摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3245-3252 New Approach to Sensitivity Analysis of Multiple Equilibria in Solutions llie Fishtik Institute of Chemistry, Academy of Sciences of Moldova ,277028 Kishinev, str. Academiei 3, Republica Moldova lstvan Nagypal and Ivan Gutman? Institute of Physical Chemistry, Attila Jozsef University, P.O. Box 105,H-6701 Szeged, Hungary We consider multiple equilibria in solutions in which the interaction of n chemical species is described by means of m stoichiometrically independent reactions (SIRs). For the study of certain thermodynamic properties of such systems, in particular, for sensitivity analysis, it is important to know the determinant A of the Hessian matrix of the Gibbs energy, as a function of the extent of the SIRs.Any linear combination of SIRs, in which (at least)rn -1 species are not involved, is called a Hessian response reaction (HR): Several properties of the HRs are pointed out, in particular, the equivalence of A to the sum of contributions originating from each HR. The effect of temperature and pressure on chemical equilibria in ideal solutions is analysed. It is shown that the sensitivity coefficient of a chemical species A, may be presented as a sum of contributions coming from all HRs in which A, is involved. Each of these contributions is a product of the stoichiometric coefficient of Ai, the enthalpy or volume change of the respective HR, and a concentration-dependent term which is always positive. It is also shown that the relaxation contribution to the heat capacity is a sum of contributions over all HRs.The position of chemical equilibrium is governed by a series of parameters, such as equilibrium constants, temperature, pressure and initial concentrations. The effect of the change of these parameters on the position of chemical equilibria is the subject of the sensitivity analysis.' Sensitivity analysis requires the calculation of a set of derivatives of the equi- librium concentrations with respect to the above parameters. These derivatives are called sensitivity coefficients. The calculation of the sensitivity coefficients for gas-phase reactions was discussed in general terms and the results are summarized in ref. 1. Complex equilibria in solutions and other special cases were considered elsewhere.2-' By means of sensitivity coefficients, one can predict the response of the system to the change of the respective param- eters.In the case of a single equilibrium process, it is easy to predict the sign of the sensitivity coefficients by means of the Le Chatelier principle. For complex systems, however, the required calculations are complicated and there is no qualit- ative way to predict the sign of the overall effect. The main idea of the present paper is to show how the overall sensitivity may be given as a linear combination of contributions originating from individual reactions, which we call response reactions, and which strictly obey the Le Cha- telier principle. In other words, the Le Chatelier principle is extended to multiple equilibrium systems.The effect of tem- perature and pressure is analysed in this paper, the effect of the other parameters will be discussed elsewhere. We arrive at the concept of response reactions by examin- ing the determinant of the Hessian matrix. Namely, when examining different aspects of the equilibrium conditions in terms of stoichiometric formulation we encounter the second derivatives of the Gibbs energy, with respect to the extents of the reactions occurring in the system ~0nsidered.l~ These derivatives may be arranged into a square matrix of order m, where m is the number of independent reactions, which is called the Hessian matrix of the Gibbs energy. Of special importance is the determinant of this matrix, the Hessian determinant, as well as its minors.In particular, the Hessian 7 Permanent address : Faculty of Science, University of Kraguje-vac, Kragujevac, Yugoslavia. determinant and its diagonal minors are known to appear in the expressions for sensitivity coefficients. l4 The Hessian determinant also plays an important role in different algo- rithms for calculating the equilibrium compositions or equi- librium constants of systems with simultaneously occurring chemical reactions, as well as in the sensitivity analysis of complex chemical equilibria. '*14 In spite of the importance of the Hessian determinant, it seems that the only thing we know about it, is that the Hessian determinant and its diagonal minors are of positive va1~es.l~It turns out, however, that the Hessian determinant has a number of other interesting properties which, as far as we are aware, have not been observed before.We study the Hessian determinant and introduce the concept of Hessian (response) reactions. We then apply the Hessian reactions to the analysis of certain sensitivity coefficients. Definitions and Notation In this paper we are concerned with the equilibrium state of a system in which chemical reactions take place. We examine the general case in which the interaction of n distinct chemi- cal species A,, A,, . . .,A,, is described by means of m chemi-cal equations: V11A1 + VlzA, + *. . + V1,An = 0 vZ~A,+ v~,A,+ * * + ~2,,A, = 0 ...............................I vmlA1+ vm2 A, + --+ v, A,, = 0 J Throughout this paper it is assumed that the above equations are stoichiometrically independent, i.e. that It is customary to call the relations in eqn. (1) chemical reac- tions; they will be referred to as stoichiometrically indepen- dent chemical reactions (SIRs). Recall that a set of SIRs can be chosen in many different, but (from the point of view of chemical equilibrium analysis) equivalent ways, i:e. the value of the Hessian determinant does not depend on the choice of the SIRS. The extent of the jth SIR, namely of cf=lvjiAi = 0, at equilibrium will be denoted by tj,j = 1, 2,. ..,rn. Let G be the Gibbs energy of the system considered. G can be viewed as a function of the parameters tl, t,, ..., t,.Then the elements of the respective Hessian matrix are d2G/a<,at,. Instead of a2G/at,at, it is cu~tomaryl*'~ to use the quantities G,s, where G,~ RT = -1 a2G/at,at, (3) In this study we examine the determinant (4) which is the Hessian determinant of the system considered. Hessian Reactions It is necessary to define the determinant D (of order m) I V1,il V1,iz .** v1, i,-1 VI, i, I D = D(il, i,, ..., i,,,) = '2, il ... '2, iz . . . ... ... '2, im-1 ... '2, i, I vm,il ... (5) and its minors Djk, where Djk is obtained from D by deleting its jth row and kth column. Now, any linear combination of eqn. (1) leads to a new chemical reaction between the species A,, A,, ...,An,which is acceptable from a stoichiometric point of view. The general form of such a linear combination is m Tn 1 Consider now the special linear combination of the above kind obtained by choosing: Aj = (-ly'+"Djm; j = 1, 2, .. . ,rn Then eqn. (6)becomes m rn 1 n rm 1 and bearing in mind that j=1 is just the expansion of the determinant D with respect to its mth column, we arrive at the reaction n C vimAim= 0; vim= mil, i, ,..,i,,,) (7)i,= 1 The reaction, eqn. (7), has the noteworthy property that vim = 0 whenever i, = i, or i, = i, or ...i, = i,-1, because then two columns in the determinant D(il, i,, ..., i,,,) are equal [see eqn. (5)]. This means that the species A,,, Aiz,..., Aim-l have zero stoichiometric coefficients and are thus not involved in eqn.(7). For reasons that will become evident later, we call eqn. (7) a Hessian response reaction, or shorter J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 a Hessian reaction (HR). We denote eqn. (7) by H = X(i,, i,, ...,im-l), where i,, i,, ...,im-indicate the species which are absent from S. There are (m-1) possible choices of the species A,,, Ai2, ...,Aim-,whose absence defines a particular Hessian reaction H,but not all such reactions need to be distinct and only rn of them can be linearly independent. We can now easily deduce an important property of the HRs. Let X be a state function whose change in thejth SIR is AXj, j = 1,2, ...,rn. Then the change of X in eqn. (6) is m AX = CljAXj j=1 Applying the above formula to a Hessian reaction, eqn.(7), we obtain: m AX(H) = 1(-1)j+'"DjmAXj j=1 Property of the Hessian Determinant In the subsequent section we demonstrate that the Hessian determinant A, eqn. (4), is equal to the sum of certain increments, each being associated with a particular Hessian response reaction : where the summation is over all HRs, and T(X),the contri- bution of the reaction H = S(il, i,, . . .,im-1), has the form UJf) (10) In the above formulae and later on, [A,] denotes the (equilibrium) concentration of the species Ai. D(il, i, ,. . .,i,) is defined by eqn. (5). Note that owing to eqn. (l), at least one of the determi- nants D(il, i,, ..., i,,,) must be non-zero. Consequently, at least one HR must have a non-zero contribution and, there- fore, A is necessarily of a positive value.Proof of Eqn.(9) Our starting point is the thermodynamic relation14 for the quantity G,, eqn. (3), where n, is the amount of the species A, in the system con- sidered, and n, the total amount of all species present in the system [including also those which do not participate in eqn. (l)]. Introducing the auxiliary quantity psi n J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 eqn. (1 1) can be rewritten as i Vri Psi Grs = i=l ni (13) We shall now exploit the following property of determi- nants: If alj = aaij + /?a';j+ yayj + * * -holds for all j = 1, 2, ....m,then a;, a,, ... a;, a,, *.-Q2m=a ............ amm a';, a,, *-. a2m + p/ ............a?, a,, ... a1m a;, a22 ... a2m + ...+y ............ a;, am2 *.-amm Let A' be the transposition of A, i.e. Then A' = A. From eqn. (13) and applying eqn. (14) to all of the columns of A', we obtain D*(il, i,, ...,i,,,) = P2, il Pz. i2 ' * P2. i, ............ IPm,il Pm,i2 *** If any two of the indices i,, i, ..... i, are mutually equal, then the respective columns in eqn. (16) coincide, and, there- fore, the corresponding summand in eqn. (15) is equal to zero. We thus have to examine only those summands on the right- hand side of eqn. (15), in which the indices i,, i, ..... i, are all mutually distinct. Let (hl, h,, ....h,) be an ordered m-tuple of integers, such that 1 < h, < h, < -< h, < n.Then one of the summands in eqn. (1 5) will be of the form There will be additional m!-1 summands in which the indices i,, i,, .... i, coincide with h,, h,, ....h,, but not in that order. Then the determinant D*(i,, i,, .... i,) can be brought into the form D*(h,, h,, .... hd by a number of transpositions of columns of D*(i,, i2..... id. Each transposi- tion of two columns of a determinant causes the change of its sign. Consequently, it will be D*(i,, i,, .... i,) = (-ly'D*(h,, h,, ....h,,,) (17) where # is the number transpositions required in the mapping (il, i,, .... i,,,) +(h,, h,, .... hm). As known from algebra, fi is just the parity of the permutation (i,, i, ..... i,,,) relative to (hl, h,, .... h,,,). If we sum the m!terms on the right-hand side of eqn.(15) in which the indices i,, i,, .... i, (when appropriately ordered) coincide with (h,, h, ..... h,,,), and take into account eqn. (17),then we obtain nn D(i,, i,, .... i,,,)D*(i,, i,, .... i,)A= (rn!)-' 11 il=l i2=l i,=l nil ni2 -* nim This can also be expressed as Eqn. (19a)-(19c) are based on eqn. (1 1) and, thus, they rep- resent thermodynamic identities, whose range of applicability coincides with that of eqn. (ll), i.e. eqn. (19) holds for ideal systems. The first summation on the right-hand side of (19c) can be interpreted as being over all Hessian reactions, and thus we may write (19c)as where Y(=@)=-1 D(il, i2, .... im)D*(il,i2, ....i,) (21)nil ni2 -- - nimm im=l In the case of solutions, the terms (ndn,) in eqn.(12) can be considered as negligibly small. Then psi = vSi, and as a special case of eqn. (19b) and eqn. (21) we have and 1 D(i1, i2, .., i,)2Y(=@)=-m im=l (22)ni, ni2-nim In the case of solutions it is usual to exchange ni with [Ai], which can be done in view of the fact that the volume of a solution may be considered as constant. Then eqn. (20) and (22) simply reduce to eqn. (9) and (10). 3248 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Note that the term In order to interpret the Hessian determinant in terms of contributions from Hessian reactions, we first have to con- A(i,, i,, ...,i,-,) = ~(i,,i,, . . . ,iJ2 (23) struct the HRs. To do this, two of the chemical species from a i,=l [Aim] total of five are to be excluded from the above reaction occurring in eqn.(10) for r(X)can be understood as the scheme. If these two species are, for example, A, and A,, then Hessian determinant of the Hessian reaction X(il,i,, ..., by applying eqn. (7) we have: i, -l). To see this, recall that, in the case of a single reaction -1 -1 1 -1 -1 0 (m= l), the Hessian determinant is just G,,. On the other -1 -2 0 A3+ -1 -2 1 A4hand, eqn. (1l), when appropriately modified for solutions, yields -1 -3 0 -1 -3 0 -1 -1 0 + -1 -2 0 A, = 0 -1 -3 1 If the single reaction considered is a HR, then vli is given by eqn. (7) and the right-hand sides of eqn. (23) and (24) which is equivalent to: coincide. A, + A, = 2A4 When this procedure is repeated for all possible combinations Example 1 of two species, one arrives at the following set of HRs (in In order to illustrate the above results we consider a solution brackets are indicated the species which are eliminated from in which five chemical species A,, A,, A,, A, and A, react the initial SIRS): according to the following reaction scheme: (A,, A,) A, +A5 = 2A4 (i) A, + A2=A3 (AlY A31 A2 + A4 = A, (ii) A, + 2A2 = A, I (25) (iii) A, + 3A2 = A, (4, A4) 2A2 + A3 = A, The above scheme may, for example, represent a stepwise (A,, A,) A2 + A3 = A4 complex formation or a protonation process : (A2, A,) A, + 2A5 = 3A4 (i) M+ L=ML (A2, A4) 2A1 + A5 = 3A3 (ii) M + 2L = ML2 (A23 A,) A, + A4 = 2A3 (iii) M + 3L = ML, (A3, A4) A, + 3A2 = A, The Hessian determinant for eqn. (25) is evidently of the third order : (A33 A,) A1 + 2A2 = A4 (A4, + = Gll G12 G13 To each of these reaction equations we associate a function G12 G22 G23 A(il, i2), eqn.(23):G13 G23 G33 1 4with A(1, 2) = -+ -+-1 [A31 CA4I CAsI1G --+-1 +-1 1 1 l1 -CAiI [A21 CA3I A(1, 3) = -+ -+-1 CA2I CA4I CAsI1 4G22 = -+ -+-1 4 1CA1I [A21 CA4I A(1, 4) = -+ -+ -1 [A21 [A31 CA5I1 9G33 =-+ -+ -1 1 1[A11 CA2I CASI A(1, 5) = -+ -+-1 CA2I [A31 CA4I1G12= GZ1 = -+ -2 1 9LA11 [A21 A(2, 3) = -+ -+-4 CA1I CA4I CAsI 4 9A(2, 4) = -+ -+-1 CAI] [A31 CA5I 1 4A(2, 5) = -+ -1 + -CA2I CA4I CA3I The stoichiometric matrix for the reaction scheme eqn. (25) 1 9reads : A(3, 4) = -+ -+-1 [A11 CA2I [As] 1 4A(3, 5) = -+ -+-1 ;'-1 -1 1 0 CAI] CA2I CA4I /i -1 -2 0 1 1 1A(4, 5) = -1 + -+ -;I,'-1 -3 0 0 CAiI [A21 CA3I J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Then the Hessian determinant for the reaction scheme eqn. (25) is given by Thus, in this case, the Hessian determinant is equal to the sum of contributions from the ten Hessian reactions. Some Basic Notions of Sensitivity Analysis Consider the system of stoichiometrically independent reac- tions (SIRS) given via eqn. (1). Let tj,Kj, Aj H" and Aj V" stand for the extent, equilibrium constant, standard enthalpy and standard volume change, respectively, of the jth chemical reaction, j = 1, 2, . . . ,rn. Let [Ailo and [Ai] be the initial and equilibrium concentrations of the species Ai, and let pLp be its standard chemical potential, i = 1,2, ... ,n.Then, we have: m cAi1 = CAiIO + vki tk (26) &=1 n n / in The eqn. (27) may be considered as an implicit system of equations in the unknowns tj,[Ailo, Kj, pLp, T and P, i = 1, 2, ..., n, j = 1, 2, ..., rn. The latter five quantities are parameters influencing the position of chemical equilibrium. In what follows we denote them briefly by X. Then, eqn. (27) may be written as : fi(t1,52, ..., tm, j= 1, 2, ..., (28) where we use the abbreviation n in From eqn. (28) we have m aj-. at1 Jk= ah. j = 1, 2, ..., rn k=l atkax ax' Differentiating eqn. (28) with respect to the extents of the reactions (l), we obtain and thus, the basic problem of the sensitivity analysis is to solve the following system of rn linear equations in unknowns atj/ax,j = 1, 2, ...,m: The solution of eqn.(29) is given as: j = 1, 2, ..., m (30) where Ajp is the minor of the Hessian determinant obtained by deleting its jth row and pth column. The sensitivity coefficients a[Ai]/aX are now readily calcu- lated by differentiating the mass-balance conditions, eqn. (26): at.aCAi1 =-aCAilo + Cvjil; i = 1, 2, ..., nax ax j=l ax where atj/aX are given by eqn. (30). To obtain more specific results we have to specify the parameter X. The derivation of the expressions for the sensitivity coeffi- cients of the specified parameters X is analogous to that used in the preceding section for the Hessian determinant.There- fore, in the following we present only the main results. Effect of Temperature In the case when X is the temperature T, we have from eqn. (27) and (28) Substituting eqn. (31) into eqn. (30) and using a similar way of reasoning as in the preceding section, we arrive at 1atjaT RT~A where D(il, i, ,...,im-AH")is the following determinant of order rn D(il, i,, ..., im-l, AH") ... and Dimis the minor of the determinant D(il, i,, ..., im-l, AH"), obtained by deleting itsjth row and mth column. Eqn. (33) should be compared with eqn. (7). Through a similar derivation it can be shown that for the sensitivity coefficients of the equilibrium concentrations, the following is valid : (34) where D(il, i,, .. . ,im-1, i) is defined uia eqn. (5). Eqn. (34) has a simple chemical meaning. We have already shown that D(il, i,, ...,im-1, i) is just the stoichiometric coef- ficient of the species Ai in the Hessian response reaction if = &(il, i,, . . ., irn-l),whereas D(il, i,, . . ., im-l, AH") is the enthalpy change of the same response reaction. Thus we can write eqn. (34) in the form 1----Ic(%)vAJf')AH"(Jf') (35)aT RT2, 3250 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 Hessian reactions for Example 2 reaction eliminated species HR AHlkJ mol-' H+, OH- C03'-+ H2C03= H+,C032-HC0,-+ H20 = H+, HC03- C032-+ 2H20 = H+, H2C03 C032-+ H20 = OH-, C03'- HC03-+ H+ = OH-, HC03- 2H+ + C032-= OH-, H2C03 H+ + C03'-= C03'-, HC03- H20 = C032-, H2C03 H20 = HC03-, H2C03 H20 = 2HC03-2A2H" -A3 H" = -5.71 H2C03+ OH-A,H" -A2 H" + A, H" = 47.19 H2C03+ 20H-2A,H" + A3 H" = 88.67 HC03-+ OH-A,H" + A2 H" = 41.48 H2C03 -A2 H" + A3H" = -9.37 H2C03 A3 H" = -24.45 HC03-A2 H" = -15.08 H+ + OH-A,H" = 56.56 H+ + OH-A,H" = 56.56 H+OH-A,H" = 56.56 11:: where 4%)= {A x [Ail] x CAiJ x *.* x CAi,-iI)-' (36) Observe that 4%) depends on the equilibrium concentra- tions of the species Al, A,, ..., A,, whereas vi(M) and AH"(&')are concentration-independent. Because A is always positive-valued, it is clear that also 4%)is positive-valued for all Hessian reactions &'. The determinant D(il, i,, ..., i,,,-l, i) = vi(&')is equal to zero whenever i = il or i = i, or .. . i = i,-'. Consequently, vi(&')= 0 whenever the Hessian reaction &' does not involve the species Ai. We thus see that only those HRs in which the chemical species Ai is involved have non-zero contributions to the summations on the right-hand sides of eqn. (34) and (39, i.e. non-zero contributions to the sensitivity coefficient a[AJ/aT. The sign of the contribution of any particular Hessian reaction &' is easy to determine: it is equal to the sign of the product vi(&')AH"(H). Example 2 Consider the proton-carbonate system : (1) H,O = Hf + OH-AIH" = 56.56 kJ mol-' (2) H+ + C03,-= HC03-A, H" = -15.08 kJ mol-' (3) 2H+ + C03,-= H2C03 A3 H" = -24.45 kJ mol-' The stoichiometric matrix for this system reads : H+ OH-H,O C03,-HC03-H2C03 1 1 -1 0 0 0 0 -1 1 110 0 -1 0 1 We first find the Hessian reactions by eliminating all pos- sible groups of rn -1 (= 2) species from the initial set of inde-pendent reactions.Concomitantly, using eqn. (33) the enthalpy changes of these reactions are calculated. For instance, when the species to be eliminated are H+ and OH-, the resulting reaction equation is: 1 1 -1 1 10 -1 0 0 H,O+ -1 O -I C03,--2 0 0 -2 0 -1 110 110 + -1 0 1 HC03-+ -1 0 0 H2C03=0 -2 0 0 -2 0 1 1 1 AIH" AHo(Hf,OH-) = -1 0 A2 H" = -2A2 H" + A3 H" -2 0 ASH" Continuing the procedure we arrive at the 10 HRs given in Table 1, of which only eight are distinct. From this data it follows that the sensitivity coefficients a[H']/dT and a[OH-]/aT are positive.For instance, H+ is involved in reactions (v)-(x) and in each of these reactions the product of its stoichiometric coefficient and the enthalpy change is positive. Analogous considerations reveal that the sensitivity coefficients c?[CO,~ -]/dT, d[HCO, -]/dT and d[H2C03]/aT may be both positive and negative. From the 10 Hessian reactions in Table 1 one can easily derive expressions for the sensitivity coefficients. To do this, the product of the stoichiometric coefficient and enthalpy change of an HR has to be divided by the equilibrium concen- trations of those species which were eliminated from that HR, and the results summed over all the Hessian reactions in which the species under consideration is involved.For example: a[H+] 1 9.37 2 x 24.45 [OH-][C032-] [OH-][HCO,-]+ 15.08 56.56 [OH-][H,COJ + [CO, 2-][HC03 -1 56.56 [C032-IL-H,C031 + CHCO,-IL-H,CO,I Some numerical calculations for this system are presented in Fig. 1 and 2. In a solution of strong acid or base and in the absence of carbonate, the temperature dependencies of H+ and OH- are given by dCOH-1 d[H'] -[H+][OH-] AIH"-=--dT dT [H'] + [OH-] RT2 As a function of pH, both d[OH']/dT and d[H']/aT have a maximum at pH = 7. It is seen from Fig. 1 that in the pres- 0.030 10.025 O.OIO 0.008 --. n5 0.020 + 0.006n & 0.015 0, rg Q5 0.010 0.004 5 0.005 0.002 0.000 0.000. ~~~ 2 4 6 8 10 12 PH Fig. 1 Temperature derivatives of [H'] and [OH-] as functions on pH in the C032--H+ system.[C0,2-] + [HC03-]+ [H2C03] = 0.1 mol drn-'. (a)a[H+]/aT; (b)a[OH-]/aT. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 325 1 equilibrium heat capacity, while the second sum is the relax- 3c ation contribution : 2 -2 " 6 7 8 9 10 11 12 PH Fig. 2 Temperature derivatives of [CO,'-], [HCO,-] and [H,CO,] as functions on pH in the C03*--H+ system; notice that their sum is equal to zero. [CO,'-] + [HCO,-] + [H,CO,] = 0.1 mol dm-,. 1, X = HCO,-; 2, X = H'CO,; 3, X = CO,*-. ence of carbonate, the peak of a[H+]/aT is shifted to the acidic region while that of a[OH-]/aT is shifted to the basic region. Note the extremely low value of a[H']/dT. Effect of Pressure In the case when pressure, P,is the parameter influencing the chemical equilibrium, we have from eqn.