + el 2H2 50SO;Na+, pH 9, Z 0.009 5.9 11 0 25 lysozyme + CloH210SO;Na+, pH 3.6,10.004 6 11 1 17 Iysozyme + CloHz 10SO;Na+, pH 9,10.009 5.9 11 0 10 ribonuclease A + C1 2H250SO;Na, pH 7,I0.005 10 4 0.4 68 trypsin+ C1 2H2 5osOiNa+, pH 3.5, Z 0.01 14 2 3 38 trypsin + C1 2H2 ,OSO;Na+, pH 5.5,IO.Ol 14 2 2.3 33 /3-hctoglobulin + C1 2H250SO;Na+, Cl2HZ50SO;Na+, pH 5.5,IO.Ol 2 3 1.6 46 ovalbumin + C1 2H250SO;Na+, pH 7,10.005 20 15 0.7 97 bovine serum albumin+ C12H250SO;Na+, pH 7.9, 10.005 58 23 1.7 198 pH 3.5,10.01 2 3 2 48 ~-1~toglobulin + - 168.8 - 118.7 - 88.7 - 30.0 - 17.1 - 193.4 -116.5 - 600 - 150 - 220 - 660 -237 22223 -223 22223 -159 22223 -64 2za3 -133 25023 -156 -280b -147 ~ 2 8 0 b -89.6 190 24 -84.3 19024 -384 ~ 5 3 0 b -758 ~ 7 9 0 b - 154 - 118 - - - 134 - 317 - 250 - 700 - 256 - 366 - 692 - 22 - 15 I - - 2.5 - 17 - 17 - 17 - 6.5 - 6.0 - 6.0 18 18 18 18 3 5 5 4 4 3 3 a At total surfactant concentration of 4 mmol ~Irn-’~; estimated on the basis of 12 J g-l.being calculated from the number of such interactions, and eqn (2)-(4) above, (ii) that the proteins are completely unfolded, involving a denaturation enthalpy term AHd, and (iii) that following saturation of the cationic sites, the remaining bound surfactant ions interact with apolar aminoacid residues in the protein with an enthalpy AH,,, then AHexp = AHC +AH, 4- AH,.( 5 ) The enthalpies of unfolding have been taken from the literature 2 3 9 24 or were estimated from the work of Privalov and Knechinashivili 23 on the basis of 12 J 8-lM. I . PAZ-ANDRADE, M. N . JONES AND H . A. SKINNER 2929 protein. Values of AHb were then calculated from eqn (5). Assuming that surfactant ions not bound to cationic sites bind to (V-number of cationic residues) apolar sites, AH, can be expressed in terms of the enthalpy per mole of bound surfactant, values for which are given in column 10.The lysozyme + Clo surfactant system is a special case, in that the bound surfactant is insufficient to cover the total number of cationic sites ; hence, in calculating AHc, it was assumed that some of the arginine sites remain unoccupied, as the lysine residues are known to be exposed on the protein surface. Furthermore, the net enthalpy of interaction is less than that arising from interaction with the cationic sites, so that in this particular case the protein is not fully unfolded. For the other systems it is noteworthy that AHb is substantially larger than the enthalpy of a micellization process,25 (e.g., for sodium n-dodecyl sulphate in a medium of ionic strength 0.023, AH(micellization, = -0.64 kJ mol-1 at 25°C). This implies a stronger interaction between surfactant molecules and protein chains than between surfactant molecules in micelles, consistent with the fact that complex formation with proteins occurs below the critical micelle concentration of the surfactant concerned.Despite the approximate nature of the present analysis, it seems clear that the surfactant non-cationic residue interactions can make a significant contribution to the overall enthalpies of interaction of sodium dodecylsulphate with proteins, but there is no obvious correlation with the number and nature of apolar aminoacid residues in the protein. We thank Mr. A. Butler and Miss M. Griffin for experimental assistance and One of us (M. I. P.-A.) Dr. G. Pilcher for calibration of the microcalorimeter. expresses her thanks to the Spanish authorities for sabbatical leave.M . N. Jones, Biological Interfaces (Elsevier, Amsterdam, 1973, chap. 5. C. Tanford and J. A. Reynolds, Biochim. Biophys. Actu, 1976,457,133. E. Tipping, M. N. Jones and H. A. Skinner, J.C.S. Faruduy I, 1974, 70, 1306. M. N. Jones and A. E. Wilkinson, Biochem. J., 1976,153,713. M. N. Jones, Biochim. Biophys. Actu, 1977,491,121. I. Satake and J. T. Yang, Biochem. Biophys. Res. Comm., 1973,54,930. ' I. Satake and J. T. Yang, Biopolymers, 1975, 14, 1841 ; 1975, 15,2263. ' M. N. Jones, H. A. Skinner, E. Tipping and A. Wilkinson, Biochem. J., 1973, 135,231. H. A. Skinner in Biochemical Microculorimetry, ed. H. D. Brown (Academic Press, New York, 1969). lo M. N. Jones, G. Agg and G. Pikher, b, Chem. ThermodyMmics, 1971, 3,801, l1 E. D. Goddard and G. C. Benson, Csnud. J. Chem., 1957,35,986. l2 G. Nemethy and H. A. Scheraga, J. Chem. Phys., 1%2,36, 3401. l3 S. J. Gill and I. Wadso, Proc. Nat. Acad. Sci., 1976, 73, 2955. l4 E. M. Arnett, W. B. Kouer and J, V. Carter, J. Amer. Chem. SOC., 1963,91,4028. l6 B. Cabane, J. Phys. Chem., 1977,81,1639. l8 M. N. Jones and P. Manly, unpublished results. l9 J. A. Reynolds and C. Tanford, Proc. Nut. Acud. Sci., 1970,66, 1002. 'O R. Pitt-Rivers and F. S. A. Impiombato, Biochem. J., 1968,109, 825. 21 T. Takagi, K. Tsiyii and K. Shirahama, J. Biochern., 1975,77,939. 22 Analytical Methods ofprotein Chemistry, ed. P. Alexander and H. P. Lundgren (1966), vol. 4, 23 P. L. Privalov and N. N. Knechinashvili, J. Mol. Biol., 1974, 86, 665. 24 S. Lapanje, M. Launder and J. Skerjanc, Proc. 4th Int. Con5 Chem. Thermodynamics (Nouvelle 2 5 M. N. Jones, G. Pilcher and L. Espada, J. Chem. Thermodynamics, 1970,2,333. I. Satake and J. T. Yang, Biopolymers, 1976, 15,2263. P. Mukerjee and K. J. Mysels, National Standard Reference Data System, N.B.S. 36, 1971. chap. 5. Berthier, Arles, 1975), sect. 5, no. 16, 92. (PAPER 8/897)

ISSN:0300-9599    

DOI:10.1039/F19787402923

出版商:RSC    

年代:1978

数据来源: RSC