(27) aln K. AVOafi -afi --1=I ax -ap ap RT It follows that eqn. (31)-(35) may be applied also in the case of pressure: it is only needed to exchange the terms A,H" by -TAj Voin all formulae. Chemical Equilibria Contributions to Heat Capacity Owing to the temperature dependence of the chemical equi- libria, the heat capacity is the sum of two contributions. The first is the conventional compositional contribution (Cp,camp), while the second is the relaxation contribution (Cp,,,J. Recent studies16-" show that Cp,rel may range from negligibly small to several hundreds of J K-' mol-l. The above derived equations can be used to interpret and evaluate this latter contribution.The total enthalpy of a homogeneous system can be expressed in terms of the partial molar enthalpies, H(Ai), of the components and the amounts, n,,of each component at equilibrium: n H = 1niH(Ai) i= 1 We deal with molar concentrations and, assuming that the volume of the solution is temperature independent, we con- sider the enthalpy in a unit volume of solution 4 n H = 1[Ai]H(Ai) (37)v i=l Differentiation of eqn. (37) with respect to temperature at constant pressure leads to a formula for the heat capacity at equilibrium The first sum on the right-hand side of eqn. (38) is identified as the conventional compositional contribution to the total where C,(Ai) stands for the molar heat capacity of the species Ai, i = 1, 2, ..., n.Here we are principally concerned with C, re1 * In order to obtain the relaxation contribution to the heat capacity, the mass-balance conditions, eqn. (26), are to be dif- ferentiated with respect to temperature: a[Ail C vji atj-= aT j=1 Combining eqn. (39) and (40) we obtain where Aj H"is the standard enthalpy change of jth SIR [see eqn. (l)], satisfying the condition n AjH" = 1vijH(Ai) i= 1 Taking into account eqn. (32), the relaxation contribution to heat capacity becomes 1 D(il, i,, ..., i,,,-l, AH")'7Cp,re1 = -1 cRT'A il<i2<...<im-l [Ail] x x *** x CAi,,,-,I (41) In order to deduce eqn. (41) we used the identity m l(-l~+"DjmAjH"= D(il, i,, ..., AH") j= 1 In analogy to (35), formula (41) can be written as Hence, also Cp,relis equal to a sum of contributions coming from Hessian response reactions. In contrast with eqn.(35), each summand in eqn. (42) is non-zero (provided that AH"($f) differs from zero). Example 3 We evaluate the relaxation contribution to heat capacity for the system NH,-H+, for which the following data are avail- able: H,O = H+ + OH-; A,H" = 56.56 kJ mol-' NH, + Hf= NH,+; A, H" = -52.22 kJ mol-' As the concentration of H,O is constant we have to consider only four chemical species: H', OH-, NH, and NH4+. The stoichiometric matrix for the above reaction scheme reads : H+ OH-NH, NH,' H20 1 1 0/1 -I-l 0 -1 O1 0 Eliminating consecutively from the above reaction equa- tions H+, OH-, NH, and NH4+, we arrive at the following PH Fig.3 Relaxation contribution to heat capacity in the NH,-H+ system. [NH,] + [NH,'] = 0.1 mol dm-,. Hessian reactions : (i) (H+) NH, + H,O = NH,+ + OH-A,H" + A2H0= 4.34 kJ mol-' (ii) (OH-) NH, + H+ = NH4+ A2 H" = -52.22 kJ mol-' (iii) (NH,) H,O = H+ + OH-A,H" = 56.56 kJ mol-' (iv) (NH4+) H20 = Hf + OH-A,H" = 56.56 kJ mol-' We thus see that the equilibrium in this system is influenced by four HRs [although reactions (iii) and (iv) are identical]. In accordance with eqn. (41), the relaxation contribution to the heat capacity is given by (4.34)2 (-52.22)2 (56.56)2 (56.56)2 +[H'] [OH-] +-[NH,] +-][NH, 3 The dependence of Cp,rel on pH is shown in Fig. 3.Discussion and Concluding Remarks The main finding reported in this paper is that various quan- tities, occurring in the sensitivity analysis of multiple chemi- cal equilibria in solutions, can be expressed as linear combinations of contributions that are associated with certain chemical reactions. Results of this kind apply both to the Hessian determinant [see eqn. (9)] and to the sensitivity coeficients [cf: eqn. (34)]. In view of this we speak about 'Hessian response reactions', HRs. We mention in passing that in the sensitivity analysis of gas-phase equilibria, response reactions other than Hessian are encountered.' The Hessian determinant is shown to be the sum of certain well defined contributions from all HRs. The same applies to the sensitivity coefficients of a chemical species Ai, except that here only those HRs in which Ai participates have non- zero contributions.We applied our general theory to the effects of temperature and pressure on the position of multiple chemical equilibria. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 These effects may be quite complicated. It was shown, however, that the overall effect is equal to the sum of contri- butions from HRs, having very simple forms: In the case of temperature each such contribution is equal to the product of the stoichiometric coefficient of Ai in the respective HR, the enthalpy change of that HR and of a concentration-dependent term; the latter is always positive. Fully analogous results hold also in the case of pressure.This makes it pos- sible to deduce easily the sign of each such contribution. On the basis of the above considerations the following pro- cedure for analysing the effect of temperature (pressure) on the complex chemical equilibria may be proposed. First, we have to derive all the Hessian reactions. Then, by means of eqn. (33) we evaluate the enthalpy change (volume change) of each HR. In order to determine the sensitivity coefficient of the chemical species Ai we select those HRs in which Ai is involved. For each such HR the stoichiometric coefficient of Ai is to be multiplied by the enthalpy change (volume change) and by c(&), eqn. (36). This yields the contribution of the given HR. The sensitivity coefficient is then the sum of such contributions over the selected Hessian reactions.References 1 W. R. Smith and R. W. Missen, Chemical Reaction Equilibrium Analysis: Theory and Algorithms, Wiley, New York, 1982. 2 M. Beck and I. Nagypal, Chemistry of Complex Equilibria, Aka-demiai Kiado, Budapest, 1990. 3 A. A. Bugaevski and B. A. Dunai, Zh. Anal. Khim., 1972,26,205. 4 A. A. Bugaevski and L. E. Rudnaya, Zh. Neorg. Khim., 1976, 21, 2827. 5 I. Nagypal, I. Paka and L. Zekany, Talanta, 1978,25, 549. 6 A. Avdeef and K. N. Raymond, Inorg. Chem., 1979,18,1605. 7 A. A. Bugaevski and L. E. Nikishina, Zh. Neorg. Khim., 1980,25, 2854. 8 I. Nagypal, M. T. Meck and A. Zuberbuhler, Talanta, 1983, 30, 593. 9 L. Zekany and I. Nagypal, in Computational Methods for the Calculation of Stability Constants, ed. D. Leggett, Plenum Press, New York, 1985. 10 I. F. Fishtik and I. G. Povar, Zh. Neorg. Khim., 1990,35,102. 11 I. F. Fishtik and I. G. Povar, Zh. Neorg. Khim., 1990,35, 108. 12 A. Braibanti, E. Fisicaro, F. Dallavalle, J. D. Lamb and J. L. Oscarson, J. Phys. Chem., 1993,9l, 8062. 13 I. Fishtik, 1. Nagypal and I. Gutman, in the press. 14 I. Prigogine and R. Defay, Chemical Thermodynamics, Long-mans, London, 1954. 15 G. G. J. Mains, J. W: Larson and L. G. Hepler, J. Phys. Chem., 1984,88,1257. 16 E. M. Woolley and L. G. Hepler, Can. J. Chem., 1977,55,158. 17 C. Jolicoeur, L. L. Lemelin and R. Lapalme, J. Phys. Chem., 1979,83,2806. 18 G. G. Allred, J. W. Larson and L. G. Hepler, Can. J. Chem., 1981,59,1068. 19 J. C. Peiper and K. S. Pitzer, J. Chem. Thermodyn., 1982,14,613. 20 J. W. Larson, K. G. Zeeb and L. G. Lepler, Can. J. Chem., 1982, 60,2141. Paper 4/02175D; Received 12th April, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003245

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3253-3259 3253 Effectof Complexation on the Excited State of Uranyl p-Diketonato Complexes: Application of the Energy Gap Law Tomoo Yayamura, Sugio Iwata, Shun-ichi lwamaru and Hiroshi Tomiyasu Research Laboratory for Nuclear Reactors, Tokyo Institute of Technology, O-okayama , Meguro-ku, Tokyo 152,Japan Excited-state properties of four series of dioxouranium(vi) (uranyl) fl-diketonato complexes UO,(L-L), thf, UO, (L-L), d mso [L-L = acac (acetylacetonate), tfacac (trifIuoroacetylacetonate), hfacac (hexafIuoroacetyl-acetonate), btfa (benzoyltrifluoroacetylacetonate), ba (benzoylacetylacetonate), dbm (di benzoylmethanide) and tta (tenoyltrifluoroacetylacetonate) ; thf = tetrahydrofuran, dmso = dimethyl sulfoxide], UO,(acac), L and UO, L,(CIO,), [L = dmso, dma (N,N-dimethylacetamide), hmpa (hexamethylphosphoric triamide), dmf (N,N-dimethylformamide) and tmp (trimethyl phosphate)] have been studied at 296 and 77 K.The lifetimes of the excited states of UO,(L-L), thf were first measured at ambient temperatures in the liquid solvent and range from 0.90 ns for L-L = acac and ba, to 484 ns for L-L = hfacac. Good correlations between €,,, (0'4''band energies in the first electronic transition) and the lifetimes of the excited states, and also between€,,, and v3 (asymmetric vibration) of U-0 in uranyl ion, were obtained. The energy gap law accounts for €,,, being the major perturbing factor of the lifetimes. One of the major themes in photochemistry is to understand from uranyl nitrate hexahydrate [U0,(N03), -6H20]" the nature of the charge-transfer excited state and its applica- and purified by recrystallization using methylene tion to a sensitizer.' Therefore it is desirable to design energy chloride.UO,(acac),dmso, UO,(acac),dma, UO,(acac),dmf, levels, emission quantum yields and lifetimes of the excited UO,(acac),hmpa and UO,(acac),tmp were prepared by dis- state by using a variety of ligands. solving UO,(acac), ,which was obtained by the sublimation Since the work by McGlynn and Smith,, the electronic of UO,(acac),H,O at 403-413 K, in DMSO, DMA, structure of uranyl complexes has been investigated. The DMF, HMPA (hexamethylphosphoric triamide) and TMP molecular orbitals of uranyl ion have been discussed on the solvent, respectively.UO,(btfa),dm~o,'~ UO,(ba),drns~,'~ multiple-scattering Xa UO,(dbm),drns~'~ and UO,(tfacac),dmso' were prepared basis of the electronic ~pectrum,~ meth~d,~the relativistic Dirac-Slater molecular orbital (MO) by the method described above. UO,(a~ac),thf,'~ method5 and the relativistic extended Huckel method;6 the UO,(tfacac), thf,I3 UO,(hfacac), thf,14 UO,(dbm),thf," highest occupied molecular orbital (HOMO) was ascribed to UO,(ba),thf l2 and UO,(btfa),thf l2 were prepared by a 0, and the lowest unoccupied molecular orbital (LUMO) to method described previously. The formation of these com- 6,. Although the electronic structure of the excited state is plexes was confirmed by comparing IR and UV-VIS mea-still uncertain, a good approximation to the energy level was surements with data reported previously.' ',' given by the calculation of the excited-state energy assuming DMSO, DMA, DMF, HMPA and TMP (Tokyo Kasei or that one electron is transferred into the LUMO. Wako) were reagent grade and used after distillation under It is important to have fundamental data available on how reduced pressure.All solvents [ethanol, carbon tetrachloride, neutral monodentate or anionic bidentate ligands govern the 1,2-dichloroethane, methylene chloride, chloroform, benzene, nature of the photoexcited states of uranyl complexes, such as cyclohexane and hexane (Kanto Kagaku)] were of optical the energy levels, Eo,o, and lifetimes of the lowest excited grade and used without further purification.In nanosecond states. flash photolysis measurements, sample solutions of CCl, , Recently, excited-state lifetimes of ruthenium and osmium CH,Cl,, hexane or cyclohexane were degassed more than complexes were successfully explained by the energy gap five times by the freeze-pumpthaw method. la^.^-^ In view of the earlier studies, it is also expected that the effects of coordination and solvation on the lifetimes of Measurements of IR and W-VIS Spectra the excited states of the uranyl ion may be interpreted by the UV-VIS spectroscopy was carried out on a Shimazu UV-365 energy gap law. spectrophotometer with a thermostatted cell compartment (at This paper presents the results of a systematic study of the 298 K) using a 1 cm quartz cell.The concentration of solu- excited states of four series of uranyl complexes; tions was adjusted to lop4 mol dm-3. IR spectra wereU0,Ls(CI04), , UO,(acac),L (L is a monodentate neutral recorded on a Hitachi 270-30 spectrometer in the 4000-400 ligand e.g. thfj, UO,(L-L),dmso and UO,(L-L),thf (L-L is a cm-' region with KBr disks, by the Nujol method, at bidentate B-diketonate ligand e.g. acetylacetonate). Most of ambient temperature. the spectroscopic and kinetic measurements were carried out at 77 K in a rigid matrix, while nanosecond flash photolysis Emission and Excitation Spectra was used at 296 K to determine the lifetimes of UO,(L-L), thf. A Hitachi 850 fluorospectrophotometer was used for emis- sion and excitation spectrophotometric measurements.Mea- surements at 77 K were taken with a cryostat, or with an Experimental EPR tube inserted into a Dewar vessel, and both the cryostat Materials and the Dewar vessel were filled with liquid nitrogen. In a series of U0,L5(C104)2, UO,(acac),L and UO,(L-L),dmso U02(dmf)5(C104)2 9 U02(dma),(C104)2 7 U02(dmso)S(C104)2 samples, the measurements were carried out in ethanol, and UO,(hmpa),(ClO,), were prepared by a method ,re- whereas isopentane-methylcyclohexane was used as a mixed ported in an earlier paper." UO,(acac),H,O was prepared solvent for UO ,( L-L) ,thf. Quantum Yield Measurements Emission quantum yields, 4,,, were measured at 77 K in degassed ethanol solutions by the freeze-pumpthaw method. Observed quantum yields were corrected by using a standard, degassed cyclohexane solution of perylene16 (4em= 0.78).Correction was also made to the refractive index, n, using Eykmann’s empirical formu1a”t and eqn. (1). Flash Pbotolysis Measurements Microsecond flash photolysis measurements were carried out at 77 K by using a giant-pulse ruby laser (Toshiba; 20 ns pulse width) and oscillator wavelength, 694.3 nm) with a KDP (KH,PO,) crystal for conversion into 347.2 nm optical second higher harmonics, and a PIN silicone photodiode as a trigger, a monochromator (Ritsu Ohyoh Kogaku P-2ON) and a photomultiplier (Hamamatsu Photonics R1509). Nanosecond flash photolysis measurements were carried out at ambient temperature by means of the time-correlated photoelectron-counting method using a Horiba NAES-500 flash photometer. Samples were irradiated by a Horiba N2-dye laser NDL-100 with Stilbene 3 (at 424 nm) and BiBuQ (at 390 nm).Results Spectroscopic data of a series of complexes, UO,(L-L),thf, UO,(L-L),dmso, UO,L,(ClO,), and UO,(acac),L, are sum- marized in Tables 1 and 2, together with the kinetic data. Uranyl complexes exhibit characteristic absorption bands in the wide range of 400-500 nm, showing well defined vibra- tional structures depending on the coordinated ligands. Although the absorption bands attributed to the uranyl ion are observed in the series of U02L,(C10,), , they are more or less masked by charge-transfer bands’ of uranyl /I-diketonate complexes. Thus the energy gap between the 0 vibrational level of the ground state and the 0 level of the excited state (E0,*) are determined by emission spectroscopic measurements in ethanol or the isopentane-methylcyclo-hexane mixture at 77 K and in carbon tetrachloride at 296 K.The spacings of these vibrational structures correspond to vl, the symmetric vibration of uranyl U=O.’* v3 and vl(C=O) are the asymmetric vibrations of the uranyl U=O and the symmetric stretching vibration of the fi-diketonate C=O, respectively, determined at ambient tem- perature by the Nujol method using IR spectroscopy. The symmetric and the asymmetric stretch vibration frequencies of the uranyl moiety exhibit significant shifts with both /I-diketonate ligands and neutral ligands (Fig. 1 and Fig.2). Irradiation at 347.2 or 337 nm initiates the ligand-centred n-n* transitions of the /I-diketonate ligands. Observed emis- sion spectra in the range of 500-600 nm are assigned to the excited state of the uranyl moiety.” Thus, the emission shows that the intramolecular energy transfer occurs from the f According to Eykmann’s empirical formula, n(n2-1P(n= c[n(T)+ 0.4) the following formula is gained using a first-order approximation on the formula; n(T)= (no + An) = (no + ca(T -To)) (c and a are parameters, c = 0.79360, a = 0.0011). The corrective coeflicient (ncthsnol/ncyclohcrpnc)2= 0.6767 is calculated from ni:2g$;:34*7 = 13 700 and n~~~,,~ohexanc= 1.4362. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 950 900 850 3-I E 4 ‘L5 800 750 700 2 4 6 8 10 12 14 16 PK, Fig.1 Correlation of symmetric stretch frequencies with pK, of fi-diketonate ligands. [O, UO,(L-L),thf, 77 K in isopentane-methylcyclohexane; 0, UO,(L-L),thf, 77 K in CCl,; A, UO,(L-L),dmso, 77 K in ethanol.] Numbers are cited in Table 1. ligand-centred (LC) excited states, localized on the /3-diketonate ligand, to metal-centred (MC) excited states. The intensity of the emission of uranyl compounds varied with the bi-dentate ligands in the following order; hfacac > tfacac > tta > btfa > acac > ba > dbm. Although the emission spectra are observed in all compounds at liquid- nitrogen temperature, those of acac, ba and dbm are not observed by the conventional spectroscopic method at 296 K in liquid solutions.The lifetimes, zo,of UO,(L-L),thf were first determined in liquid solutions at an ambient temperature of 296 K by nano- second flash photolysis. The value of zo at 296 K for UO,(L-L),thf was greatest for L-L = hfacac (484 x s) 950 I I I I I I 1 900 850 r I L,E -% a \a 800 750 700 I I I I I I I 10 15 20 25 30 35 40 45 50 DN Fig. 2 Correlation between the symmetric stretch frequencies of the uranyl U-0 bond and the donor number (DN) of coordinated neutral ligands. 10, UO,(acac),L at 77 K in ethanol; 0, U0,L5(CI0,)2 at 77 K in ethanol.] Alphabetical codes are cited in Table 2. Table 1 Spectroscopic and kinetic data of UO,(L-L),thf and UO,(L-L),drnso complexes ~~ ~ UO,(L-L), thf UO2( L-L),dmso E0,o '0.0 (296 K) 'O(295 K) E0,o '0 (77K) L-L pK,(L-L)" /lo3 cm-' /lo3cm-' vl/cm-' v3/cm-' sc ko/s-' ' /lo3 m-l d vl/m-l s k,/io2 s-~ k,,/io3 S-' 1 hfacac 6.00 19.16 19.05 858 940 484 2.07 x lo6 2 tfacac 8.33 18.98 19.01 843 930 30.6 3.27 x lo7 19.03 840 250 0.22 8.8 3.1 3 tta 9.10 19.05 18.94 936 41.1 2.43 x lo7_J 4 btfa 9.14 18.80 18.87 828 926 49.2 2.03 x lo7 18.99 833 148 0.15 10.0 5.7 5 acac 12.70 18.75 n.0.' 2 922 0.9 1.11 x 109 18.75 826 206 0.17 8.3 4.0 6 ba 12.85 18.62 n.o.e 814 914 0.9 1.11 x 109 18.69 817 148 0.11 7.4 6.0 c 7 dbm 13.75 18.52 n.0.' 805 908 -B _-B 18.56 813 103 0.11 11.0 8.6 r 0 In a solution of dioxane-water (75: 25, w/w%) at 303 K (H.Yokoi and T.Kishi, Chem.Lett., 1973, 749). 'Measured at 77 K in isopentane-methylcyclohexane. Measured at 296 K in CCl,. 8 Measured at 77 K in ethanol. n.0.: not observed. Vibration spectra are not suficiently well defined to determine vl, because those complexes are insoluble in CCl, . t~The lifetime at 296 K for the dbm complex was too short (z < s) to be measured by the present equipment. Table 2 Spectroscopic and kinetic data of U0,L5(C10,), and UO,(acac),L complexes u02L5(c104)2 UO ,(acac),L a tmp 23 18.74 818 194 0.17 8.8 4.3 b dmf 24 19.28 856 219 0.27 1.2 3.3 18.76 824 206 0.17 8.3 4.0 c dma 27.3 19.10 844 172 0.31 1.8 4.0 18.76 808 181 0.17 9.4 4.6 d dmso 29.8 19.01 835 168 0.24 1.4 4.5 18.75 826 206 0.17 8.3 4.0 e hmpa 38.8 19.35' 848' 357 0.15 0.4 2.4 18.76 811 190 0.17 9.0 4.4 ~~ ~ V.Gutman and E.Wychera, Znorg. Nucl. Chem. Lett., 1975, 11,635. Measured at 77 K in ethanol. 'Four-coordinate complex, i.e. U02(hmpa),(C10,)2. and decreased in the same order as described for the emission intensity. The lifetime of the dbm complex at 296 K was too short (zo < lo-’ s) to be measured by the present equipment. Results of the lifetimes and quantum yields for the emission of UO,(L-L),thf are summarized in Table 1. The quantum yield q5em is described in terms of k, and k,, by eqn. (2) and (31, kd = k, + k,, where kd refers to the deactivation rate constant, which is the sum of two processes, k,, the radiative decay rate constant and k,, the non-radiative decay rate constant. The values of k, and k,, determined by eqn.(2) and (3), are shown in Tables 1 and 2. Discussion Effect of Coordination on the Stretch Vibration of the Uranyl Ion Fig. 2 shows v1 to have considerable dependence on the equatorial neutral ligands of UO,L,(ClO,), . The stretch vibration of the uranyl ion may be considered as a parameter of the complexing ability of the equatorial ligand toward the uranyl ion, as proposed by Jung et ~1.~’v1 decreases pro- portionally with increasing donor number (DN), which is a measure of the enthalpy of complex formation between the given ligand and antimony pentachloride in 1,2-dichloro-ethane, except for the four-coordinated complex of U0,(hmpa)4(C10,),.2 This correlation between v1 and DN is parallel to that between v1 and AGO, the Gibbs energy for the one-base exchange reaction of UO,(hfacac),thf relative to THF, reported by Kramer et In the series of UO,L,(ClO,), complexes that have been studied, only the hmpa forms a four-coordinate complex, UO,(hmpa),(C10,), , which has a larger v1 value than expected.Bray and Kramer2 showed that UO,(hfacac), and the neutral ligand, L, form a Lewis acid-base adduct, and v3 of the adduct complex is a good parameter for determining its Gibbs energy relative to THF. According to their idea, the Gibbs energy of complex formation of UO2LS2+ is related to the enthalpy of complex formation between the given base and SbCl, in CH,ClCH,Cl, reflected in DN, and the Gibbs energy of complexation for UO,(hmpa),’ will be smaller + than expected from DN.Unlike UO2LS2+,v1 changes very little with L for UO,(acac),L (Fig. 2) because the anionic acetylacetonate ligand has a much larger complexing effect toward the equa- torial plane of uranyl than the monodentate neutral ligand, and hence the substitution of L yields little perturbation on the equatorial plane of UO,(acac),L. In the case of the uranyl /I-diketonato complexes, UC,(L-L), thf and UO,(L--L),dmso, the symmetric stretch frequency, v,, of the uranyl U=O bond displays marked shifts according to the complexing ability of the b-diketonate, which is evaluated from the pK, values of the B-diketone. Indeed, as seen in Fig. 1, a plot of v1 us. pK, gives a linear relationship, indicating that an increase in the basicity of the P-diketonate weakens the uranyl U-0 bond.A linear corre- lation between the pK, of a series of ligands and log K for a certain metal ion (where K is the complex formation constant) is often found, as seen in the copper P-diketonato complex.24 This relationship can be understood on the basis that the complex formation reaction involves the substitution of a proton by a metal ion, thus the ligating ability increases J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 4 900 920 940 960 vJcm -’ Fig. 3 Correlation of symmetric stretch frequencies of the uranyl U=O bond at 77 K in isopentane-methylcyclohexane (right axis) and the carbonyl C-0 bond of B-diketonate ligands at 296 K in a Nujol mull (left axis) with the asymmetric stretch frequencies of uranyl U-0 bond for UO,(L-L),thf complexes as the formation constant between the proton and the ligand increases.Owing to O=U=O being linear, the asymmetric stretch is IR-active while the symmetric stretch is Raman-active. The symmetric stretch, vl, observed in the UV-VIS spectra in a rigid matrix at 77 K corresponds to reported Rarnan data15 and correlates well with the asymmetric stretch, v3 , observed at ambient temperatures (Fig. 3). v3 also shows a linear relationship with the symmetric stretch frequency v1 (CEO) as shown in Fig. 3. Excited-state Energy Levels Unlike the small variation in the ligand-field splittings in a series of P-diketonato complexes of iron(m) and chromium(~~~),~’there exists a large variation in Eo,o in the uranyl complexes.This large difference in Eo,o, amounting to CQ. 6500 cm-’ in UO,(L-L),thf, could be attributed to the difference in the nature of the lowest excited states of the uranyl complexes and the d-transition-metal complexes. The lowest absorption band was attributed to the tran- sition from nu rather than that from nu by Denning et d3 Pyykko and Lohr determined the ionization energies, Ei, for nu and nu of UOZ2+ and U0,C142- from the relativistic extended Hiickel method,26 which supported the result of Denning et al. Based on the large value of the squared ampli- tude (59%) calculated for the 5f contribution to the numolec-ular orbital (HOMO) by Boring et ~l.,~Jarrgensen stated that the lowest excited state would be due to the transitions from a bonding 5f u,,orbital to non-bonding 5f orbitals27 in addi- tion to charge transfer from the uranyl oxygen to the central uranyl atom.Although the assignment of the orbital responsible for the lowest excited state and the mechanism by which the equatorial ligands perturb the energy levels of the uranyl MO are still controversial, it is valuable to elucidate how the equatorial ligand affects the excited energy levels. The lowest excited energy levels Eo,o are plotted against the symmetric stretch frequency v1 (Fig. 4), which appears to be a parameter of the complexing ability. The energy levels increase linearly with increasing symmetric stretch frequency for two series of complexes, UO,(L-L),thf and UO,(L-L),dmso.In contrast to the above complexes, the J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3257 This can be interpreted by Einstein's relation3' for sponta- neous emission [eqn. (4)]. 64n4v3 k, = -(4) l9 500 t 3hc3 m2 c I 8 On the assumption that the dipole moment matrix element sE 8 for transition, m2,is constant, then k, should vary as the cube A Y b of the emission energy. In fact, k, is 9.4 x 10' s-l with 19000 -A' 'A Eo,o= 18.76 x lo3 cm-' for UO,(acac),dma, which lies in c 9 the middle of the series of UO,(acac),L complexes, while for LIP UO,dmf5(C10,), , k, = 1.2 x lo3 s-' with Eo,o = 20.19 x lo3 cm-', which lies in the middle of the series of 18 500 U02L,(C104), complexes. This indicates that k, changes by ca.24% with the 7% change in Eo,o, close to the value of 23% expected from the cubic relationship of k, with respect to E0.0 * 18000780 800 820 840 860 880 The uranyl complex is known to destroy its organic ligand v, (77 K)/cm-' Fig. 4 Correlation between 0-0 excited energy levels and the sym- metric stretching frequencies of the U-0 bond. [A,UO,(L-L),thf; A, UO,(L-L),dmso; 0,UO,(acac),L and 0,UO,L,(CIO,),.] energy levels for a series of UO,(acac),L complexes are nearly constant, in the region of 20 cm-' in vl. The results shown in Fig. 4 are attributed primarily to the a(L-M) electron donation into the empty uranyl 4, or 6, orbitals which extend into the equatorial plane of the uranyl ion.23 The increased electron density on these antibonding orbitals destabilizes the O=U=O bond, because uranyl 4, or 6, orbitals are similar to atomic-like orbitals on the uranium atom and give little contribution to the O=U=O b~nding.'~.~~Therefore the decrease in v1 corresponds to a decrease in the bond strength, and hence to an increase in the bond distance of O=U=0.29 Simultaneously, Boring et al.pointed out that the separa- tion between unoccupied 4, and 6, levels and occupied levels of the uranyl ion increased with decreasing bond distance r(U=O).4 The bond distance of uranyl U=O, r(U-0), can be estimated by McGlynn's28 and formulae. On applying these formulae to a series of UO,(hfacac),thf, r(U=O) decreases from 1.740 to 1.727 A as the value of v3 increases from 908 to 940 cm-'.These values of r coincide with the usual value of ca. 1.71 A for the uranyl U=O bond. Thus, the above analysis explains the trends observed in the series of uranyl complexes. As the U=O bond distance decreases, the interaction between the atomic orbitals of oxygen and uranium in the forming uranyl molecular orbitals (MO) increases, and hence the energy gap between the occupied orbitals, which are localized mainly on the 2p orbitals of uranyl oxygen, and the unoccupied orbitals local- ized mainly on uranium 5f orbitals, decreases. Consequently, the increasing donation weakens the uranyl U-0 bond, and lowers the energy of the lowest excited state. Excited-state Lifetimes The values of the radiative rate constant, k,, calculated by eqn.(2) and (3) are listed in Tables 1 and 2. It can be seen in these tables that k, is sensitive to the emission energy levels. t According to McGlynn's f~rmula,~ force constants of the uranyl axial U-0 bond and asymmetric stretch frequency, v3, are related by Fu-o = (8.2558 x 10-8)(v3)2. Furthermore, Jones introduced the formula indicating the relationship between the bond length of U-0 and force constants of the U-0 symmetric stretch vibration as follows; ru-o = (5.013 x 10-'0)(Fu-d-'/3 + 1.17 x lo-''. photolytically and to exhibit photochemical reactions with organic compounds, and therefore the quantum yield of the uranyl emission is suppressed by these chemical pro-ces~es.~~-~~Many organic compounds are reported to react with *U02'+, including aromatic hydrocarbon^,^^ alco-hol~,~~ethers,35 aldehydes,36 organic acids37 and other com- pound~.~~,~'There is no evidence that saturated alkylketones effectively quench the excited uranyl ion.The reported value of the decay rate constant for the emission of uranyl per- chlorate is 5.2 x lo5 s-' in a 50% acetone-water mixture,40 which is comparable to the value 5.2 x lo5 s-' under the condition: [UOZ2'] = 0.01 mol dm-3 in aqueous solution, pH 1.0 and ionic strength 3.0.41 This indicates that acetone, the saturated alkyl ketone, has no significant effect on the quenching of the excited uranyl ion. Our interest lies in the extent to which the chemical deacti- vation of *UOZ2 leads to photolysis of B-diketonate uranyl + complexes. In order to examine this, the following experiment was carried out: a solution of UO,(acac),thf in cyclohexane ([UO,(acac),thfl = 1 x mol dm-3) was irradiated at 436 nm by using a high-pressure mercury lamp with filter L42 and B390.The absorption spectrum was unchanged during the irradiation time of 10 min. This result indicates that the quenching yield of the photochemical reaction is minimal in the uranyl b-diketonato complex. Note in Tables 1 and 2 that k, varies by more than two orders of magnitude with the emission energy gap, Eo,o, in UO,(L-L),thf and k,, varies by ca. 250% with Eo,o in UO,(L-L),dmso. As predicted by the energy gap law,42,43 a negative linear relationship between the non-radiative deacti- vation rate constants, k,, , and emission energy gaps, Eo,o, is observed (Fig.5 and 6). Caspar and Meyer reported7v8 the deactivation rate constants of MLCT excited states of ruthenium(I1) and osmium(I1) complexes based on the energy gap law. The energy gap law is given by eqn. (5) by Jortner and co-w~rkers~~~~~ YEo*oIn k,, = (In /?-&) --A% where y=ln--E0,o 1 (7)'OM sM and O, = 2nvM is the angular frequency of the deactivating vibration, Vz the electron tunnelling matrix element corre- sponding to the transition between two states and S,, the 3258 lo I I I I I 18000 19 000 20000 E0,olcm- Fig. 5 Correlation between the logarithms of non-radiative-decay rate constants of the excited state and 0-0 excited energy levels.A, UO,(L-L),dmso; 0,UO,(acac),L; 0,U02L5(C104), at 77 K in a methanol glass; 0,(UO,),Ni(L-L-L-L),py. at 77 K in methyl- tetrahydrofuran glass. (L-L-L-L: 1,7-diphenylheptane-1,3,5,7-tetra-one, see ref. 44.) half sum of the dimensionless fractional displacement of the normal mode between the equilibrium configurations of the ground and the excited states. Although there is no knowledge of the deactivating vibra- tion mode of the excited uranyl ion, it is reasonable to assume that the intense symmetric or asymmetric stretch vibration of the uranyl ion corresponds to the deactivating vibration. The logarithms of k,, are plotted against the emis- sion energy gaps Eo,o in the series of UO,(L-L),dmso, UO,(acac),L and UO,L,(ClO,), in the rigid matrix at 77 K (Fig.5). The values of k,, for /3-diketonato complexes decrease linearly with increasing Eo,o [see eqn. (5)]. The values of k,, for UO2LS2+ are slightly greater than those of 25 -20 -0 Y -c 15-10. 5. \ -~~ 18000 19 000 20 000 21 000 Eo,o/cm-l Fig. 6 Correlation of the logarithm of decay rate constants of the excited state of UO,(L-L),thf complexes and other complexes with the 0-0 excited energy levels at 296 K in liquid solutions. 10, UO,(L-L),thf at 296 K in CCl,. Numbers are cited in Table 1. For complexes 6 and 7, the values of Eo,o at 77 K in isopentane- methylcyclohexane are used. 0,(a) UO,F,(CIO,), , (b) UO,(H,O),(ClO,), at 296 K in aqueous solutions; from ref.41.1 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 k,, for the /3-diketonato complexes. Assuming that the slope of k,, vs. Eo,o is the same for both U02Ls2+ and for the /3-diketonato complexes, i.e. that the deactivating vibrations (CUM)are the same in both these complexes, lower values of k,, for the #I-diketonato complexes at the same Eo,o value com- pared with UO2LS2+ result in the smaller electron tunnelling matrix elementsg for the /3-diketonate complexes. This is also supported by the values of lifetime and Eo,o for the complex with the oligomer analogue of the /3-diketonate ligand re- .~~ported by Tamilarasan et ~1 The /3-tetraketonate complex, (U02)2Ni'yL-L-L-L)2py4, has two uranyl moieties per mol- ecule, and four carbonyl groups and two pyridine ligands coordinated to one uranyl moiety, whose ligating groups are quite similar to those of the P-diketonate complexes in our study.The kinetic and spectroscopic data of the tetraketonate complex reproduce the plot of the #I-diketonate complexes, but not the plot of the neutral ligand complex (Fig. 5). Values of In kd are plotted against Eo,o in a series of U02(L-L),thf complexes in CCl, solutions at 296 K (Fig. 6). In this figure, the value of kd can be satisfactory used in place of the value of k,,, since the radiative quantum yields in liquid solutions are quite small [+em = 0.0028 for UO,(hfacac),thfl. The energy gap law is also found to be applicable to UO,(L-L), thf in liquid solution. Excellent correlations between kd and Eo,o are found for /3-diketonato complexes except for those with monophenyl substituents.If the same slope is also assumed for two monophenyl-substituted /3-diketonato complexes, the smaller values of k,, are responsible for the stabilization of excited states because of the smaller value of the electron tunnelling matrix element for monophenyl substituents. Note that the uranyl fluoro complex has a considerably long lifetime, z = 0.149 x s ([UO,"], = 0.01 mol dm-3, [NaF] = 1.0 mol dm-3, pH 1.0, ionic strength 1.0):' compared with other uranyl complexes in aqueous solutions and a point from the plot of In kd vs. Eo,o coincides with the plot for the #I-diketonato complex (Fig. 6). The average coor- dination number of F-of the fluoro complex is almost five, if excess fluoride ion exists in solution,46 suggesting that U02Fs3- has no H20 molecules coordinated in the equato- rial plane of the uranyl ion.A lack of water molecules in the coordination sphere might affect the lifetime of the excited uranyl ion, because the high-frequency stretch vibration of the 0-H bonding in H,O is often observed to be responsible for the large rate constants of non-radiative deac- tivation in emissive lanthanide ions in aqueous Thus the behaviour of the excited state of the uranyl fluoro complex differs from that of the aqua complex U02(H20)4(C104)2 (Fig. 6).The values of the slopes in Fig. 5 and 6 are listed in Table 3. It was reported for ruthenium bipyridine complexes that the solvent effe~t,~ on life- as well as the complexing effe~t,~ times in the MLCT excited states was explained by the energy gap law.The slopes for the plots of In k,, 11s. Eo,o were similar for a variety of N-coordinate ligands in Ru(bpy),Ln2 + Table 3 Values of slopes and intercepts obtained from the plots of In k,, us. Eo,o reaction complex conditions slope/eV - ref. UO,( L-L) ,dmso UO ,(acac),L u02L5(c104)2 77 K in EtOH 77 K in EtOH -16.1 -16.1 this work this work UOJL-L), thf" Ru(bpy),L,'+ 296 K in CCI, 200 K in CH,Cl, -100 -7.3 this work b Values are derived from plots of In k, us. Eo,o.* See ref. 7. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3259 (L, N-coordinated ligands, n = 1, 2) and for a series of sol-vents dissolving Ru(bpy),’+.In comparison with the ruthe- nium complexes, the large values of the slopes for the uranyl complexes may be attributed to significant vibrational contri- bution. For U02(hfacac),thf complexes, intense vibrations, v3(U=O) = 956 cm-’, were reported in the vapour phase.” These high-frequency vibrations may be responsible for the effective deactivation of the excited state of uranyl ion. 24 25 26 27 28 29 30 M. 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Harada and H. Tomiy- asu, J. Chem. SOC., Faraday Trans., 1990,86, 55. R. Engleman and J. Jortner, Mol. Phys., 1970, 18, 145. K. Freed and J. Jortner, J. Chem. Phys., 1970,52,6272. R. Tamilarasan, T. Ramakrkhnan and J. F. Endicott, Inorg. 19 20 1976,22,262; 1968,25,312. Y-Y. Park, Y. Sakai, R. Abe, T. Ishii, M. Harada, T. Kojima and H. Tomiyasu, J. Chem. SOC., Faraday Trans., 1990,86,55. W-S. Jung, H. Tomiyasu and H. Fukutomi, Bull. Chem. SOC. 45 46 Chim. Acta, 1987,142,321. M. Moriyasu, Y. Yokoyama and S. Ikeda, J. Znorg. Nucl. Chem., 1977,39,2199. M. Harada, Y. Fujii, S. Sakamaki and H. Tomiyasu, Bull. Chem. 21 Jpn., 1985,58,938. G. J. Honan, S. F. Lincoln and E. H. Williams, Znorg. Chem., 47 SOC.Jpn., 1992,65, 3022. J. L. Kropp and M. W. Windsor, J. Chem. Phys., 1963,39, 2769; 1978,17,1855. G. Stein and E. Wurzberg, J. Chem. Phys., 1975,62,208. 22 G. M. Kramer, E. T. Maas Jr. and M. B. Dines, Znorg. Chem., 23 1981,20,1415. R. G. Bray and G. M. Kramer, Znorg. Chem., 1983,22,1843. Paper 4103614J;Received 14th June, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003253

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3261-3265 -EPR Study of NaCl :CO,' and NaCl :SO, Peter D. Moens,' Sabine E. Van Doorslaer,t Freddy J. Callens,$ Filiep R. Maes and Paul F. Matthys Laboratory of Crystallography and Study of the Solid State, Krijgslaan 28lSl,8-9000Gent, Belgium Johan M. D'heer Laboratory for General and Inorganic Chemistry, Krijgslaan 28 1-S3,8-9000Gent, Belgium Two new triatomic molecular ions with C,, symmetry i.e.C0,-and SO,-have been detected in NaCl by means of EPR. By analysing the 13C hyperfine tensor of the C0,-radical the spin densities in the 2s and 2p, carbon atomic orbitals can be calculated. Since the 14N hyperfine tensor of the isoelectronic NO, molecule in the same host lattice is also known, a comparison between the spin densities in both defects can be made.In addition, the spin densities for both defects are calculated using a6 initio methods and are compared with the experimental results. As is the case for the smaller C0,-ion, the SO,-radical is assumed to substitute for only one halide ion. The monovacancy model for the SO,-ion is discussed. Such a model is in accordance with that for the smaller ozonide 0,-ion. On the other hand, the larger S3-and Se,- ions occupy a trivacancy site. Owing to their inherent high degree of symmetry, alkali- metal halide single crystals are very interesting host matrices for the study of atomic and molecular impurities. Once the species are trapped in the alkali-metal halide lattice, the elec- tronic structure and the orientation of the impurity can be studied with EPR and/or electron nuclear double resonance (ENDOR) to yield a wealth of information for both theo- reticians and experimentalists.This paper reports the EPR spectra of two related triatomic molecules, C0,- and SO,-, with C,,symmetry, trapped in NaCl. The C0,- ion has already been studied extensively in a wide variety of host lattices.'-* However, it has only very recently been detected in alkali-metal halide single crystals, specifically in KC19 and KBr." In these matrices, the defect exhibits axial symmetry around a (111) axis and hence is denoted as a C0,- { 11 l} species. It was assumed that the C0,-ion substitutes for a single halide ion. The spin den- sities in the carbon atomic orbitals of the CO, -radical could be obtained by analysing the 13C hyperfine tensor. A com-parison with the isoelectronic NO, molecule in the same host lattices and in the same configuration (both (111) species) could be made, yielding results in contradiction with theoreti- cal deductions made by Atkins et ul." The present experi- ments were carried out in order to check whether the C0,- ion can be incorporated in the NaCl lattice and also to deter- mine whether the comparison with the NO, molecule in NaCl yields results that contradict the theory.Theoretical calculations were performed to test the hypothesis of Atkins et ul." regarding the spin distribution in NO, and CO,-. Recently, there has been interest in the C0,- ion among EPR spectroscopists because it serves as a possible candidate for a low-dose (doses < 10 Gy) EPR do~imeter.~~*'~ Both the C0,- and the SO,-ions are triatomic molecules possessing C,, symmetry.In the alkali-metal halides, the SO,-ion has also already been detected in single crystals of KCl and KBr.14*15 The monovacancy model for the larger SO,-ion, proposed by these authors, will be tested using the data presented in this paper. As the 33S hyperfine tensor could not be determined, the spin density of the SO,-ion cannot be discussed. Research Assistant of the N.F.S.R. (Belgium).3 Senior Research Associate of the N.F.S.R. (Belgium). Experimental Materials NaCl single crystals were grown using the Bridgman tech- nique. The quartz growth capsule contained a graphite cylindro-conical crucible, filled with, typically 15 g NaCl (Merck Suprapur Powder).The NaCl powder was vacuum dried at 200 "C for one week using a two-stage rotation pump (ALCATEL 2012 AC, chemical series). Approximately 0.7 wt.% Na, l3Co3 was added with a I3C enrichment of 99% (MSD Isotopes, Miinchen, Germany) together with Na metal. The capsule was re-evacuated and sealed off. The samples thus grown were X-irradiated at room temperature for 30 min with a tungsten anticathode Philips X-ray tube, operated at 60 kV and 40 mA, corresponding to a dose of CQ. 40 kGy. Crystals were cut for rotation around a (1 10) axis. There was no intentional doping with a sulfur-containing species. The sulfur may originate from the carbon crucible where it is present as contamination.Methods The EPR spectra were recorded using a Bruker ESP300 X-band spectrometer. The maximum power of this spectrom- eter is 200 mW. The magnetic field was modulated at 100 kHz with a peak-to-peak amplitude of 0.5 x T. All spectra were normalised to the same frequency uiz. 9.47 GHz and hence can be compared directly. The magnetic field was measured using a Bruker NMR035M Gaussmeter. With this equipment it is possible to measure accurately the relative line positions of the EPR signals present. Small shifts of the magnetic field position down to 0.1 x lop4 T can be detected. Using an Oxford ESRlO flow cryostat, temperatures down to 4 K can readily be obtained. For absolute g value determination, a cali-bration using the g standard DPPH at 0.1 mW (g = 2.0036) was performed.Results No resonances could be detected in unirradiated crystals. After X-irradiation, resonances due to two different paramag- netic species were visible. The first defect (defect I) is visible J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 tk 3485e 3285' I <110> Cool>aldegrees Fig. 1 Angular variation in a (1 lo} plane of the resonances of the C0,-ion in NaCl. The height and the width of the rectangles is proportional to the line width and the signal height, respectively. The solid lines represent the theoretical angular variation, assuming zero line width. from 4 K to 180 K. The optimum detection temperature and microwave power for this defect are 4 K and 1 mW, respec- tively.The second defect (defect 11) is visible from 80 K up to 150 K and is preferentially measured at 100 K and 10 mW. Both defects will be discussed in more detail separately. Defect I The angular variation of the resonances of defect I in a { 110) plane is shown in Fig. 1. The observed spectra consist of three major groups of lines i.e. one group due to the unen- riched species (I = 0, "C) and two groups of lines (doublet) caused by the interaction of the electron spin with a 13C nucleus; this is the only nucleus with I = 1/2 in view of the doping procedure described above. The central line around g = 2 can be explained by considering the use of the graphite The central line has twice the intensity of the 13C hyperfine lines.The paramagnetic centre exhibits trigonal symmetry around a (1 11) axis. Owing to the large line width of the EPR resonances (compared with the g and A tensor anisotropy), the individ- ual resonances cannot be resolved in most orientations. Only when the applied magnetic field is oriented approximately along a (1 11) or a (1 10) direction, a small additional reson- ance is visible on the low-field hyperfine satellite (see Fig. 1). No additional substructure is detected on the central and high-field hyperfine lines. As was argued by Moens et al.," one must be very careful in determining the spin Hamiltonian parameters from such an angular variation by the 'classical methods' (ix.by fitting the resonance positions).As a result of the strong line overlap of the individual resonances, the observed line positions can be shifted from the real positions. Therefore, a program was developed for fitting the line profiles of all the observed spectra instead of only fitting the line positions of the EPR resonances. The resulting spin Hamiltonian parameters thus obtained are (hyperfine param- eters in MHz): g1 = 2.0022 (2), A, = 413 (3); gll = 1.9951 (3), All = 351 (3). The line width was optimized to be 11 (1) x T. The numbers in parentheses indicate the error on the last digit. The gI1axis is oriented parallel to a (111) axis. The theoretical angular variation, calculated with the above-mentioned spin Hamiltonian parameters is also shown in Fig.1. As can be seen, the small additional resonances of the low-field hyperfine line lie somewhat above the theoreti- cally predicted positions. That this is an effect of the large line 3350:: <loo> <110> aldegrees Fig. 2 Angular variation in a (100) plane of the resonances of the SO,-ion in NaCl. Same remarks as for Fig. 1. width and not of the fitting procedure, can be demonstrated by computer simulations. lo Defect I1 The angular variation of defect I1 exhibits orthorhombic symmetry and is shown in Fig. 2 in a (100) plane. The reson- ance positions can be adequately reproduced with the follow- ing g tensor: gx = 2.0017 (2), g,, = 2.0113 (2), gz = 2.006312). Here, the x and y directions are along the (110) and (110) axes, respectively, the z direction is along a (001) axis.As the line width is relatively small compared to the g tensor aniso- tropy, the individual resonances can be isolated in most orientations and hence the 'classical' optimization routines can be applied. The theoretical angular variation using the g values listed above, is also shown in Fig. 2. Discussion It will be demonstrated that both defects have to be ascribed to two different molecular ions. Hence defect I and I1 are discussed separately. Defect I As defect I exhibits hyperfine interaction with an I = 1/2 nucleus and as 13C is the only nucleus having nuclear spin I = 1/2, this defect must contain one 13C nucleus. Comparing the spin Hamiltonian parameters for defect I with data from the literature,'-'' the radical giving rise to the resonances of defect I is undoubtedly a C0,- ion.Table 1 lists the spin Hamiltonian parameters for the C0,- ion in KCl, KBr and NaCl. As can be seen, there is a strong resemblance between the C0,-ion in NaC1, KCl and KBr i.e. the spin Hamilto- nian parameters for the C0,- ion do not vary much with the host matrix. This can be readily understood by considering the electronic configuration of the radical. The CO, -molec-ular ion exhibits CZVsymmetry with a 2Al ground state.'v2 Thus, as the ground state is orbitally non-degenerate, the crystal field will merely cause a shift of the energy levels, having little effect on the g and A values. Table 1 Spin Hamiltonian parameters for the C0,-ion in several alkali-metal halide single crystals 91 II A, A,, AB/10-4T ref.~ ~~ KCl 2.0026 (1) 1.9962 (3) 388 (1) 328 (1) - 9 KBr NaCl 2.0026 (2) 2.0022 (3) 1.9948 (3) 1.9951 (3) 383 (3) 413 (3) 341 (3) 351 (3) 13 (1) 11 (1) 10 this work Hyperhe parameters in MHz. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Most probably, the C0,-radical substitutes for one halogen ion in the alkali-metal halide host lattice. This model is depicted in Fig. 3. As a result of the 'A, ground state of the C0,- ion, the one-electron molecular orbital containing the unpaired elec- tron can be written as: Here, the z axis is oriented along the main symmetry axis of the molecule (i.e. the C2 axis).The y axis is parallel to the direction connecting the two oxygen atoms, whereas the x direction is perpendicular to the molecular plane. The super- script in the notation for the atomic orbitals denotes the atom on which the orbital is centred. Normally, the C0,-species exhibits orthorhombic g and A Typical values for g and A, measured along the distinct molecular axes, are: gx = 2.0030, g, = 1.9970, gz = 2.0015 and A, = 513 MHz, A,= 506 MHz, A, = 630 MHz. The axial g value found for defect I indicates that the C0,-radical is rapidly rotating around the y axis (i.e. the axis connecting two oxygens). That such a rapid tumbling can occur at such low temperatures, has already been established for the isoelec- tronic NO, m01ecule.'~*'~ From the principal values of the 13C hyperfine tensor of the C0,-ion, one can easily calculate the spin densities in the 2s and 2p, carbon atomic orbitals [see eqn.(l)]. The spin density in the 2s carbon atomic orbital can be estimated by comparing Ai, with 2p0 g, gN 1,pN IYzs(o)12/3. The latter being calculated to be 3108.66 MHz.18 On the other hand, the spin density in the 2p, carbon atomic orbital is estimated by comparing A, -with 0.29, gN 1,pN(r2p-')//I, the latter being calculated as 45.36 MHz.'* Because the "O hyperfine tensor is unknown, only the spin density in all the oxygen atomic orbitals can be deduced from the normalisation to unity of the total spin density. The results are given in Table 2. As the 14N hyperfine tensor of the isoelectronic NO, mol- ecule in the same configuration and in the same host lattices is also known (i.e.an NO,{lll) rnolec~le),'~ we are in an excellent position to compare the spin densities of both the C0,-and NO, ions in KCl, KBr and NaCl. The results are also shown in Table 2. <001> Fig. 3 Model for the C02-ion in NaCl. The different ions are drawn to scale, taking into account their ionic radii. 3263 Table 2 Experimental spin densities on the individual carbon and nitrogen orbitals: S&N) and P&%N) for co2-and NO, in KC1, KBr and NaC1;" ck is the spin density on the oxygen atoms, derived from the normalisation to unity of the total spin density co,-NO, ','C CL 'i CZN ';,N '6 KCl 0.118 (1) 0.43 (2) 0.45 (2) 0.099 (2) 0.47 (2) 0.43 (2) KBr 0.119 (2) 0.32 (4) 0.56 (4) 0.099 (2) 0.40 (4) 0.50 (4) NaCl 0.127 (2) 0.45 (4) 0.42 (4) 0.102 (1) 0.49 (3) 0.41 (3) The values in parentheses denote the error on the last digit.From these results, it can be noticed that the spin density on the oxygen atoms is always larger for the C0,-ion than for the NO, molecule. These results are at variance with the theoretical deductions about C0,-and NO, ,''9l9 that predict that the spin density on the oxygen atoms is larger for NO, than for CO,-. In order to check this contradiction between theory and experiment, theoretical calculations of the spin densities of both free molecular ions using the GAMESS9O2O program were performed. The program was run on an IBM RS/6000, model 320 workstation.The spin densities of C0,-and NO, were calculated using five different basis sets. These were thus chosen to include minimal (ST0-3G),,' split valence (6-31G,,' 6-31 lG2') and polarized (6-31G*, 6-311G*)24 basis sets. The geometry of the molecules was optimized for each basis set. The spin densities calculated using the Lowdin spin population analysis are summarized in Table 3. For the oxygen atoms, only the total spin density is given. As can be seen from Table 3, the spin density on the oxygen atoms is always larger for NO, than for CO,-, in agreement with the theoretical deductions made by Atkins et al.' The discrepancy with the experimental results most '9" probably has two origins. First, the theoretical calculations are performed on free ions whereas the experiments are carried out on molecules embedded in a crystal matrix and secondly, the calculations do not account for the rapid rota- tion of the C0,- ion around the oxygen-oxygen axis.Such a rotation introduces a spin-rotational coupling term in the Hamiltonian which may have an effect on the magnetic res- onance spectrum. This was considered in some detail by Bojko and Silsbee16 for the NO, molecule. Ab initio calcu-lations, performed on radicals embedded in a host matrix can possibly give an indication of the importance of the effect of the crystal field. Calculations incorporating the effect of the crystal field through the Madelung potential will be carried out in our laboratory in the near future.Table 3 Comparison of the theoretical spin densities in the differ- ent atomic orbitals of carbon and nitrogen in C0,- and NO,, and the total spin density on the oxygen atoms ion basis 1s 2s 2p, 2py 2p, 0 C0,-STO-3G 0.0018 0.1301 0.0323 -0.0347 0.4060 0.4645 6-31G 0.0025 0.2205 0.0439 -0.0267 0.4394 0.3204 6-311G 0.0017 0.2298 0.0428 -0.0271 0.4294 0.3234 6-31*G 0.0015 0.1739 0.0347 -0.0242 0.4206 0.3171 6-311*G 0.0017 0.1806 0.0341 -0.0248 0.4148 0.3116 NO, STO-3G O.ooo6 0.0558 0.0102 -0.0450 0.3188 0.6597 6-31G 0.0013 0.1093 0.0389 -0.0274 0.3930 0.4845 6-311G 0.0010 0.1085 0.0406 -0.0266 0.3947 0.4819 6-31*G 0.0011 0.0846 0.0293 -0.0232 0.3718 0.4832 6-311*G 0.0011 0.0861 0.0315 -0.0258 0.3792 0.4826 Data obtained using the Lowdin population analysis calculated with the GAMESS90 program package.Defect I1 The paramagnetic radical giving rise to the resonances of defect I1 exhibits no hyperfine interaction. Hence it cannot contain a carbon nucleus. In view of the doping procedure, the most plausible candidate for this radical would be the 0,-ion. However, the ozonide ion has to be rejected on the basis of the g values, as can be seen from Table 4.However, comparing the g values of defect I1 with those of the SO2-ion in KCI and KBrl49I5 indicates that this paramagnetic species must be identified as an SO,-molecular ion (see Table 4). The sulfur most probably originates from the carbon crucible where it is present as contamination.Reac- tion of the sulfur with the dopant during the growth process can result in the formation of SO2and SO, molecules. The SO2-ion exhibits CZy symmetry and has a ,B, ground ~tate.'~ Hence, within the LCAO scheme, the unpaired electron resides in a molecular orbital of the form: where the same notations and conventions for the molecular axes as for eqn. (1) are used. As can be seen from eqn. (2), the electron resides mainly in the px atomic orbitals i.e. the orbitals perpendicular to the molecular plane. Unfortunately no conclusions can be drawn about the spin densities in the distinct atomic orbitals as no data are available for either the 33Sor the 170 hyperfine tensors. The SO,-ion most probably substitutes for one halide ion. The model in depicted in Fig.4.The incorporation of the SOz-ion will be governed by the electrostatic interactions between the paramagnetic px lobes of the defect (as these contain the unpaired electron) and the surrounding Na' and C1- ions of the host matrix. In the case of a monovacancy model, the electrostatic interaction will be mainly between the px lobes and the six nearest Na' ions. As the p, lobes are oriented perpendicular to the molecular plane, the only Na' ions which are of interest here are the two Na+ ions situated Table 4 g tensor values for the 0,-and SO,-ions in several alkali-metal halide single crystals lattice defect 9x gv 9, ref. KCI 03 -2.0032 2.0182 2.0118 22 K Br 03 -2.0027 2.0180 2.0113 23 KI 03 -2.0030 2.0185 2.0116 24 KCI SO,-2.0025 2.0110 2.0071 14 KBr SO,-2.0050 2.0100 2.0075 14 NaCl defect I1 2.0017 2.0113 2.0063 this work <110> i <OOl> Fig.4 Model for the SO,-ion in NaCI. Same remarks as for Fig. 3' J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 above and the two situated below the molecular plane. The two Na ions lying above the molecular plane are indicated + in Fig. 4 by 1 and 2. Thus, it can be seen that the SO,-ion has to be placed slightly off-centre, so that the px lobes avoid the electron clouds of the Na+(l) and Na'(2) ions as much as possible. However, the total C,, symmetry is preserved. The situation is in some respects analogous to the case of the smaller 0,-ion in KCl, KBr and KI.25-28 B y assuming that the incorporation of this ion is also governed by electro- static interactions, it was predicted that the ozonide ion could not be incorporated untilted in the KC1 lattice because of the too small dimensions of the C1- vacancy.On the other hand, the Br -and I -vacancies offer enough space for the 0,-ion to be untilted. All these aspects were indeed observed with EPR, giving experimental evidence for the proposed model for 0,-to be extended in this paper to SO,-. By the same arguments the larger S,-and Se,- ions should occupy a tri- vacancy site, as was confirmed e~perirnentally.'~*' Conclusions The EPR spectra of two small triatomic molecules with C,, symmetry in NaCl uiz. C0,-and SO,-, are presented. The C0,-ion rotates rapidly around the axis connecting the two outer oxygen atoms, even at liquid-helium tem- peratures. The same peculiar feature was observed for the iso- electronic NO, molecule. By analysing the 13C hyperfine tensor, the spin densities in the 2s and 2p, carbon atomic orbitals could be calculated.In addition, the total spin density on the oxygen atoms could be deduced from the nor- malisation to unity of the total spin density. When these data are compared with those for the NO, molecule, results in contradiction with theoretical calculations using the GAMESS90 program are obtained. Most probably, this dis- crepancy between theory and experiment must be explained by the fact that the calculations were performed on free rad- icals, whereas the experiments concern radicals embedded in a crystal matrix.The monovacancy model for the SOz-ion, as proposed by other authors, is confirmed by our measurements. The authors wish to thank the Executieve van de Vlaamse Gemeenschap-Departement Onderwijs and the Inter-universitair Instituut voor Kern Wetenschappen (IIKW) for financial support. References 1 D. W. Ovenall and D. H. Whiffen, Mol. Phys., 1961,4, 135. 2 S. A. Marshall, A. R. Reinberg, R. A. Serway and J. A. Hodges, Mol. Phys., 1964, 8, 225. 3 R. J. Cook and D. H. Whiffen, J. Phys. Chem., 1967,71,93. 4 S. Schlick, B. L. Silver and 2. Luz, J. Chem. Phys., 1971,54,867. 5 P. Meriaudeau, J. C. Vedrine, Y.Ben Taarit and C. Naccache, J. Chem. SOC.,Faraday Trans. 2,1974,71,736. 6 G.Bacquet, V. Quang Truong, M. Vignoles, J. C. Trombe and G. Bonel, Calcif Tissue Znt., 1981,33, 105. 7 F. J. Callens, R. M. H. Verbeeck, P. F. A. Matthys, L. C. Martens and E. R. Boesman, Calcif. Tissue Int., 1987,41, 124. 8 R. Debuyst, M. Bidiamambu and F. Dejehet, Bull. SOC. Chim. Belg., 1990,99, 535. 9 F. Callens, P. Matthys and E. Boesman, J. Phys. Chem. Solids, 1989,50,377. 10 P. D. W. Moens, P. F. A. Matthys, F. J. Callens, F. R. C. Maes and E. R. Boesman, Phys. Status. Solidi,B, 1991,168,289. 11 P. W. Atkins, N. Keen and M. C. R. Symons, J. Chem. Soc., 1962, 3, 1873. 12 P. N. Keizer, J. R. Morton and K. F. Preston, J. Chem. SOC., Faraday Trans., 199 1,87, 3 147. 13 R. Gray and M. H. Stevenson, Bruker Report 91192,1992, p. 9. 14 J.Schneider, B. Dischler and A. Rauber, Phys. Status. Solidi, 1966, 13, 141. J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 3265 15 16 17 18 J. Suwalski and H. Seidel, Phys. Status. Solidi, 1966,13, 159. I. Bojko and R.H. Silsbee, J. Magn. Reson., 1971,5, 339. J. R. Brailsford and J. R. Morton, J. Magn. Reson., 1969,1,575. A. H.Koh and D. J. Miller, At. Data Nucl. Data Tables, 1985, 23 24 R. Krishnan, J. S. Binkley, R. Seeger and J. A. Pople, J. Chem. Phys., 1980,72,650. P. C. Chariharan and J. A. Pople, Theoret. Chim. Acta, 1973,28, 213. 19 20 21 33,235. P. W. Atkins and M. C. R. Symons, The Structure of Inorganic Radicals. An Application of Electron Spin Resonance to the Study of Molecular Structure, Elsevier, Amsterdam, 1967. M.W. Schmidt, K. K. Baldridge, J. A. Boatz, J. H.Jensen, S. Koseki, M.S. Gordon, K. A. Nguyen, T. L. Windus and S. T. Elbert, QCPE Bulletin, 1990,10,52.W.J. Hehre, R.F. Stewart and J. A. Pople, J. Chem.Phys., 1970, 25 26 27 28 F. Callens, P. Matthys and E. Boesman, J. Phys. C, 1989, 21, 3159. F. Maes, F. Callens, P. Matthys and E. Boesman, Phys. Status. Solidi, B, 1990,161, K1. P. Moens, F. Callens, S. Van Doorslaer, F. Maes and P. Matthys, Phys. Status. Solidi, B, 1994, 182, 21 1. F. Callens, P. Moens, S. Van Doorslaer, F. Maes and E. Boesman, J. Chem. Soc., Faraday Trans., 1993, W,3691. 22 52,2769. W. J. Hehre, R. DitcMeld and J. A. Pople, J. Chern. Phys., 1972, 56,2257. Paper 4/03002H;Received 20th May, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003261

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3267-3271 Time-resolved EPR Study of Spin Polarization and Dynamics of F+ Centres in Additively Coloured CaO Crystalst Carlo Corvaja, Lorenzo Franco, Luigi Pasimeni and Antonio Toffoletti Department of Physical Chemistry, University of Padova, Via Loredan 2,35131 Padova, Italy Optical excitation of additively coloured CaO crystals generates 'T, and 3T, excited states and F+ defects, the latter being revealed by EPR experiments. Interaction of the excited 3T,u triplet state and the F+ centre is documented by the emissive character of the transient EPR signal detected just after a laser pulse. A quenching of the initial polarization, observed at high microwave power, is ascribed to microwave saturation of the F+ spin leveIs.In additively coloured CaO crystals the defect present at the largest concentration is the F colour centre consisting of an oxygen ion vacancy at which a pair of electrons are trapped with singlet ground state. Besides F centres other defects have been identified in CaO crystals, among them the Ff centre formed by a single electron trapped at the oxygen ion vacancy.' Neighbouring F+ centres can associate giving rise to aggregates such as dimers F; + (a nearest-neighbour oxygen divacancy containing two electrons) and trimers Fi + (a self-trapped hole, adjacent to an F;' centre).' The F+ centre is paramagnetic (S = 1/2) and Henderson and Wertz' reported that its continuous wave (CW) EPR spectrum consists of a single narrow line (linewidth <20 pT), being F+ free of hyperfine splitting because of the low (0.14%) natural abundance of the 43Ca nuclide with I = 7/2.We have observed that irradiation of the crystal by UV or visible light causes a marked enhancement of the EPR line intensity. Moreover, in some conditions, by using pulsed irra- diation, the signal intensity just after the light pulse corre- sponds to microwave emission. It has been shown that when CaO crystals are excited with light within the F band (F absorption band), photoionization of the F centres takes place, with F -+ Ff centre conversion. The most detailed study of photoionization at the F centre in CaO was reported by Welch et d..' They also showed that in CaO the F absorption band, which peaks at 400 nm, arises from the allowed transition from the 'A, ground state to the 'Tlu state, rapidly converted into the 9T,, state by inter- system crossing (ISC).At temperatures above 55 K electrons excited to the 'T,, and 'TI,, levels of the F centre are involved in a thermal ionization process which raises them to the conduction band. The activation energy for the excitation from the 'Tlu is 0.1 eV. Electrons excited into the conduction band are then trapped at shallow states below the bottom of the conduction band, from whence they are subsequently slowly released to form the ground-state F centre. The enhanced CW EPR signal intensity we observed under optical excitation could be due to the increased concentration of the doublet species F generated by irradiation.+ Recently some papers have appeared in the literat~re~.~ in which a transient change in intensity of the EPR spectrum of an S = 1/2 species under illumination was observed. This effect is explained in terms of a tripletaoublet (TD) inter- action. The latter causes a change in the relative spin sublevel population of the doublet species and consequently of the EPR signal intensity. Our observation on CaO might also be explained by this mechanism. The occurrence of an inter- action between triplet and doublet states in CaO crystals was t This paper was presented at the 27th International ESR Con-ference at the University of Wales, Cardiff, 21st-25th March, 1994. suggested by Glasbeek and Hond' in order to explain the temperature behaviour of the relaxation properties of the Fi+ defect in the excited triplet state in the presence of the F+ doublet species.In order to distinguish the possible causes of EPR signal intensity variation and to have further insight into the TD interaction, we have performed two types of time-resolved EPR experiments. In the first the excitation light was pulsed, which instantaneously generates the excited species and the time-evolution of the CW EPR spectrum was recorded.6 In the second, a pulsed EPR experiment was performed7 in the dark as well as during continuous illumination. The triplet- doublet perturbation switched on by the triplet species gener- ation, was expected to produce initial spin polarization.It was possible to study the spin dynamics of the F+ centres by comparing their spin relaxation properties with and without illumination. Experimenta1 Additively coloured crystals of CaO were purchased from Spicer Ltd. CaO is face-centred cubic and the unit cell (a = 4.80 A) contains four molecules of CaO.* The crystals were orientated by the X-ray technique and attached at the bottom of a Plexiglas rod which allowed us to rotate them around an axis perpendicular to the static magnetic field, B,. The rod was inserted into a quartz tube to avoid any exposure of the crystal to humidity. EPR measure- ments were performed with the magnetic field exploring the [OlO] plane. CW and pulsed EPR spectra were recorded by using an EPR Bruker ESP 380 X-band spectrometer with a dielectric resonator.The crystal was irradiated by a xenon lamp (Cermax 300 W). Time-resolved (TR) EPR measurements were carried out with a Bruker ER 200 D EPR X-band spec- trometer equipped with a standard TElo2 rectangular cavity and nitrogen flow cryostat. Pulsed light was supplied by an excimer XeCl laser (A = 308 nm, pulse duration = 20 ns) (Lambda Physik LPX 100). The EPR signal was directly recorded without field modulation and lock-in detection. The signal was collected and averaged by a digital oscilloscope (Nicolet 4094C, maximum resolution 5 ns per point). An averaged off-resonance signal was subtracted from that on-resonance to avoid noise coming from the laser trigger. Results The room-temperature CW EPR spectrum of the CaO crystal recorded in the dark is shown in Fig.1. The spectrum consists of the superposition of two lines with different line- widths and saturation properties. The measured linewidth Fig. 1 CW EPR spectrum of the F+ defects in additively coloured CaO crystals recorded in the dark at room temperature. Experimen- tal detection conditions are: mw power 0.54 pW,field modulation frequency 6.25 KHz, field modulation strength 0.6 pT. ABPpof the narrow component is 1.3 pT which is at the limit of field homogeneity of the EPR spectrometer, superimposed on a broader component having linewidth of 4.0 pT. This extremely narrow EPR line was recorded by using a low modulation frequency (6.5 kHz) in order to avoid line broadening due to sidebands of the modulation frequency.When the crystal is illuminated, the two components merge in a broader single line (ABpp= 6.5 pT) and, most inter- estingly, the signal intensity is enhanced considerably. After the light is turned off the intensity of the EPR line regains its previous value in times of the order of several tens of minutes (Fig. 2). Spin-lattice TI and spin-spin T2 relaxation times were measured from pulsed EPR experiments, performed in the dark as well as under continuous illumination. The decay time T; was obtained from the free-induction decay (FID) signal, which monitors the time-evolution of the y component of magnetization, M,(t), after a n/2 microwave pulse. TT is defined as = 1/T2h + 1/T2inh (1) where T2h and TZinh are the homogenous and inhomogeneous contributions.At room temperature the values of 7’;mea-sured in the dark and under continuous illumination are 4.8 f0.2 and 1.1 f0.1 ps, respectively (Fig. 3). The homogeneous contribution T2h has been measured from the decay of the spin-echo intensity in a Carr-Purcell- Gill-Meiboom sequence of microwave pulses J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0 10 20 0 1 2 t/P frequency/M Hz Fig. 3 Experimental FID (left) and its fast Fourier transform (right) of F+ centres in CaO recorded at 290 K (a)in the dark and (b)under continuous illumination of the sample. The FID consists of a damped oscillation with damping parameter T:. The frequency of oscillation represents the offset of the magnetic field from the resonance. 7r/2-(-z-n-t-echo-), .To obtain the correct TZhvalue the suitable phase cycle was used in order to compensate for the incorrect length of the pulses. At room temperature in the dark the measured value of T2h is 30 k2 ps, which would correspond to a linewidth of only 0.19 pT in the CW spec-trum. No spin-echo was detected under continuous illumi- nation. The TI values were measured by the inversion recovery technique applied to the FID signal, consisting of the sequence of a n pulse followed after a delay time, t, by a n/2 pulse.’ After the first microwave pulse, magnetization of the spin system reverts to its initial value and evolves in time to recover its equilibrium value with the characteristic time TI.We measured the peak intensity of the spectra obtained by Fourier transform of the FID as a function of the delay, t, between the pulses (Fig. 4) The value of TI, measured in the dark at room temperature, is 48 & 4 ps. In the TR-EPR experiments’signals detected with the mag- netic field, B,, set in resonance with the EPR line, exhibit a marked dependence on the microwave (mw) power. At low 1 MHz 0 100 200 T time/min Fig. 4 Stacked plot of a series of FT EPR spectra of F+ centres in Fig. 2 Decay of CW EPR signal intensity of F+ centres in CaO, CaO recorded at different delays, T, in a n-z-n/2 pulse sequence. after illumination with visible light. Experimental points are fitted Typical pulse lengths are 16 ns (n/2 pulse) and 32 ns (n pulse).The with the function I = A/(B+ t)where A = 9 min, B = 11 min. real part of the transform, after a linear phase correction, is shown. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 ~ ~~~ emission 6 12 18 ti me/ms Fig. 5 TR-EPR signal of the F+ defects in the CaO crystal recorded at 250 K and 50 pW of mw power mw power, particularly, the signal, initially in emission, con- verts into absorption and eventually disappears with a single- exponential decay function as shown in Fig. 5. On increasing the mw power the emission reduces and the initial growth of the absorption becomes faster. At sufficiently high mw power the signal displays characteristic damped oscillationsY6 whose frequency depends on the strength of the oscillating mw field.A selection of experimental transient signals is shown in Fig. 6, together with computer simulations which will be discussed later. Discussion Measurements of T,and T,by Pulsed EPR In the dark the Tz value measured from the FID signal at room temperature is 4.8 ps corresponding to a linewidth AB,, = 1.2 pT which compares nicely with the value of 1.3 pT measured for one component of the CW spectrum. The dephasing-time, TZh,obtained from the decay of the spin- echo intensity by applying the Carr-Purcell-Gill-Meiboom sequence of microwave pulses is 30 ps. This time is extremely long and corresponds to a homogeneous linewidth of 0.19 pT. (4 20 mW 20 mW 2 mW 0.2mW 0.2 mWII7 0 18 ti me/ps Fig.6 (a) Experimental and (b) simulated TR-EPR signals of F+ defects in a CaO crystal recorded after the UV laser pulse at 290 K with different values of mw power. Fitting parameters are: TI = 4 p, T, = 0.8 ps. The above results are explained by allowing that in the dark the CaO crystals may host two kinds of F+ defects: (i) isolated F+ centres; (ii) weakly interacting F centres, giving + rise to the narrow and the broad components in the CW spectrum, respectively. The observed FID and spin-echo signals are assigned to the former species whose narrow com- ponent in the CW spectrum is considered inhomogeneously broadened by the inhomogeneity of the static magnetic field. The broad line arises from the interacting F+ defects with a homogeneous linewidth.Its FID signal decays rapidly, obscured by the slower decaying FID signal of the isolated species and no spin-echo signal is expected. When the light is turned on, new Ff species are generated from the F centres and the amount of interacting F+ defects increases at the expense of the isolated ones. The CW spec- trum then shows a single broad line, the FID signal decays with a shorter time TZ and no spin-echo is detected due to the homogeneous linewidth of the interacting species. All these features have been observed experimentally. EPR Experiments with Pulsed Light Two time regions could be clearly identified in the time- evolution of the EPR signal shown in Fig. 3. In the first, the EPR signal, initially in emission, converts into an absorption signal whose maximum intensity is achieved within a few hundred ps.After that time the variation is much slower and the signal decays exponentially with the characteristic time of 5.4 ms. This decay time is much longer than the spin-lattice relaxation time TI, as measured by the pulsed EPR experi- ments. At times longer than a few hundred ps any polariza- tion effect is over and the subsequent decay of the signal has to be ascribed merely to a process which depletes the species responsible for the EPR absorption. On the other hand the EPR absorption recorded after continuous illumination shows a much slower signal decay (see Fig. 2). The results of the two experiments indicate that there are two independent decay channels. The studies of luminescence and photoconductivity carried out on CaO crystals' have shown that the light absorption by F centres is followed by thermal photoionization from the excited ITlu and 3Tlu states, with the production of F+ centres and electrons in the conduction band.The photoion- ization process causes a decrease of the excited-state lumines- cence lifetime which drops from ca. 3 ms at 4.2 K to a fraction of an ms at temperatures higher than 70 K. The con- duction electrons are eventually trapped by relatively shallow impurity levels and are successively released to combine again with the F+ centres. The EPR signal decay should reflect the recombination process. The fact that the EPR signal decays with two very different characteristic times (of the order of lo3 and s) suggests the presence of two types of electron traps having different energy depths.Concerning the short time range, we note that the mecha- nism of photoionization does not contain any spin-selective process capable of generating spin polarization. Therefore, the contribution to the EPR signal due to this mechanism is initially zero. To account for the observed initial emissive polarization one is forced to look for a different process. Very recently, we reported the transient EPR emission spectra of a free radical trapped in a single crystal of chlora- nil.' When the crystal was illuminated inside the cavity of an EPR spectrometer, the intensity of the free radical spectral lines changed.After a laser pulse, the transient EPR signal was initially in emission. This behaviour was accounted for by assuming that triplet species, produced by the light excita- tion, interact with the stable radical species, according to the following kinetic scheme. T+ D The same model could explain the initial spin polarization observed in the transient EPR spectra of CaO F+ centres. Accordingly, the excited triplet states 3Tlu generated in the crystal, interact with Ff defects in the ground state giving rise to triplet-doublet pairs. If the spin-exchange interaction between the partners of the pair is non-zero the energy levels of the quartet and doublet spin states of the pair are split. The pair doublet state deactivates by conserving its spin multiplicity and giving rise to triplet quenching, while the quartet state is unreactive.However, before it dissociates back into a triplet and a doublet, it acquires a small admix- ture with the doublet components. The doublet-quartet mixing is caused by the zero-field splitting (zfs) interaction. Such a mixing is spin selective and if J < 0 (the doublet state lies lower in energy than the quartet), one finds that the rate constant of the process populating the 1/2 spin component of the excited doublet species D* emerg- ing from the doublet state of the pair is larger than that of the -1/2 component. The D* levels decay rapidly into those of the ground-state D species by a spin-allowed transition.Hence, the doublet species D participating in the pair forma- tion acquires emissive spin polarization. Another source of doublet-quartet mixing is the hyperfine interaction, which, however, is completely absent in CaO F+ centres. It should be noted that in our crystals the initial spin polarization does not involve all the F+ centres because those formed by photoionization are not polarized. Dependence ofthe TR-EPRSignal on mw Power Fig. 5 shows the EPR signal evolution in the short time range. The time dependence of the EPR signal is affected by the mw power. On increasing the mw power: (i) as expected, the signal displays transient nutations due to the spin dynamics;" (ii) the initial spin polarization diminishes until it disappears.Indicating by o1= yB,, the Rabi frequency proportional to the strength, B,, of the oscillating mw field, by following Atkins et al." one derives, at resonance, the following rela- tion for My@)which is proportional to the EPR signal where and = &[(T,' + Ti1) T [(T,' -Ti')' -4~:]~/'] (4) where T, and T, are the spin-lattice and the spin-spin relax-ation times, respectively, while M,, and M,(O) = 2, M,, are the magnetization along z at thermal equilibrium and just after the laser pulse. At high mw power eqn. (2) reduces to the form My(t)= (al/b)exp(at)sin(bt) (5) where the parameters a and b are given by: J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 At moderate mw powers the signal decays in a single-exponential manner with a time constant T;" given by l/T"," = l/Tl + a:T2 (7) The measured values of Ttff extrapolated to zero mw yield Tl-In order to simulate the whole series of EPR signals, recorded at different mw powers, one has also to take into account how the mw power affects the initial spin polariza- tion.This decreases as the mw power is increased. In order to account for this behaviour, it is worth noting that there should be a mechanism acting on the F+ centres before the laser pulse. In fact the initial polarization is generated in a time much shorter than the reciprocal of the Rabi frequency, o;l, which is the time the mw field needs in order to produce an appreciable change in the spin sublevel population. Moreover, the EPR line of the F+ centres is very narrow, as we have shown, and it can be easily saturated. Thus, one could infer that saturation of the spin levels of the F+ doub-lets existent before the laser pulse, is responsible for the reduced initial polarization.Let us consider the spin levels of the F+ centre, those of the T-D pair in the doublet state ('[T, D]) and the two spin sublevels 4[T, D]*33/2 of the doublet-triplet pair; the latter would be in a pure quartet state in the absence of doublet- quartet mixing. They are represented schematically in Fig. 7. When an F+ centre in the spin state -1/2 forms a pair with an F centre in the triplet excited state having -1 spin component, the [T, D] pair produced is in the 4D,D]-3/2 quartet state. Because of the mixing caused by the zfs inter- action within the F-centre triplet state, the quartet substate is mixed with the + 1/2 component of the pair doublet state as indicated by the double arrows in Fig.7. Therefore the pair could evolve, conserving the total spin, into an F+ centre in the excited doublet state (spin component +1/2) and a ground-state F centre. The successive deactivation of the excited F+ centre occurs at the + 1/2 sublevel. An analogous process takes place from the + 1/2 level to the -1/2 level. The overall process due to the doublet-triplet interaction can be envisaged as a population transfer from 'D+,/, to 'Ddl/, and vice versa. By indicating with k+-and k-+ the 8 0. D" 1/2 -112 Fig. 7 Energy levels of the F+ doublet state (either in the ground state or in the excited state, D*) and of the triplet-doublet interacting pair.Continuous straight arrows indicate the population transfer path from F-to F+ through interaction with a triplet-state F centre in the I -1) spin substate (T-J. Dashed arrows show the path from F+l,2+to F- through interaction with a triplet-state F centre in the I + 1) spin substate (TI).Double arrows symbolize the mixing between quartet and doublet states of the interacting pair (the larger the arrows the greater the mixing). J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 rate constants of the two processes, the net spin polarization, P, transferred is given by P = k+-n+l12-k-+ n-l12 where n,,/, are the populations of the corresponding spin levels.When the exchange interaction J is large and negative (in the pair the quartet has higher energy than the doublet) so that the only effective mixing occurs between 4[T, D]-3/2 and ’[T, D]+ 112, the rate constant k-+ % k+ -and a net emissive polarization is delivered to the F+ species inter- acting with the 3Tlutriplets. However, if J is fairly small, the doublet-quartet leads to the result k, -x k-+ . In these conditions the transfer of polarization, P,relies merely on the difference (n+1/2 -n-lj2) that can be depleted by saturating the EPR transition by a strong mw field, before the interaction with the triplet species. In order to compute the transient signals we assumed an initial polarization proportional to the saturation factor : The simulated signals, shown in Fig.6, fit rather satisfacto- rily the experimental ones, although the fitting was achieved by considering a single set of relaxation times, neglecting the presence of two species and, furthermore, the effect due to the chemical decay was not taken into account. The above observations indicate that in [T, D] pairs in CaO the exchange interaction is rather small. This is not sur- prising because the excited triplet species 3Tlu in CaO crys- tals are generated at the localized F sites lying rather far from the F+ defects, and they do not diffuse. Conclusions We have shown that the generation of triplet states by illumi- nation with UV-visible light in the additively coloured CaO crystal produces new excited F+ besides the stable F+ species in the ground state, as revealed by the enhanced intensity of the CW EPR spectrum.We have also shown that F centres excited in a triplet state give rise to interacting tripletdoublet pairs responsible for the initial spin polarization observed in the TR-EPR signal after the laser pulse. On increasing the mw power, this polar- ization is quenched. The saturation of the EPR transition, occurring between the levels of F’, is responsible for the quenching of the initial spin polarization. This work was supported by the Italian National Research Council (CNR) through the Centro Studi sugli Molecolari Radicalici ed Eccitati and by the Minister0 dell’universita e della Ricerca Scientifica e Tecnologica (MURST). References 1 B. Henderson and I. E. Wertz, Adv. Phys., 1968,17,749. 2 L. S. Welch, A. E. Hughes and G. P. Summers, J. Phys. C, 1980, 13, 5801. 3 C. Blattler, F. Jent and H. Paul, Chem. Phys. Lett., 1990, 166, 375. 4 A. Kawai, T. Okutsu and K. Obi, J. Phys. Chem., 1991,%, 9130. 5 M. Glasbeek and R. Hond, Phys. Rev. B, 1981,23,4220. 6 R. Furrer, F. Fujara, C. Lange, D. Stehlik, H. M. Vieth and W. Vollmann, Chem. Phys. Lett., 1980, 75, 332. 7 A. Schweiger, Angew. Chem. Znt. Ed. Engl., 1991,30,265. 8 C. W. Bunn, Chemical Crystallography, Oxford Univerisity Press, London, 1961. 9 C. Corvaja, L. Franco, L. Pasimeni, A. Toffoletti and L. Monta-nari, Chem. Phys. Lett., 1993,210,355. 10 (a)P. J. Hore and K. A. McLauchlan, J. Magn. Reson. 1979, 36, 129; (b)P. J. Hore and K. A. McLauchlan, Mof.Phys., 1981,42, 533. 11 P. W. Atkins, K. A. McLauchlan and W. Percival, Mof.Phys., 1973, 25,281. Paper 4/02242D; Received 14th April, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003267

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3273-3280 EPR, ENDOR and TRIPLE Resonance Characterization of Three Paramagnetic Redox Stages of 5-Methylene-5H-dibenzo [a,d]cycloheptene M. Luisa T. M. B. Franco and M. Celina R. L. R. Lazana lnstituto Superior Tecnico, Laboratorio de Quimica Orgdnica Av. Rovisco Pais, P-1096 Lisboa Codex, Portugal Radical anions and radical trianions derived from 5-methylene-5H-dibenzo[a,d]cycloheptene 1, lO-deuterio-5-met h y Iene-5H-dibenzo[a,d]c yc loheptene 1 d and 5-d ideu te r iomet h y Iene-5H-dibenzo[a ,d]cyclo he ptene 1d, by reduction with lithium, sodium, potassium and caesium in ethereal solvents have been studied by EPR, ENDOR and TRIPLE resonance spectroscopy. The spin and charge distribution in 1'-and la3-are discussed both in terms of HMO-McLachlan and INDO calculations.The energies of the lowest unoccupied molecular orbital (LUMO) and of the next lowest unoccupied molecular orbital (NLUMO) are influenced differently by the strong interaction of the supercharged radical trianion with the counter-cations. A change in the orbital sequence of LUMO and NLUMO was predicted theoretically for the radical trianion relative to that found for the radical anion, in accordance with the experimental results. The radical cations lo+,Id'+ and Id," obtained via oxidation by AICI, in dichloromethane solution were also studied by EPR, ENDOR and TRIPLE resonance spectroscopy. The experimental results were also interpreted in terms of HMO-McLachlan and INDO methods.Although numerous reports on radical anions and dianions of polycyclic conjugated hydrocarbons have been published, only very limited data on higher negatively charged species have appeared in the literature. Since the earlier reports on the EPR of radical trianions were p~blished,'-~ the interest in these species has increased. In the review published by Gerson and Huber4 only 16 compounds were then referred to as giving radical trianions which had been characterized by ~1.~9~EPR. Since then, Hirayama et have examined the radical trianions of 2,2'-(4,9-dihydronaphtho[2,3-c][1,2,5] thiadiazole-4,9-diylidene)bis(propanedinitrile) and 2,2'-(4H, 8H-benzoC 1,2-c :4,5-c']bis[ 1,2,5]thiadiazole-4,8-diylidene)bis (propanedinitrile). Gerson and co-worker~~.~ described the radical trianions of 1,3,5,7-tetra(tert-butyl)dicyclopenta[u,e] pentalene7 and phenyl-substituted 1,2:9,10-dibenzo[2.2]-paracyclophane- 1,9-dienes'.The radical trianion of 4,7-phe- nanthroline was reported by Fujita and Ohya-Nishiguchig and the radical trianions of thio-and dithio-esters of benzenedi- and benzenetri-glyoxylic acid were studied by Sawluk and Vass." Radical anions, dianions and trianions are known to be powerful bases which are effective in abstracting a proton from very weakly acidic compounds. Nevertheless, they can also undergo electron-transfer reactions which are by far the most common process occurring even in systems favourable to proton abstraction. '' Significant differences in reactivity are expected depending not only on the electron affinity and charge or spin density distribution, but also on ion pairing.The interest in radical trianion research follows from the expected enhanced reactivity of these species compared with the corresponding less reduced radical anions and dianions. Here we report a detailed EPR, ENDOR and TRIPLE res- onance investigation of the radical anions and radical tri- anions which are formed in the course of the reduction of 5-methylene-5H-dibenzo[u,d]cycloheptene 1 with alkali metals in aprotic solvents. This study has been extended to the radical cation obtained by oxidation with AlCl, in CH,Cl, . For the assign- ment of hyperfine splitting (hfs) to protons in the 10,ll and 12 positions, isotopically labelled derivatives were also included, namely l0-deuterio-5-methylene-5H-dibenzo[u,d-J-cycloheptene (Id) and 5-dideuteriomethylene-5H-dibenzo [a,d]cycloheptene (ld2).For the remaining positions the assignment of hfs was based on theoretical spin density calcu- lations carried out by McLachlan and INDO methods. 1 R', R2, R3=H Id R' = D; R2, R3= H ldp R', R2 = H; R3=D Experimental 5-Methylene- 5H-di benzo [a,d] cyclo hep tene (1) was syn t he- sized in this laboratory from SH-dibenzo[a,d]cyclohepten-5-one (A) (Aldrich) according to the procedure reported in the literature.12 The same procedure was used in the synthesis of 5-dideuteriomethylene-5H-dibenzo[u,d]cycloheptene Id, using CD,MgI instead of CH,MgI. The final product was obtained in 75% yield as colourless crystals, mp 112.5-1 13 "C (from 95% ethanol).6, (300 MHz, solvent DCCl,) 6.812 (2H, s), 7.300 (SH, m); m/z 206, 179, 165, 152, 102. Mass spectral analysis and 'H NMR were both consistent with a 99% deu- teriation. The monodeuteriated derivative 1-d was synthesized from the same ketone A, in the following steps: bromination fol- lowed by dehydrobromination of A according to literature procedure^'^ led to l0-bromo-5H-dibenzo[u,d-Jcyclohepten-5-one (B). The Grignard reaction of methyl iodide with B fol-lowed by dehydration of the alcohol with iodine under vacuum, similar to the procedure described in the synthesis of 1,'' gave l0-bromo-5-methylene-5H-dibenzo[u,d]c~clo-heptene (C)in 52% yield.Recrystallization in hexane gave colourless crystals, mp 78-79 "C (lit.14 mp 80-82 "C). 6, [300 MHz, solvent (CH,),CO] 5.309 (2H, s), 7.400 (7H, m), 7.652 (lH, s), 7.843 (lH, d, J 7 Hz); m/z284, 282, 202, 101. 5 cm3 of butyllithium (7.5 mmol drn-,, 1.5 mol dm-j hexane) was added dropwise to a solution of C (2.72 g, 7.5 mmol drn-,) in 75 cm3 dry tetrahydrofuran cooled with a liquid-nitrogen- ethanol solution ( -100"C) under argon. The solution was stirred for 30 min, after which excess deuterium oxide (ca. 5 cm3) was added dropwise. The reaction mixture was further stirred for 30 min at -100°C and was slowly warmed to room temperature. After the solvents were distilled, the residue was extracted with hexane and the new residue after solvent evaporation was recrystallized from 95% ethanol, giving colourless crystals of Id in 72% yield, mp 114-1 15 "C.6, [300 MHz, solvent (CD,),CO] 5.210 (2H, s), 6.874 (lH, s), 7.360 (SH, m); m/z 205, 190, 179, 164, 152, 102. The isotopic purity achieved was 91%, as was verified by 'H NMR. 2-Methyltetrahydrofuran (MTHF), tetrahydrofuran (THF), 1,2-dimethoxyethane (DME) and hexamethylphosphoric tri- amide (HMPT) were dried by known techniques, stored under vacuum over Na-K alloy and distilled under vacuum directly into the sample tube. Acetonitrile (ACN), dimethyl- formamide (DMF) and dichloromethane (DCM), spectro- scopic grade, were further purified by being passed through a neutral alumina column (activity grade 1), under argon, directly into ampoules where each solvent was kept over molecular sieves 3A (Fluka).These solvents were finally dis- tilled under vacuum directly into the sample tubes. Tetra- butylammonium tetrafluoroborate (Bu,NBF,) and sodium tetraphenylborate (NaBPh,) were dried under vacuum at 100°C using P205 as the drying agent. Commercial alu- minium trichloride (Merck) was sublimed under high vacuum into the sample tube. Preparation of Radical Ions for Spectroscopic Studies The radical cations l*+,Id" and Id2*+were generated uia oxidation of the corresponding neutral compounds by AlCl, in CH,Cl,. The radical anions lo-,Id-and ld2*-were immediately obtained via reduction of the neutral compounds by the alkali metals Li, Na, K and Cs in MTHF, THF, DME as well as in mixtures of these solvents with HMPT.Metal films were used in the experiments involving Na, K and Cs, and metal sand was used for the reductions with lithium. The radical trianions lo,-, ldS3-, ld2'3-were produced by pro- longed contact with the alkali metal (after a few hours for reductions with potassium and caesium and after a day for reductions with lithium and sodium). EPR and ENDOR spectra were recorded on a Bruker ER 200D spectrometer equipped with an EN 810 ENDOR unit and an A-300 RF power amplifier. The temperature of the samples was monitored by means of a Bruker variable- temperature unit ER 400 VT. Voltammetric Measurements The cell used for low-temperature cyclic voltammetry was similar to that described by Terahara et all5 with a platinum wire working electrode and counter-electrode instead of gold wire.The preparation of the sample solutions was carried out in the vacuum line. The reduction studies were run in ACN or DMF containing Bu,NBF, (0.1 mol dm-3) as well as DME containing NaBPh, (0.1 mol drn-,). The oxidation of 1 was studied in DCM containing Bu,NBF, (0.1 mol dm-3) as supporting electrolyte. The solvent was distilled directly from an ampoule into the cell containing 1 (lo-, mol drn-,) and the supporting electrolyte. The solutions were degassed by the freeze-pumpthaw method under a dry and oxygen-free argon atmosphere. Ferrocene (Fc) (lo-, rnol dm-,) was added subsequently as an internal standard in order to cali- J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 brate the potential with the Fc+/Fc couple (Eo = 0.400 V us. NHE16). The cyclic voltammograms of 1 were taken in a PAR 273A Potentiostat/Galvanostat interfaced with a PC equipped with PARC 270 Electrochemical Software. Quenching Experiments These experiments were performed using a double Schlenk tube with the two arms separated by a break-seal. One arm contained excess dry methanol (protonation) or excess dry CH,I (methylation) under nitrogen, which were degassed several times before being sealed under high vacuum. The radical trianion was generated in the other arm, as usual. After exhaustive reduction of 1 and observation of the EPR spectrum from the radical trianion, the solution was trans- ferred through the break-seal to the other arm and mixed with the electrophilic reagent.After treatment with water, the organic layer was separated, washed with 5% HCl and water and dried over MgSO,. The residue after solvent evapo- ration was analysed by mass spectroscopy. Reoxidation Experiments These experiments were carried out in a double Schlenk tube with the two arms separated by a break-seal. The reduction of 1 by the alkali metal was carried out, as usual, in one arm. In the other one, two parts of the neutral compound were dissolved in the same solvent under nitrogen, degassed and then the tube was sealed under high vacuum. After 1 had been exhaustively reduced, the EPR spectrum revealed the presence of lo3-which was then mixed with the unreacted solution via the break-seal.The final spectrum showed that lo,-had been reconverted into lo-. Results Cyclic Voltammetry of 1 The cyclic voltammetric measurements were carried out at temperatures ranging from 210 K to room temperature. As the reduced species, which are formed through electron trans- fer, are very sensitive to low concentrations of protic and electrophilic impurities, the system, including cell, solvent and electrolyte, was carefully purified as described in the Experi- mental. During the cathodic sweep, in the voltage range from -1.2 to -3.0 V, three reduction waves were clearly distin- guishable. The first, corresponding to the formation of the radical anion, was clearly reversible with El,, estimated to be -2.26 V in ACN at 210 K and -2.24 V in DMF at 220 K us.NHE. The anodic peaks of the other two waves are very small compared with the cathodic peaks. The values of the observed cathodic peak potentials are -2.41 and -2.85 V in ACN at 210 K and -2.50 and -2.94 V in DMF at 220 K, us. NHE. In DME containing NaBPh, the voltammograms consisted of ill-defined consecutive electron transfers even at temperatures as low as 210 K. The cyclic voltammograms of 1 exhibited an oxidation wave in the voltage range 0.0 to +2.0 V with an oxidation peak potential of + 1.40 V in DCM at 210 K, us. NHE. Radical Anions Under the preparative conditions described in the Experi- mental, the radical anions la-,Id'-and ld2*-were persistent and could be studied by EPR spectroscopy in the tem-perature range 183-298 K.Analysis of the EPR hyperfine pattern was secured by ENDOR and TRIPLE resonance spectroscopy. The relative signs of the hfs were determined by J. CHEM. SOC. FARADAY TRANS., 1994, Fig. 1 EPR, ENDOR and TRIPLE resonance spectra of 1'-K+ in DME at 193 K: (a)EPR spectrum; (b)ENDOR spectrum; (c) general TRIPLE resonance spectrum, irradiating on the line marked with an arrow; (d) special TRIPLE resonance spectrum. general TRIPLE resonance assuming that the largest absol- ute value is negative. Fig. 1 shows the EPR, ENDOR and TRIPLE resonance spectra of 1'-K+ in DME, observed at 193 K. The values of the proton and deuteron hfs are col- lected in Table 1. The proton hyperfine splitting pattern is not very sensitive to changes in the counter-cation, the solvent Table 1 Experimental hfs of the radical anions l'-, Id-and ld2'-obtained by reduction with different alkali metals" hyperfine splitting constant/mT 1'-b -0.196 +0.037 -0.409 +0.119 -0.434 0.119 (2H) -0.433 (1H) Id--0.198 +0.036 -0.408 +0.119 0.064 (1D) 0.119 (2H) ld2*--0.196 +0.037 -0.409 +0.119 -0.434 0.017 (2D) For the assignments see the Discussion section.Identical hyperfine patterns were obtained with Li, Na,K or Cs in MTHF, THF, DME and in mixtures of DME and HMPT. 3275 or the temperature. Assignment of the largest hfs to protons at the 10,ll-positions were based on the effects of selective deuteriation of one of these positions. The effects of labelling on the EPR and ENDOR spectra were a change of the multi- plicity of the proton hfs (0.434 mT) from two to one and the appearance of an hfs of 0.064 mT, corresponding to one deu- teron.The effect of deuteriation of the methylene group was to change the multiplicity of the proton hfs (0.119 mT) from four to two, and for the deuteron hfs of 0.017 mT, with a multiplicity of two, to appear. This observation allowed two protons with an hfs of 0.119 mT to be assigned to the methyl- ene group. The assignment of the remaining hfs was based on theoretical calculations as discussed later. Hyperfine inter- actions due to the counter-cation were not observed. The g value of 1'-K+ in DME at room temperature was found to be 2.0028 fO.OOO1.Radical Trianions Exhaustive reductions of 1, Id and Id, upon prolonged contact of their ethereal solutions with the alkali metal led to the disappearance of the EPR spectra of the radical anions and subsequent observation of new well resolved EPR spectra identified as the radical trianions 1'3-, ldo3-and ld203-. These species were very persistent and could be studied by EPR and ENDOR spectroscopy in the tem-perature range 183-298 K. Contrary to observations for the radical anions, the ENDOR spectra of the lithium and sodium radical trianion gave evidence for the hyperfine inter- action of two counter-cations with two different alkali-metal hfs. On the other hand, the potassium and caesium radical trianions exhibited two identical hfs from the alkali-metal cations.The simulation of the corresponding EPR spectra ruled out any interaction with a third counter-cation despite the expected strong association of the radical trianions with three positively charged counter-cations. When 1'-was reduced in mixtures of ethereal solvents and HMPA, the spectrum indicated the presence of the above species together with a second radical. After some hours, this new species was the only one to be detected and had a lifetime of up to several months. As will be shown below, this corresponds to a radical trianion probably with a different structural arrange- ment. In the absence of HMPT reduction of 1'-also led to this new spectrum after several days which remained unchanged for long periods of time.Table 2 gives the hyperfine data for the two radical tri- anions (hereafter referred to as A and B, respectively) with different counter-cations and solvents. The relative signs and multiplicities of the hfs were determined by general and special TRIPLE resonance, respectively. The EPR and ENDOR spectra of lo3-3Cs+ in DME at 193 K and lo3-3K+ in DME-HMPA at 193 K are shown in Fig. 2 and 3, respectively. The proton hfs vary only slightly with the counter-cation (Table 2) as well as with the solvent and the temperature. However, the alkali-metal hfs are strongly dependent on the environmental conditions. As for the radical anion, from the analysis of the EPR and ENDOR spectra of the monodeuteriated ldo3-and the dideuteriated ld203-, respectively, the largest hfs were assign- ed to protons in the C( 10) and C(11) positions and the smal- lest one to the methylene protons for both radical trianions.The assignment of the remaining hfs was based on theoretical spin density calculations as described in the Discussion. No alternating linewidth effects were observed in the tem- perature range used for the EPR experiments which might give evidence for the existence of any dynamic equilibrium. The g value of l.3-3K+ in DME (species B) at room tem- perature is 2.0040 & 0.0oO1. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Experimental hfs for the radical trianions la3-,id3-and Id,''-obtained by reduction with different alkali metals in ethereal solvents' hyperfine splitting constant/mT radical trianion C(1, 9) C(2, 8) C(3, 7) C74, 6) C(l0, 11) C(12) (2H) aM1 aM* lS3--DME Li A B -0.171 -0.23 1 +0.026 +0.059 -0.37 1 -0.417 +0.097 +0.099 -0.463 -0.434 0.010 +0.027 -0.073 0.073 ca.0.002 0.013 Na A B -0.167 -0.235 +0.024 +0.058 -0.367 -0.42 1 +0.092 +0.098 -0.447 -0.45 1 0.010 +0.026 -0.135 0.102 0.024 0.051 K A B -0.158 -0.230 +0.029 +0.059 -0.367 -0.418 +0.096 +0.098 -0.444 -0.437 0.007 +0.026 0.053 0.027 0.053 0.027 cs A B -0.153 -0.232 +0.028 +0.059 -0.366 -0.404 +0.097 +0.097 -0.438 -0.428 +0.015 +0.027 0.280 0.280 0.280 0.280 1-3--DME-HMPT Na A B -0.167 -0.235 +0.026 +0.056 -0.368 -0.419 +0.100 +0.102 -0.445 -0.429 +0.010 +0.026 0.133 0.032 0.036 0.01 3 K A B -0.158 -0.232 +0.030 +0.055 -0.376 -0.419 +0.095 +0.100 -0.443 -0.434 +0.011 +0.027 0.021 0.015 0.021 0.015 cs A B -0.156 -0.232 +0.027 +0.057 -0.367 -0.422 +0.102 +0.102 -0.43 7 -0.437 +0.008 +0.027 0.280 0.280 0.280 0.280 la3--THF Na A B -0.170 -0.235 +0.024 +0.059 -0.370 -0.425 +0.102 +0.104 -0.464 -0.454 0.01 1 0.027 0.088 0.1 10 0.066 0.027 la3--MTHF Na A B -0.174 -0.238 +0.059 +0.023 -0.373 -0.430 +0.095 +0.107 -0.455 -0.456 +0.010 +0.033 0.062 0.105 0.011 -0.050 ld' --D ME K A -0.158 +0.029 -0.367 +0.096 -0.444 (1H) 0.007 0.053 0.053 B -0.230 +0.059 -0.419 +0.098 -0.437 (1H) 0.071 (1D) +0.026 0.027 0.027 0.070 (1D) 1dza3--DME K A B -0.158 -0.232 +0.029 +0.055 -0.367 -0.419 +0.096 +0.096 -0.444 -0.437 -- 0.053 0.027 0.053 0.027 For assignments see Discussion.The identification of the radical trianions was corroborated by three further experiments. Reoxidation of the samples with dry oxygen yielded the starting compound as the sole product at any stage of the reduction. This finding demon- strated that the molecular framework of the highly reduced species remained intact excluding the occurrence of any chemical reaction on the primary radical anion. By reoxidation with two equivalents of the neutral com- pound the EPR spectrum of either radical trianion disap- peared, being replaced by the EPR spectrum of the radical anion, according to the comproportionation reaction 1.3-+ 21 -+31'-(1) This finding gave further evidence for the reversibility of the formation of the radical trianion.Quenching experiments performed by adding electrophilic reagents such as methanol or excess dry methyl iodide allowed the number of negative charges in the highest reduced species to be determined. These reactions were per- formed with solutions exhibiting either EPR spectrum of the radical trianions A and B. Di- and tetra-hydro or di- and tetra-methyl adducts were accordingly obtained, and were identified by mass spectrometry. This result is compatible with the presence of a radical trianion in either case. In effect, in highly reduced n systems such as these, intermolecular electron-exchange processes compete with electrophilic addi- tion reactions according to the following disproportionation equilibrium 21'3--12-+ 14-(2)A even when the anion is slowly added to an excess of the elec- trophile.Radical Cation The radical cations l*+,Id*+and Id," generated in CH,Cl, containing AlCl, were studied by EPR, ENDOR and TRIPLE resonance spectroscopy in the temperature range between 213 K and room temperature. The relative signs and multiplicities of the hfs were determined by general and special TRIPLE resonance, respectively. The value of the proton and deuteron hfs are included in the Table 3. In order to assign the hfs to protons in individual postions of lo+,the effect of selective deuteriation at the C(10) and C(12) posi- tions was investigated whilst the remaining hfs were assigned on the basis of theoretical calculations, which reproduced satisfactorily the observed signs of hfs, as discussed later.The undeuteriated radical cation 1'+has two hfs (a = -0.770 and -0.680 mT) with a multiplicity of one, whereas the remain- ing hfs have multiplicities of four (0.276 mT) and two (0.757, Table 3 Experimental hfs of the radical cations l'+,Id+ and Id,'+ in CH,CI, at 223 K" ~ ~ hyperfine splitting constant/mT 1'+ 0.276 0.115 0.276 0.142 0.680 (1H) 0.757 (2H) 0.770 (1H) ld+ 0.276 0.115 0.276 0.142 0.674 (1H) 0.766 (2H) 0.102 (lD) Id2'+ 0.276 0.115 0.276 0.142 0.683 (1H) 0.109 (2D) 0.803 (1H) a For assignments, see Discussion. J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 (c) Fig. 2 EPR, ENDOR and TRIPLE resonance spectra of la3-3Cs+ in DME at 193 K (species A): (a) EPR spectrum; (b) ENDOR spec- trum; (c) general TRIPLE resonance spectrum irradiating on the line marked with an arrow 0.142 and 0.115 mT). This set of hfs denoted a reduced sym- metry of the paramagnetic species, as was clearly shown by its EPR spectrum without a central line (Fig. 4). The two hfs with a multiplicity of one are attributed to the protons in the C(10)-and C(11)-positions while a = 0.757 mT is attributed to the methylene protons. This assignment was made by analysis of the EPR and ENDOR spectra of the radical cations generated from Id and Id, under the same condi- tions. Discussion Radical Anions and Trianioos The observed invariance of the proton hfs of the primary radical anion with the counter-cation, the solvent and the temperature indicated that throughout the temperature range studied these radical anions behaved as solvent-separated ion pairs or even as free-radical anions.The final paramagnetic species obtained upon prolonged reduction with the alkali metals must be identified as the radical trianion. In effect the absence of any structural change resulting from the occurrence of a chemical reaction of the radical anion, together with the reversibility found for the reduction pro- cesses, as well as the results of the quenching experiments, supported its assignment to the radical trianion, 1'3-3M'. Further evidence was given by the observation of two alkali- metal hfs even at low temperatures, e.g.183 K. The observed stability of these radical trianions may result from a strong association with the three positively charged counter-cations. (c1 Fig. 3 EPR, ENDOR and TRIPLE resonance spectra of lS3-3K+ in DME-HMPA at 193 K (species B): (a) EPR spectrum; (b)ENDOR spectrum; (c)general TRIPLE resonance spectrum irradiat-ing on the line marked with an arrow Even in strongly cation-solvating mixtures of HMPT and ethereal solvents such as DME, the interactions of the radical trianion with the counter-ions is very significant. This is sup- ported by the observation of hfs from two Na' ions when the reduction was carried out in this medium (Table 2). The cycloheptatriene ring with an exocyclic carbon double bond must be the relevant structural feature on the anion moiety favouring conversion into the radical trianion.This arrange- ment may be responsible for rendering the energies of the LUMO and of the NLUMO sufficiently close. ZNDO Calculations on the Free-radical Anion 1'-Information about the molecular geometry of the radical species is necessary for the theoretical analysis of the experi- mental hfs. However, while unambiguous structural determi- nation of these paramagnetic species in solution is not possible, the only information has been derived from the theoretical spin-density calculations performed in order to fit the experimental results. Knowledge of hfs including their signs as well as the assignment of the hfs corresponding to C(lo), C(11) and C( 12) positions provided the experimental support for discussing the molecular geometry.The INDO method seems to be the most suitable for spin density calcu- lations on these non-planar systems. Even for the parent neutral compound 1, there is no single-crystal X-ray structure determination reported in the literature. For that reason we began the calculations with the X-ray structural data published for the related compound 4-(5-rnethylene-W- dibenzo[a,dJcycloheptenyl)-1-methylpiperidine. ' Accord- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 "H I I' (d ) Fig. 4 EPR, ENDOR and TRIPLE resonance spectra of 1'+ in CH,Cl, containing AlCl, at 223 K: (a) experimental EPR spectrum; (b) simulated EPR spectrum with linewidth 0.015 mT and the hfs given in Table 3; (c) ENDOR spectrum; (d)special TRIPLE reson- ance spectrum ingly, the central seven-membered ring adopted a boat con- formation with dihedral angles, 8, and e,, of 26.9" and 50.6", respectively, and a dihedral angle, 8,, between the plane of the two benzene rings of 124.1" (Fig.5). Both the order and the signs of the INDO calculated hfs for the free-radical anion, assuming a geometry analogous to that of the above related compound, showed poor correlation with experimental results (Table 4). For that reason we have tried to improve the INDO calculations by making some Fig. 5 Configuration of the central seven-membered ring for 4-45 methylene-5H-dibenzo[a,~cycloheptenyl)-1-methyl-piperidine' adjustments to the structure, particularly by varying the dihe- dral angles el, 8, and 8, .t The range of variation of 8, and 8, was restricted owing to the steric interaction between the protons bound to C(12) and C(4) or C(6) (e.g., for 8, =30", 8, (140" in order to have distances greater than 2 A between the sterically interacting hydrogen atoms).8,,19, and 8, were varied in the ranges 30-lo", 60-10" and 160-120", respec-tively. A much better agreement between the observed and calculated hfs was obtained for a more flattened structure with 8, =20", 8, =30" and 8, =140". This structure is illus- trated in Fig. qa). An adjustment was also made to the C(lO)-C(ll) bond length in order to improve the results [C( lo)--( 11) =1.366 A].The calculated hfs assuming such a structure are shown in Table 4. INDO Calculations on the Contact Radical Trianion lo,-3Li' The molecular geometry of the two radical trianions was much more difficult to establish owing to the uncertainty of the location of the three associated counter-cations. There- fore we only adjusted the calculations for the radical trianion (a) (b ) Fig. 6 Molecular structure assumed in the INDO calculations: (a) radical anion; (b) radical trianion fGeometry optimization by MIND0/3 was also tried, but it was restricted to the ionic R system neglecting both ion pairing and solva- tion effects. The INDO calculated hfs based on such a structure were substantially improved, but the magnitude of the hfs still showed large errors (Table 4).Table 4 INDO calculated hfs for the radical anions 1'-and lS3-assuming different structures hyperfine splitting constant/mT ~~ ~ 1'-l2 +0.214 -0.3 10 +0.065 -0.028 +0.843 1'-b -0.355 +0.241 -0.429 +0.284 -0.347 1'--0.241 +0.151 -0.337 +0.181 -0.397 1'3 --0.29 1 +0.164 -0.309 +0.214 -0.428 l2 Geometry based on the X-ray structure published for the related compound cited in the text.I7 Geometry adjusted as described in the text [Fig. qa)]. ~ ~~ -0.554 ---0.045 --0.099 -0.03 1 Li( 1) -0.01 Li(2) -0.02 Li( 3) -0.036 ~ ~ ~~____________ Optimized geometry by MIND0/3. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 referred to as A.As a starting point, any change in the geometry of the anion moiety which might occur on attrac- tion of the counter-cations was neglected. Note that the cal- culations performed were restricted to the lithium radical trianion owing to the limitations of the available QCPE INDO version which includes only first-row atoms. Note, also, that every satisfactory structure we find is only from a theoretical approach. The location of the counter-cations was tentatively deduced from simple models of electrostatic inter- action and symmetry arguments. Hence, they must occupy positions in the mirror plane of the molecule since no asym- metric spin density distribution was observed. However, they are expected to be situated in the proximity of the highest negative charge-density carbon atoms.Fig. 7 shows the excess negative charge distribution in the radical trianion calculated by the INDO method, excluding the influence of any counter-cation. The largest calculated excess negative charge lies on the set of the exocyclic carbon atom C(12) and its vicinal hydrogen atoms; the next largest excess negative charge lies on the pairs of carbon atoms of the central seven-membered ring C(4a)-C(5a) and C( 10)- C(11) with its bonded hydrogen atoms. Some positions of the benzo rings such as C(2) and C(3) [or C(7) and C(8)] and neighbouring hydrogen atoms are also expected to have considerable excess negative charge density. On the other hand, the folded structure of the mol- ecule (as illustrated in Fig.6) has two sterically different sides relative to the counter-cation attraction, the concave side of the central seven-membered ring being less sterically hindered than the convex one. Therefore, we considered that one Li ion [Li(l)] is more associated with the C(12) atom, whereas the other two [Li(2), Li(3)] are polarized by the remaining excess negative charge centres. We have looked for positions of the Li ions which yielded the best overall agreement of all hfs with experiment. From the countless structures attempt- ed, the one that gave the best fit with experimental results is represented in Fig. qb). In this structure the Li ion associated with C(12) and the neighbouring hydrogen atoms was con- sidered to be situated in the concave side and considerably deviated away from the bulky molecule.The most suitable distances between Li ions and the neighbouring atoms were found to be Li(1)-C(12) = 2.05 A; Li(2)-C(10) = 2.47 8,; Li(3)-C(4a) = 2.52 A whereas Li(1)-Li(2) = 3.13 A. In this structure some adjustments to C(5)-C(12) (1.44 A) and C(lO)-C(ll) (1.41 A) bond lengths were finally made. The calculated hfs are included in the Table 4. MO Calculations The hyperfine data determined experimentally for 1'-and 1'3 -reflect, respectively, the shapes of the LUMO and NLUMO of 1. These MOs were calculated by the Huckel- McLachlan's procedure so that a satisfactory correlation was obtained between the theoretical and experimental hfs. The resonance parameters, ycc , used in these calculations were based on the molecular geometry referred to above as giving the best fit on the INDO calculated hfs.They were corrected -0.142.--0.098 Fig. 7 Excess negative charge distribution in the radical trianion both for the bond length and the twist angle between the two p-orbitals of the bond, according to Streitwieser." A positive Coulomb parameter on the exocyclic carbon atom C(12) was used in order to take into account the strong electron- donating influence of the cycloheptatrienyl ring. A better fit was obtained by further adjustment of the Coulomb param- eter on the remaining carbon atoms according to the w-technique'* with o = 0.9. The resulting LUMO and NLUMO of 1, neglecting the effect of counter-cation associ- ation, are schematically depicted in Fig.8(a). The areas of the circles are proportional to the squares of the LCAO coeff- cients. Blank and filled circles symbolize the different signs of the coefficients. Whereas the hfs of the protons in 1.-, listed in the Table 1, are consistent with the single occupancy of the LUMO, as shown in the Table 5, those for the radical tri- anion lS3-,included in Table 2, disagree with the ones calcu- lated for the NLUMO. Nevertheless, a good correlation was obtained when the effect of the tighter ionic association in the radical trianion was considered in the MO calculations. The cation association was introduced as usual by further adjust- ment of the Coulomb parameters with positive corrections on the neighbouring carbon atoms, in order to account for the +I effect of the positive charge.As illustrated in Fig. 8(b), the calculated energetic sequence of the MO is interchanged. Two reasons for this orbital exchange can be advanced: (i) the LUMO and NLUMO in 1 are almost degenerate, being separated by only a small energy gap (0.025 b) and (ii) the LUMO has two nodes on the C(5) and C(12) positions whereas these carbon atoms have high orbital density in the NLUMO. As discussed previously, those positions are expected to be the most perturbed by ion pairing. Therefore, while the NLUMO of 1 is strongly stabilized owing to the attraction between the oppositely charged ions, the energy of the LUMO is almost unaffected, the cation effect being SUE-cient to shift the Huckel energy of the NLUMO below that of the LUMO as shown in Fig.8. The good fit obtained in the hfs thus calculated for the radical trianion la3-3M+, included in Table 5, allowed the assignment of its SOMO to the NLUMO of Fig. 8(b). The ion-pairing effect decreases the energy necessary to transfer the second and third electrons in the successive reduction steps of the radical anion. Radical Cations Reduced symmetry was observed in the EPR and ENDOR spectra of 1'+ relative to the parent compound. Analysis of the hyperfine pattern of Id' allowed the identification of the NLuMo* @E = a-0.736p E = a-0.6448 LUMo m m E = a-0.71 18 E = a-0.36Sg (a1 (b1 Fig. 8 Schematic presentation of the LUMO and the NLUMO of 1 calculated by the Huckel model (the parameters used in these calcu- lations are included as footnotes to Table 5): (a) not taking into account the counter-cation association ; (b) taking into account counter-cation association J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 5 Calculated hfs by the Huckel-McLachlan procedure, assuming different singly occupied MOs (SOMO) hyperfine splitting constant/mT HOMO" -0.085 +0.013 -0.228 +0.127 -0.566 0.906 HOMO*-' -0.1 34d + 0.M6d -0.230" +0.120" -0.625(10) -0.759 consistent with SOMO of 1'+ -0.798(11) LUMO~.~ -0.201 + 0.067 -0.364 + 0.061 -0.375 +0.029 consistent with SOMO of 1'-NLUMOB + 0.066 -0.303 +0.335 -0.3 10 + 0.303 -1.739 NLUMO~.' -0.189 + 0.058 -0.338 + 0.067 -0.436 +0.010 consistent with SOMO of 1'3-~~ "HMO parameters: Y,,~= Y~,~=~S,~=YS.,~= ~5.44=~5.5a=O-79;1-03; ~2.3=~3,4=~6.7=~7.8=~44,iia=~sa.9a=0~97~ ~io,9a= yll, = 0.84; ylo, ,, = 0, 9; y5.12 = 1.1; a,, = 0.22; 6,, = a,, = 0.15; QzH = -3.0 mT. HMO parameters as " except a,, = -0.05 and 6, = 0.1. aH(l'+)satisfactorily correlate with these calculated values. Two very close hfs were calculated for the two positions. These are the corresponding mean values. HMO parameters: yl, = 78, = 1.03; Y2, = Y3.4 = 76, = Y7,8 = 0.97; Y4.Q = Y6, 5" = 0.968; 75. = Y5.50 = 0-76; Y&, 11" = 75". 9a --0.975; ylo, 9a = y,,, = 0.865; ylo, ,, = 1.04; y5, 12 = 1.08; = 0.014; a,, = S,, = -0.276 (calculated by the o-satisfactorily correlates with these calculated values. HMO parameters as for '.HMOtechnique'8 with w = 0.9); Q:H = -2.3 mT. aH(lo-) parameters as for e except a,, = 1.25; a,, = a,, = -0.176; y5, 12 = 0.9; ylo, ,, = 1.0. ' aH(lS3-)satisfactorily correlates with these calculated values. non-equivalence of the two vicinal carbon atoms C(10) and already used for the radical anion and calculating the C(1 1) (Table 3). These two positions have simultaneously Coulomb parameters by the o-techniqueI8 with o = 0.9 as high spin density and the highest excess positive charge well. The loss of symmetry referred to above can be ascribed density. to some specific interaction with the medium, namely with a Fig. 9(a) shows the HOMO of the radical cation 1" calcu-counter-anion such as AlCI,-, the coordination being prefer- lated by the Hiickel-Mclachlan approximation.The areas of entially oriented to one of the highest positively charged the circles are proportional to the squares of the LCAO coef-C(l0) or C(11) positions. The effect of this perturbation was ficients. Blank and filled circles symbolize the different signs simulated using the HMO approximation by an additional of these coefficients. The excess positive charge distribution is negative Coulomb parameter defining this carbon atom. The indicated in Fig. 9(b) and is obviously symmetric. These cal- HOMO thus calculated is shown in Fig. 9(c). The calculated culations were performed assuming the resonance parameters hfs, assuming the single occupancy of this HOMO, correlated well with the observed hyperfine pattern for 1'+ as shown in Table 5 which substantiates the arguments used.References 1 F. Gerson, R. Heckendorn, D. 0. Cowan, A. M. Kini and M. Maxfield, J. Am. Chem. SOC., 1983,105,7017. 2 W. Huber, Tetrahedron Lett., 1983,24, 3595. 3 W. Huber, Helv. Chim. Acta, 1983,66, 2582. 4 F. Gerson and W. Huber, Acc. Chem. Res., 1987,20, 85. +O. 149 5 M. Hirayama, A. Seki, Y. Yamashita, T. Suzuki and T. Miyashi, Chem. Lett., 1988,67 and 769. 6 M. Hirayama, A. Seki, Y. Yamashita, T. Suzuki and T. Miyashi, J. Chem. SOC.,Chem. Commun., 1988,490. 7 F. Gerson, G. Gescheidt, K. Hafner, N. Nimmerfroh and B. Stowasser, Helv. Chim. Acta, 1988, 71, 101 1.+0.027 8 A. de Meijere, F. Gerson, B. Konig, 0. Reiser and T. Wellauer, +0.040 J. Am. Chem. SOC., 1990,112,6827. 9 H. Fujita and H. Ohya-Nishiguchi, J. Chem. SOC., Chem. Commun., 1989,1091. 10 A. Sawluk and J. Voss, Angew. Chem., Int. Ed. Engl., 1989, 28, 906. 11 M. Szwarc, Ions and Ion Pairs in Organic Reactions, 1974, vol. 2, Wiley, New York, ch. 1. 12 A. C. Cope and S. W. Fenton, J. Org. Chem., 1951, 1673. 13 W. Treibs and H-J. Klinkhammer, Chem. Ber., 1951,84, 671. 14 F. Hoffmann-La Roche and Co., A.G.; Belg. Pat. 659 599, 1965; Swiss Appl., 1964; Chem. Abs., 64, P 5023g. (c1 15 A. Terahara, H. Ohya-Nishiguchi and N. Hirota, J. Phys. Chem., 1986,90,1564. Fig. 9 Schematic presentation of the HOMO of 1 and excess posi- 16 R. R. Gagne, C. A. Koval and G. C. Lisensky, Inorg. Chem. tive charge distribution in 1" calculated by the Huckel model (the 1980,19,2854. parameters used in the calculations are included as footnotes to 17 B. Birknes, Acta Crystallogr., Sect. B, 1977,33, 687. Table 5): (a) calculated HOMO, not taking into account any 18 A. Streitwieser Jr., Molecular Orbital Theory for Organic Chem- counter-anion association; (b) calculated excess positive charge dis-ists, Wiley, New York, 1961, ch. 4. tribution, not taking into account any counter-anion association; (c) calculated HOMO taking into account counter-anion association Paper 4/02896A; Received 16th May, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003273

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3281-3285 Mean Activity Coefficients of NaCl in Glucose-Water and Sucrose-Water Mixtures at 298.15 K Jianji Wang,* Wenbin Liu, Jing Fan and Jinsuo Lu Department of Chemistry, Henan Normal University, Xinxiang, Henan 453002, People's Republic of China The mean activity coefficients of NaCl in aqueous solutions of 10, 20 and 30 mass% sugar (glucose and sucrose) have been determined at 298.15 K from emf measurements in the molality range 0.006-2.0 rnol kg-'. The results have been analysed by using the Debye-Huckel extended equation and the Pitzer equation. There is good agreement between the results obtained from these theoretical models. It has been shown that two parameters of the Pitzer equation, fl(O) and ,!?('), increase linearly with the increasing reciprocal of the relative permittivity for the mixed solvents as well as the mole fraction of sugar in the mixed solvents.A similar linear relationship was also found for the ion-interaction parameter, C, of the Debye-Huckel extended equation. There is a growing interest in the determination of activity coefficients of electrolytes in mixed solvents, particularly at high electrolyte molalities. This interest is, in part, because of applications in such areas as the quantitative determination of Gibbs energies of dilution,' calculation of thermodynamic solubility products (together with solubility) and transfer Gibbs energies of electrolyte into mixed solvents,2 and pre- diction of vapour-liquid equilibrium data of ternary systems containing salt,3 where it is essential to have accurate activity coefficient values for electrolyte in mixed solvents.There have been many studies on the determination of activity coefficients for electrolytes in mixed However, most of these are the 'by-products' of measure- ments of other thermodynamic properties, such as Gibbs energies. These data are usually limited to very low electro- lyte molalities. Relatively few studies on the activity coeffi- cients at high electrolyte molalities in mixed solvents have been rep~rted.~.~-' Recently, we have investigated the solvation of some elec- trolytes in glucose-water and sucrose-water mixtures as well as the interactions of the electrolytes with glucose and sucrose in water.' '*I2 As part of the continuing study of ther- modynamic properties of water-sugar-electrolyte systems, we now report the mean activity coefficients of NaCl in aqueous solutions of 10,20 and 30 mass% sugar (glucose and sucrose) in the electrolyte molality range ca.0.006-2.0 mol kg-'. Based on these activity coefficient data, Pitzer parameters, 8'') and /I(')and the Debye-Hiickel ion-interaction param- eter, C, for NaCl were determined in different aqueous sugar solutions and are discussed in terms of the properties of the mixed solvents It was expected that this investigation would provide additional information on ion-ion interactions in these ternary systems. Experimental Reagents Anhydrous glucose (analytical reagent, Shanghai Chem.Co.) and sucrose (analytical reagent, Beijing Chem. Co.) were dried under vacuum at 343 K to constant weight and stored over P20, in a desiccator before use. NaCl (analytical reagent, Shanghai Chem. Co.) was recrystallized and dried under vacuum. Conductivity water with a conductivity of 1.2 pQ-' cm-' was prepared by distilling the deionized water from basic KMnO, in an all-Pyrex still. In order to minimise the experimental error, stock aqueous NaCl solutions were used to prepare sugar-NaC1-water ternary solutions when the concentrations of NaCl were below 0.1 mol kg-'. In other cases, all test solutions were made by direct weighing of water, sugar and NaC1. Determination of Activity Coefficients The emf method for determination of the activity coefficients of NaCl in aqueous sugar solutions was chosen.The differ- ence in emf of the cell (A) relative to the reference cell (B): Na glass INaCl(m), H20( 100-Y) ,sugar(Y)I AgCl I Ag (A) Na glass I NaCl(m,), H20(100-Y) , sugar(Y) I AgCl I Ag (B) was determined over a range of electrolyte molalities. In cells (A) and (B), m is the molality of electrolyte defined per kg of (sugar-water) mixture and Y mass% the amount of sugar in the mixed solvents. The purpose of introducing the reference cell (B) is to compensate the asymmetry potential of the glass electrode. The molaity of NaCl in cell (B), m,,was kept con- stant for each given solvent. A sodium-glass electrode, model 312 (Jiangsu, China) was used in this work.The Ag/AgCl electrodes were prepared by a thermal-electrolytic method. I3 The measurements were per- formed only with Ag/AgCl electrodes whose potentials showed an internal difference of less than 0.04 mV. The Nern- stian behaviour of the sodium-glass electrode was checked by measuring the difference in emf between cells (A) and (B) when Y = 0 (pure water) with the molality of NaCl varying between 0.006 and 2.0 mol kg-' (see Table 1, later). Using activity coefficients of NaCl in water calculated by the empirical equation recommended by Hamer and Wu,14 an excellent linear relationship with a correlation coefficient of 0.999996 was obtained between AE and log(my,) [see eqn. (1) later]. From the slope of this linear equation, the Nernst slope was found to be 59.10 f0.04 mV, which is in excellent agreement with the theoretical value.The H-shaped measuring cells were made of Pyrex glass. Emf measurements on cells (A) and (B) were carried out using an Orion PH meter (model 720A) with a resolution of 0.1 mV. In all measurements, the temperature inside the cells was kept constant at 298.15 f0.02 K by using a thermostatted water bath. To provide highly accurate experimental results, the method for eliminating the asymmetry potential proposed by Feakins et al." was applied. The experimental details are similar to those described in our earlier papers.' The molality of NaCl in cell (A) varies from CQ. 0.006 to 2.0 mol kg-' for each given solvent, whereas that in reference cell (B) was always constant at CQ. 0.1 mol kg-'.The molality of NaCl in all the cell solutions was accurate to ca. 0.02%. 3282 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 The amount of sugar in the mixed solvents was known to Table 2 Experimental AE and yk values at different molalities of within kO.005mass%. NaCl in sucrose-water mixtures at 298.15 K rn/mol kg-' AE/mV yk rn/mol kg-' AEfmV Y+ Results Y=lO According to the Nernst equation, the difference in emf, AE, 0.005983 134.7 0.92 13 0.008001 120.4 0.9099 between cells (A) and (B) in a given solvent can be expressed 0.009998 109.5 0.9003 0.01OOO 109.4 0.9018 0.02000 76.2 0.8604 0.03000 56.65 0.8391by 0.03999 43.0 0.8210 0.O4001 42.9 0.8222 (1) 0.04999 32.65 0.8034 0.05999 23.85 0.7945AE = 2kN logCmr(y*),/myf1 0.08002 10.5 0.7723 0.09995" 0.0 0.7585 where kN = 2.303RT/F is the Nernst slope, yk and (y*), are 0.1500 -19.3 0.7358 0.2998 -51.65 0.6909 the activity coefficients of NaCl at concentrations m and M,, 0.4499 -70.75 0.6677 0.6005 -84.3 0.6512 respectively.The measured AEs for different molalities of 0.7998 -98.45 0.6439 0.9999 -109.9 0.6436 NaCl in water and in different glucose-water and sucrose- 1.1995 -119.1 0.6417 1.3985 -127.2 0.6444 1.5984 -133.95 0.6429 2.0001 -146.4 0.6546water mixtures are listed in Tables 1 and 2, respectively. Although the value of (y*), remains constant for each mixed Y = 20 0.005992 134.5 0.9 124 0.007999 120.2 0.9028solvent studied, it has to be known exactly in order to calcu- 0.01 000 109.25 0.8937 0.02000 76.1 0.8518late y* values in these mixed solvents.Therefore, the Debye- 0.02998 56.7 0.8289 0.03999 43.15 0.8089 Huckel extended equation and the Pitzer equation were used 0.04996 32.4 0.7981 0.05996 23.85 0.7854 here to determine (y*), values, together with the ion-0.08000 10.5 0.7633 0.09997" 0.0 0.7493 interaction parameters of these equations. 0.1500 -19.0 0.7230 0.3000 -51.52 0.6769 0.4500 -70.5 0.6564 0.5998 -84.2 0.6429 0.7998 -98.4 0.6356 1.OOO1 -109.65 0.6327 Table 1 Experimental AE and y* values at different molalities of 1.2000 -119.25 0.6356 1.3994 -127.35 0.6381 NaCl in water and in glucose-water mixtures at 298.15 K 1.5996 -134.65 0.6435 1.9997 -147.4 0.6597 Y = 30 m/mol kg-' AEImV y* m/mol kg-' AE/mV y* 0.007995 119.5 0.9OOo 0.01197 99.85 0.8811 0.02008 75.35 0.8452 0.02996 56.35 0.8208Y=O 0.o4001 42.95 0.7977 0.04988 32.45 0.7849 0.006000 135.7 0.9222 0.008002 122.1 0.9010 0.06Ooo 23.65 0.7744 0.06003 23.75 0.7725 0.009998 111.2 0.8915 0.02000 76.5 0.8755 0.07 9 9 9 10.4 0.7518 0.09999" 0.0 0.7363 0.02999 57.0 0.8534 0.04000 43.2 0.8369 0.1201 -8.65 0.7254 0.1500 -18.8 0.7076 0.05000 32.8 0.8197 0.06001 23.95 0.8114 0.3000 -51.35 0.6666 0.4500 -70.4 0.6439 0.08002 10.5 0.7905 0.1oOo" 0.0 0.7761 0.6002 -84.2 0.6315 0.8009 -98.65 0.6269 0.1250 -10.75 0.7653 0.1500 -18.95 0.7480 0.9998 -109.9 0.6251 1.1988 -119.8 0.6321 0.3000 -51.4 0.7033 0.4502 -70.5 0.6797 1.4005 -127.95 0.6340 1.6015 -135.4 0.6410 0.6OoO -84.3 0.6671 0.8002 -98.2 0.6556 2.001 1 -148.4 0.6606 1.0003 -109.3 0.6509 1.1985 -118.6 0.6510 1.3991 -126.5 0.6504 1.5957 -133.65 0.6554 " Molality of NaCl in reference cell (B).1.9963 -145.95 0.6655 Y=lO 0.005998 134.05 0.9307 0.007994 120.45 0.9099 For a 1 :1 electrolyte, the Debye-Huckel extended 0.00999 1 109.95 0.8930 0.02000 76.2 0.8604 equation16 is given by 0.02997 56.75 0.8383 0.03999 42.95 0.8218 log y* = -Am'/'/(l +Biim'/')0.04999 32.55 0.8049 0.05991 24.25 0.7894 0.08000 10.35 0.7748 0.09999" 0.0 0.7582 +Cm -log(1 +0.002mMJ (2)0.1495 -18.9 0.7325 0.2998 -51.6 0.6903 0.4497 -70.1 0.6596 0.5994 -84.35 0.6530 where d is the ion size parameter, C the ion-interaction 0.7995 -98.4 0.6435 0.9999 -109.7 0.6411 parameter and M, the mean molar mass of the mixed solvent.1.1985 -119.05 0.6416 1.3995 -127.15 0.6433 1 S994 -134.1 0.6444 1.9983 -146.7 0.6591 A and B are the conventional Debye-Huckel constants given by the equations Y = 20 0.007993 119.25 0.9054 0.009998 108.2 0.8975 A/(mol-'/' kg1I2 K3/') = 1.8246 x 106d'/'(&,T)-3/2 (3)0.02000 75.2 0.8528 0.02998 56.0 0.8266 B/(cm-'mol-'/' kg'" K'12) = 50.29d'/'(~,T)-'/~ (4)0.03998 42.45 0.8069 0.05000 32.1 0.7891 0.05996 23.65 0.7757 0.08000 10.35 0.7531 0.09995" 0.0 0.7373 0.1500 -19.0 0.7111 where d and E, are the density and relative permittivity of the 0.2999 -51.2 0.6655 0.4498 -70.3 0.6435 solvent, respectively. For the sugar-water mixtures studied, 0.6004 -84.25 0.6325 0.7995 -98.0 0.6207 density values were calculated from the equations given by 0.9993 -109.55 0.6217 1.1997 -118.7 0.6188 Daldrup and Schonert.' Values of the relative permittivity 1.3992 -127.45 0.6291 1.5991 -134.5 0.6314 were taken from the 1.9994 -147.65 0.6522 Inserting eqn. (2)into eqn.(l),it follows that Y = 30 0.009998 107.7 0.8940 0.01999 74.75 0.8490 AEd =AE +2kN log(m/rn,) -2k,[Am"2/(1 +B6m1/')] 0.02998 56.25 0.81 14 0.03999 42.6 0.7934 -2kN lOg(1 +0.002mM30.04998 32.0 0.7803 0.05996 23.8 0.7629 0.08000 10.25 0.7443 0.09999" 0.0 0.7270 0.1500 -18.7 0.6973 0.3001 -51.0 0.6535 0.4500 -70.2 0.6333 0.5997 -83.9 0.6204 It is clear from this equation that there is a linear relationship 0.7954 -97.7 0.6118 1.0003 -109.7 0.6145 between AEd and m. A least-squares analysis was used to 1.1999 -119.4 0.6187 1.4OOO -127.9 0.6256 1.6006 -135.3 0.6320 2.0011 -148.3 0.65 10 determine (y*), and the parameter C for each of the mixed solvents from the experimental AEs listed in Tables 1 and 2." Molality of NaCl in reference cell (B). The values obtained are shown in Table 3, together with their J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 3 Parameters of the Debye-Huckel extended equation for Table 4 Parameters of the Pitzer equation for sugar-water mix-sugar-water mixtures at 298.15 K tures at 298.15 K glucose-water 0 0.7763 0.0011 0.0536 0.0008 3.9 0.26 -0.9984 10 0.7580 0.0012 0.0666 O.ooo9 3.6 0.28 -0.9985 20 0.7366 0.0012 0.0848 O.ooo9 3.2 0.27 -0.9991 30 0.7264 0.0013 0.0903 0.0010 3.4 0.30 -0.9991 sucrose-water 10 0.7581 O.ooo9 0.0634 0.0007 3.7 0.22 -0.9989 20 0.7489 O.OOO5 0.0730 O.OOO4 3.7 0.12 -0.9998 30 0.7354 0.0007 0.0805 O.OOO6 3.8 0.18 -0.9995 standard deviations, s, and s, ,respectively, the standard devi- ation of the fit, 6,and the correlation coefficient, R.In the regressions, h was regarded as an adjustable parameter. The optimum 4 values will be those which lead to a minimum standard deviation of the fit. Values of h are also included in Table 3. The relevant equations for the activity coefficients of 1 : 1 electrolyte derived by PitzerZ0 are In yk =f, + B,rn + C,m2 (6) where f,= -A 4JX B, = 2p'"' + 2B"'y and y = [I -exp(-arn'/2)(1 + -0.5a2rn)]/(a2rn) In the above equations, b and a are parameters, A, is the Debye-Huckel coefficient for the osmotic functions. At 298.15 K A, can be calculated from eqn.(7) A mol-'/2 kg'/2) = 272.058d'/2 E~-~/~ (7) It have been shown by Koh et d2'that the same values of b and 01 for aqueous solutions could be used for the methanol- water mixtures without greatly affecting the standard devi- ation of the fit. Therefore, the values of b = 1.2 kg'/2 mol- 'I2 and a = 2.0 kg'j2 mol-'/2 were used in the present calcu- lations. In addition, for electrolytes whose concentrations do not exceed 2 mol kg-', the term containing C,in eqn. (6), which accounts for triple ion short-range interactions, may be neglected.20 At least, this is true for the water-rich mixed solvents.8 Bearing these simplifications in mind, and intro- ducing eqn.(6) into eqn. (l), the following equation was obtained AEp = AE + (2RT/F)ln(rn/mr)+ (2RT/F)f, = (2RT/F)ln(y,), -(4RT/F)B(")rn-(RT/F)@') The experimental AEs at different electrolyte molality for each of the mixed solvents were fitted to eqn. (8) by a multi- ple regression program with the A, values calculated from the density" and relative for the different sugar-water mixtures investigated. The values of (y*),, /I(') and /?('I are summarised in Table 4 together with their stan- dard deviations sr, so and sl, respectively, and the standard deviation of the fit, 6. As can be seen from Tables 3 and 4, within experimental error, there is a good agreement between the (y*), values glucose-water 0 0.7759 0.0015 0.0813 0.0024 0.231 0.016 0.25 10 0.7584 0.0018 0.0940 0.0030 0.232 0.020 0.30 20 0.7380 0.0018 0.1117 0.0030 0.251 0.020 0.30 30 0.7275 0.0022 0.1227 0.0034 0.294 0.023 0.33 sucrose-water 10 0.7588 0.0013 0.0899 0.0024 0.256 0.015 0.24 20 0.7496 0.0008 0.1028 0.0013 0.296 0.009 0.14 30 0.7371 0.0017 0.1142 0.0029 0.368 0.020 0.29 obtained from the two theoretical models for each of the mixed solvents.Therefore, the mean of two (y*), values for a given solvent was used to calculate the y* for NaCl in this solvent by using eqn. (1). The yk values thus obtained are also included in Tables 1 and 2. Discussion In the molality range studied, the standard deviation of the activity coefficients between our experimental values and those reported in the literat~re'~ was calculated to be k0.005 in pure water solvent.Considering the fact that the two sets of y* data were determined by different experimental tech- niques, it would appear that the agreement between our values and those from the literature is quite good. To the best of our knowledge, no activity coefficient data for NaCl in glucose--water and sucrose-water mixtures have been re-ported in the literature. Analysis of the activity coefficient data for NaCl in sugar- water mixtures shows some interesting features. As the pro- portion of the mixed solvent is fixed, the activity coefficients first decrease and then increase with increasing molality of the electrolyte. On the other hand, they decrease with increasing content of sugar in the mixed solvents when the molality of the electrolyte is fixed.Similar trends of the activ- ity coefficient have been noticed for NaBr in methanol-water rnixt~res,~for NaCl in ethanol-water mixtures' and for HCl in aqueous solutions of methan01,~ ethanol, glycerol and 1,4-dioxane.lo The Pitzer equation was based on a semiempirical theory of statistical mechanics. Although this equation has been used successfully for representing activity coefficients of elec- trolyte in mixed solvent^,^^^.^ not much effort has been devoted to understanding the nature of the Pitzer parameters in mixed solvents.2'.22 Values of pC0)and fl(') determined in this work are positive in all the solvents, indicating a net repulsive force in short- range interactions and positive second virial coefficients in sugar-water-NaC1 systems.22 This is the typical profile observed with 1 : 1 electrolytes (and most others) in aqueous solutions.23 Furthermore, it is interesting to note that both Po)and /3(') vary linearly with the reciprocal of relative per- mittivity for sugar-water mixtures (see Fig.1 and 2). Similar linear plots can also be obtained between these parameters and the mole fraction of sugar in the mixed solvents. On further analysis of the results of alkali-metal chlorides and HCl in methan~l-water~.~' and of NaCl in ethanol-water,' it is shown that B(") and /?(I) for these electrolytes in these solvent mixtures indeed increase linearly with increasing 1/~,, except for LEI, KC1, RbCl and CsCl in methanol-water sol-vents, where the values for /?(") are almost constant or decrease slightly with increasing 1/~,.An example is given in Fig. 3. 3284 -r6 s 's 0.0 1.2 1.3 1.4 1.5 102/E, Fig. (0)P(') and (0)B(')for NaCl in glucose-water mix ires as a function of relative permittivity In the discussion of hard-core effects on osmotic and activ- ity coefficients in terms of the extended form of the Debye- Huckel theory, Pitzer2' has pointed out that B(')would be a function of rc2. Because ic2 involves the reciprocal of the rela- tive permittivity as shown in eqn. (9), it is expected that /3(') would be a function of 1/~,.rc2 = 2e2~, E~ kT) (9)rnd/(~, 0.4 0.3 6 o^60.2 0.1 0.0 1.2 1.3 1.4 1.5 102/&, Fig. 2 (0)B(O) and (0)8"' for NaCl in sucrose-water mixtures as a function of relative permittivity 2.0 1.5 -r -P 1.0 s 's 0.5 v0.0 ' -1.o 1.5 2.0 2.5 3.0 102/&, Fig. 3 (0)B'O) and (0)B(') for NaCl in methanol-water mixtures as a function of relative permittivity J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 where k is the Boltzmann constant and other symbols have their usual meanings. According to Pitzer,20 the main contribution to /?(') comes from the short-range interactions of unlike charged ions, while B(O) is determined by the short-range interactions of both like and unlike charged ions.The radial distribution functions at hard-core contact, g+ -, g+ + , g--, therefore, contribute in weighted proportion to #I(')whereas only g+ -contribute to /3('). In Pizter's thermodynamic treatment of electrolyte solutions, there is no explicit dependence of /?(') on relative permittivity. However, the radial distribution func- tions at hard-core contact were found to be a function of the relative permittivity of the solvents.22 Thus the dependence of pCo)on E, observed in the present investigation is understand- able. The radial distribution functions, g+ -,g+ + and g--,for methanol-water mixtures have been calculated by Gupta22 from an exponential form of the Debye-Huckel theory, which has been proven by Monte Carlo calculations to give identi- cal results for aqueous 1 : 1 electrolyte up to 0.4 mol kg-'.24 It has been shown that g+ -increases whereas g+ + and g--decrease with increasing content of methanol in the mixed solvents. Using the same equations, we have calculated these radial distribution functions at hard-core contact for glucose- water, sucrose-water and ethanol-water mixtures.The same trends of g+ -, g+ + and g--with the composition of the organic component in these mixed solvents as in methanol- water mixtures were observed. Since the relative permittivity of the solvents studied decreases with increasing content of the organic component in the mixed s~lvents,~*'~*'~~~~ it is expected that /3(') would increase with 1/~,.This is consistent with the results discussed above. Moreover, as 8'') values for NaCl in aqueous solutions of glucose, sucrose, methanol and ethanol, and for HCI in aqueous methanol solutions are found to increase with increasing l/~,,it would appear that they are influenced more by the interactions of unlike-charged ions than by those of like-charged ions in these systems. On the contrary, /3(") values decrease with increasing 1/~,for the other alkali-metal chlorides in methanol-water mixtures, indicating that in these cases the parameters are influenced more by the interactions of like-charged ions. Finally, it is interesting to note that the Debye-Huckel ion- interaction parameter, C, is numerically close to @') for each of the sugar-water mixtures. Also, it increases linearly with 1/~,and the mole fraction of sugar in the mixed solvents studied.Financial support from the Natural Science Foundation of China and the Natural Science Foundation of Henan Prov- ince is gratefully acknowledged. References 1 J. A. Rard, J. Solution Chem., 1990, 19, 525. 2 J. W. Lorimer, Pure Appl. Chem., 1993,65, 183. 3 H. Pan, S. Han and Y. Yao, J. Chem. Znd. Eng. (China), 1992,43, 360. 4 A. K. Covington and T. Dickinson, Physical Chemistry of Organic Soloent Systems, Plenum Press, New York, 1973. 5 Y.Macpherson and R. Palepu, Can. J. Chem., 1993,71,2038; D. Chu, Q. Zhang and R. Liu, J. Chem. Soc., Faraday Trans. I, 1987, 83, 635; A. K. Covington and J. M. Thain, J. Chem.Soc., Faraday Trans. I, 1975,71,78. 6 Z. Kozlowski, A. Bald and J. Gregorowicz, J. Electroanal. Chem., 1990,288,75; Thermochim. Acta., 1990,170,217. 7 P. Longhi, P. R. Mussini, T. Mussini and S. Rondinini, J. Solu-tion Chem., 1988, 17, 423; J. Radosevic and I. Mekjavic, Electro-chim. Acta, 1983, 28, 1435; V. V. Sastry and C. Kalidas, J. Chem. Eng. Data, 1983, 28, 5. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3285 8 M. A. Esteso, 0. M. Gonzalez-Diaz, F. F. Hernandez-Luis and 16 R. A. Robinson and R. H. Stokes, Electrolyte Solutions, Butter- L. Fernandez-Merida, J. Solution Chem., 1989,18,277. worth, London, 2nd edn., 1959. 9 L. Zhang, X. Lu, Y. Wang and J. Shi, J. Solution Chem., 1993, 17 N. Daldrup and H. Schonert, J. Chem. SOC., Faraday Trans. I, 22, 137. 1988,84,2553. 10 H. S. Harned and B. B. Owen, The Physical Chemistry of Elec- 18 J. P. Williams, S. B. Knight and H. D. Crockford, J. Am. Chem. trolyte Solutions, Reinhold, New York, 3rd edn., 1958. 11 J. Wang, W. Liu, T. Bai and J. Lu, J. Chem. SOC., Faraday 19 SOC.,1950,72, 1277. C. G. Malmberg and A. A. Maryott, J. Res. Natl. Bur. Stand., Trans., 1993, 89, 1741. 1950,45, 299. 12 J. Wang, L. Zeng, W. Liu and J. Lu, Thermochim. Acta, 1993, 224,261. 20 21 K. S. Pitzer, J. Phys. Chem., 1973,77,268. D. S. P. Koh, K. H. Khoo and C. Y. Chan, J. Solution Chem., 13 G. J. Ives and G. J. Janz, Reference Electrodes, Theory and Prac- 1985, 14, 635. tice, Academic Press, New York, 1961. 22 A. R. Gupta, J. Phys. Chem., 1979,83,2986. 14 W. J. Hamer and Y. C. Wu, J. Phys. Chem. Ref: Data, 1972, 1, 1047. 23 24 K. S. Pitzer and G. Mayorga, J. Phys. Chem., 1973,77,2300. D. N. Card and J. P. Valleau, J. Chem. Phys., 1970,52,6232. 15 D. Feakins, R. D. O’Neill and W. E. Waghorne, J. Chem. SOC., Faraday Trans. I, 1982,78,1431. Paper 4/01955E; Received 31st March, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003281

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3287-3292 Preferential Solvation in Acetonitrile-Water Mixtures Relationship between Solvatochromic Parameters and Standard pH Values Jose Barbosa' and Victoria Sanz-Nebot Department of Analytical Chemistry, University of Barcelona, Diagonal, 647.08028,Barcelona, Spain Standard pH values; pH(s), for five reference buffer solutions, KH tartrate, KH, citrate, KH phthalate, acetate buffer and phosphate buffer in acetonitrile-water mixtures containing 10, 30, 40, 50, 70 and 100 wt.% acetonitrile at 298.15 K have been determined using the IUPAC standardization rules. The relationship between pH(s) and solvent composition, expressed as a fraction, have been studied, with a view to assessing the presence of preferential solvation effects.In order to obtain pH(s) values for all possible acetonitrile-water mixtures, the linear solvation energy relationships method, LSER, has been applied. The pH(s) values were then correlated with the Kamlet-Taft, K*, a and /3 solvatochromic parameters of the acetonitrile-water mixtures. The equation obtained permits the standardization of potentiometric sensors in these mixtures. Solvent mixtures have attracted interest because of their fre- quent use and wide field of applications. A prominent binary mixture is acetonitrile-water. Aqueous acetonitrile is used in many branches of chemistry, ranging from hydrometallurgy to liquid chromatography, from reaction media for synthesis to electrochemistry, and the properties of these mixture are therefore of considerable interest.Accurate pH measurements in the more widely used binary acetonitrile-water mixtures are needed, since pH is potentially useful for optimizing met hods. Accurate determination of the pH of reference buffer solu- tions for standardization of potentiometric sensors is the key to the pH-metric problem in aqueous organic solvent mix- tures such as acetonitrile-water.' *' The internationally recog- nized operational equation used for electrometric pH measuremen t3 is pH(x) = pH(s) + -(E, -Ex) 9 where Ex and E, denote the emf measurements, in cell A, in the sample solutions at unknown pH(x) and in the reference standard buffer solution at known pH(s), respectively, and g = log(RT/F) The reference pH values of the standard buffer solutions are influenced by the nature of the solvent and in acetonitrile-water mixtures they vary with solvent composi- tion in a manner that is not easily understood.Preferential solvation in binary solvent mixtures is impor- tant in solution chemistry for explaining spectroscopic, equi- librium and kinetic In general, preferential solvation is a composite effect determined by solute-solvent inter-actions and solvent-solvent interactions. The quasi-lattice quasi-chemical (QLQC) theory of prefer- ential solvation developed by Marcusg*'* can be applied to quantify the preferential solvation of the hydrogen ions in ace t oni t rile-w ater mixtures . The main quantity obtained from the theoretical approaches is the local mole fraction of one of the components of the mixtures, say water, around hydrogen ions, xf; .For the QLQC method 'local' means nearest-neighbour. Another way to express this is the preferential solvation, i.e.the excess or deficiency of the local mole fraction relative to the corre- sponding bulk quantity: 6, = x4 -x, is the preferential sol- vation of water around ions. The QLQC method permits the calculation of the preferential solvation parameter as a func- tion of the composition of the solvent, on the basis of infor- mation that is independent of the transfer of the ions to the solvent mixture. Central to the study of pH(s) in acetonitrile-water mixtures is the problem of how the solvated species reacts to changes in the environment. Certain microscopic phenomena of solutes are influenced by the solvent that forms their solva- tion sphere.This does not normally have the same properties as the bulk solvent, or even the same composition. In order to characterize this zone of the solvent, a group of solvatoch- romic parameters has been proposed. Recently' 'the polarity and ability to form hydrogen bonds in binary aqueous mix- tures of acetonitrile have been measured by the Kamlet-Taft solvatochromic parameters n*,a and /3."~'~ The polarity and polarizability of these solvents are measured by n*, the hydrogen-bond accepting ability is measured by fi, and the ability to donate a hydrogen atom for the formation of a hydrogen bond is measured by a.On the other hand, the solvatochromic parameter EJ30), proposed by Dimroth and Reichardt14 and the normalized EA30) parameter, E;, in ref- erence to sulfolane (EY = 0) and water (EY = I), have been used to study preferential solvation in binary solvent mix- '."ture~.~.' E: values are already known for acetonitrile- water mixtures." The purposes of the present study were: (a) to assess pH(s) values for five standard buffer solutions: 0.05 mol kg-' pot-assium dihydrogen citrate (KH, citrate); 0.1 mol 1-' acetic acid-0.1 mol 1-' sodium acetate (acetate buffer); 0.025 mol kg-' potassium dihydrogen phosphate-0.025 mol kg-sodium hydrogen phosphate (phosphate buffer); a saturated solution, at 25"C, at potassium hydrogen tartrate (KH tartrate) and 0.05 mol kg- potassium hydrogen phthalate (KH phthalate), in acetonitrile-water mixtures containing 10, 30, 40,50, 70 and 100 wt.% acetonitrile, according to the criteria recently endorsed by IUPAC;'v2 (b) to apply the QLQC theory to the study of the preferential solvation of the hydrogen ions in acetonitrile-water mixtures in order to clarify the acid-base behaviour of the solutes and (c) to study the correlation of pH(s) values with the solvatochromic parameters n*,a and fi for the acetonitrile-water mixtures by the LSER method with a view to determining the pH(s) values of the buffer reference solution studied in any of the binary solvent acetonitrile-water mixtures.The equations obtained allow calculation of the pH(s) values of the standard 3288 buffer solutions in any acetonitrile-water mixture up to 70 wt.% acetonitrile and thus permit the standardization of potentiometric sensors in these mixtures.Experimental Apparatus The emf values of the potentiometric cell were measured with a CRISON 2002 potentiometer (k0.l mV) using a Radi-ometer G202C glass electrode and a reference Ag/AgCl elec- trode prepared according to the electrolytic method17 and directly inmmersed in the solution to avoid the residual liquid-junction potentials. This electrode system gave stable and reproducible potentials within 5 min. The glass electrode was stored in water when not in use and soaked for 15-20 min in acetonitrile-water mixture before potentiometric mea- surements.The E" values used here are the average of at least 15 standardizations. l7 The standardization of the electrode system was carried out each time the solvent medium or elec- trodes were changed, and the constancy of E" values was assured by continual surveillance by means of periodic cali- brations. The cell was thermostatted externally at 25 0.1 "C. All the potentiometric assembly was automati- cally controlled by a Stronger AT microcomputer. Reagents Analytical reagent grade chemicals were used, unless other- wise indicated. All of the solutions were prepared by mixing doubly dis- tilled, freshly boiled water, the conductivity of which did not exceed 0.05 pS cm-',and acetonitrile (Merck, chromatog- raphy grade). The concentrations of the standard reference solutions were chosen as recommended by IUPAC.3 Primary standards (potassium hydrogen tartrate, potassium dihydro- gen citrate, disodium hydrogen phosphate, potassium dihy- drogen phosphate and potassium hydrogen phthalate) were Merck reagents for preparation of pH standard buffer solu- tions according to DIN 19266.Chemicals were dried at 110"C before use. Anhydrous sodium acetate and anhydrous acetic acid were Carlo Erba Reagents, RPE-ACS and RS grade, respectively. Stock 0.1 mol I-' potassium hydroxide (Carlo Erba, RPE grade) solutions were prepared with an ion-exchange resin' to avoid carbonation and standardized volumetrically against potassium hydrogen phthalate. Procedures Reference pH values of standard buffer solutions in acetonitrile-water mixtures with 10, 30, 40,50 and 70 wt.% acetonitrile, pH(s), were assigned using the procedure adopted by IUPAC.2 This procedure involves different steps:I8 (i) measurement of the emf of the cell: standard buffer + KCl glass electrode in acetonitrile-water where the reference standard buffer solution contains pot- assium chloride at different and accurately known concentra- tions. The emf, E, of this cell is directly related to the activities of the hydrogen and chloride ions in solution: E = E" + log(aH++I-) (2) where E" is the standard emf of the cell.E" values were deter- mined as in a previous study.17 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 (ii) Determination of the pH values for each concentration of potassium chloride, eel-, examined using the Nernst expression of emf, E, for cell (A): P(aH+ycl-)is a thermodynamic quantity that can be deter- mined in thermodynamically exact terms, but to obtain pH values for the mixed electrolyte in cell (B) it is essential to calculate the molar activity coefficient, pycl-, through an extra-thermodynamic assumption, i.e.a form of the classical Debye-Huckel equation (4) In compliance with IUPAC rule^,^,'^ the value of a, B in eqn. (4) is assigned at T = 298.15 K by an extension of the Bates-Guggenheim convention :2i3 (a, B)T = 1.5[(&wps)/(&spw)];'2 where E is the relative permittivity, p the density and the superscripts W and S refer to pure water and to the approp- riate solvent mixture, respectively.Calculation of pycl-from eqn. (4) requires knowledge of the ionic strength, I, of the standard buffer-KC1 mixed elec- trolyte solutions 1 = I, + cc1-but I is, in turn, a function of the H+ concentration cH+, which is expressed by (7) and of the ionization constants, pK, corresponding to the equilibria involved in the standard buffer solutions in acetonitrile-water mixtures. These pK values were deter-mined previously." Then, calculation of pycl-values proceeds by successive iterations. Initially one takes I = c, + ccl-and obtains pycl-from eqn. (4), for subsequent insertion in eqn. (7) to obtain pcH+and a better value of I by eqn. (6). Thus, one calculates again the pycl-value by eqn.(4), and so on, until constancy of I is obtained. Inserting pycl-in eqn. (3), a distinct pH value is obtained for each concentration ccl-examined. The standard value, pH(s), for standard buffer alone at the fixed concentration recommended as International pH Standards3 can finally be obtained as the intercept at ccl-= 0 from the pH us. ccl-linear regression at each mole fraction, x, of acetonitrile studied. Moreover, pH(s) values were obtained in pure acetonitrile solvent for the three standard buffers that are soluble in this medium. For this purpose, the potentiometric system described previously2' was used and the standard potential of the cell was determined by titration of picric acid solutions in acetonitrile with tetrabutylammonium hydroxide as in pre- vious Results and Discussion The emf, E, of cell B was measured at different concentrations ccl-of KCl added to the constant concentration of each stan- dard buffer in 10, 30, 40,50 and 70 wt.% acetonitrile-water solvent.For each standard buffer solution various series of measurements were performed, for a total of 560 independent measurements over the solvent interval explored. When the acidity function, pH, for each buffer solution was plotted us. ccl-, straight lines of small slope were obtained. Typical J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 regression lines are shown in Fig. 1 for the KH tartrate in the acetonitrile-water mixtures studied. The statistical analysis of variance was applied to the various independent sets of measurements.The variances that we can hope to estimate are s;, the variance within sets of data or series of measurements and si, the variance between sets of data. If applying the F-test for a 5% level of signifi-cance we can conclude that the two variances do not differ significantly, then, for a given primary standard buffer, all the points of every data set belong to the same population and it should be permissible to calculate the total average, pH,, and the standard deviation, s, by fitting all the points together and carrying our least-squares analysis. If this hypothesis is rejected, then most of the error derives from the variability between data series, the pH(s) value is obtained by averaging the different intercepts and the total variance s2 = s; + s; can be calculated.Table 1 shows the pH(s) values determined for the KH tar- trate, KH, citrate, phosphate buffer, KH phthalate and acetate standard buffer solutions in 10, 30, 40, 50 and 70 wt.% acetonitrile-water mixtures and the respective standard deviations, s, together with standard pH(s) values reported in water24 that can be used for standardization of poten-tiometric sensors in these solvents. Table 1 also shows the pH(s) values in neat acetonitrile obtained in the present work. pH(s) values of acetate and phosphate buffers cannot be obtained in 70 wt.% acetonitrile-water mixtures or in neat acetonitrile, since these substances are not soluble in these media. Although the variation of the pH(s) values obtained in acetronitrile-water mixtures with xANis approximately linear in the mixtures studied (Fig.2) these pH(s) are lower than the expected values considering the high pH(s) in the neat solvent acetonitrile, (Table 1). Thus, preferential solvation by one of 5.0 ---50%h$ 4.5-$40% P --30% 4.0-+ : : : : : e + +-t-t-t+--c+10% 3.5-I I I I I 2 -0.05 0.05 0.15 0.25 0.35 0.45 0.55 XAN Fig. 2 pH(s) as a function of the mole fraction of acetonitrile, xANin the acetonitrile-water mixtures. V,KH tartrate; +, KH, citrate; +, KH phthalate; 0,acetate buffer; A, phosphate buffer. the solvents is expected. If a solute interacts with one solvent more strongly than with the other, the solute will be prefer- entially solvated by the former.Preferential solvation in acetonitrile-water mixtures produce lower pH(s) than expected if the 'preferred' solvent is water. As a consequence of the mixture nature of the Dimroth and Richardt solvatochromic parameter, E,(30), which is sen- sitive to many solvent-solvent interactions, it has been used to study preferential solvation in many binary mix-tUres.4,8,15,25 The normalized EF values for acetonitrile-water mixtures ' are given in Table 2. Krygowsky et3 6v2 7 ~1.'~studied the contribution of the transition energy of pyri- dinium betaine due to specific interactions (E;) in aqueous binary solvent mixtures. A plot of EY values of the acetonitrile-water mixture us.the mole fraction of water is shown in Fig.3. For these mixtures there is preferential sol- vation of betaine by acetonitrile in the water-rich region. The strong self-association interactions between the water mol- ecules makes them less available for solvation. However, as more dipolar aprotic solvent is gradually added it breaks the self-associated structure of water and the resulting 'free' water molecules then interact with the solute.4*" Thus, the solute is preferentially solvated by water in the acetonitrile- Table 2 Solvatochromic parameter values for the acetonitrile-water mixtures studied acetonit rile X (wt."/,) EY 7r* a B O.oo00 0 1.00 1.14 1.13 0.47 0.0465 10 0.93 1.10 1.03 0.59 0.1583 30 0.84 1.01 0.92 0.61 0.2264 40 0.81 0.97 0.91 0.61 0.3051 50 0.79 0.92 0.90 0.61 0.5059 70 0.76 0.84 0.89 0.59 Table 1 pH(s) values for standard reference solutions in various acetonitrile-water mixtures together with their standard deviations (in parentheses) at 298.15 K acetonitrile (wt.Yo) reference standard 0" 10 30 40 50 70 100 KH tartrate 3.557 3.802 (0.006) 4.325 (0.005) 4.570 (0.005) 4.852 (0.004) 5.723 (0.006) 17.79 (0.09) KH, citrate 3.776 3.994 (0.007) 4.470 (0.008) 4.702 (0.006) 4.995 (0.003) 5.610 (0.005) 16.48 (0.09) KH phthalate 4.008 4.318 (0.005) 5.015 (0.004) 5.346 (0.004) 5.644 (0.004) 6.428 (0.005) 16.82 (0.07) acetate buffer 4.644 4.898 (0.005) 5.532 (0.004) 5.875 (0.003) 6.275 (0.004) --phosphate buffer 6.865 7.149 (0.007) 7.604 (0.003) 7.667 (0.004) 8.002 (0.005) --Ref.24. 3290 1.o 0.8 0.6 zi-LL1 0.4 0.2 0.0 -1 0.1 0.3 0.5 0.7 0.9 1.1 xw Fig. 3 ,!$ values of acetonitrile-water mixtures us. the mole fraction of water, x, rich region. The transition occurs at an intermediate solvent composition around xANz 0.35 or ca. 55 wt.% acetonitrile. The solvation of the betaine used to establish the E, scale of solvent polarity may, however, be different from that of inor- ganic ions since it is large and hydrophobic. Moreover, in water-acetonitrile mixtures, the specific solvation of iodide and chloride ions has been studied by UV and NMR, respec- ti~ely.~~,~~The results showed that iodide and chloride are preferentially solvated by water in acetonitrile-water solvent mixtures.In a parallel manner to the E, scale, halide ions, n-alkylpyridinium iodides' or several alkyl halide^,^ the solutes in the buffer solutions studied are likely to undergo preferential solvation in acetonitrile-water mixtures. The composition of the immediate surroundings of a solute may be different from the composition of the bulk mixture. Preferential solvation is attributable to an excess or a defi- ciency of molecules of one of the solvents in these surround- ing~.~~If the solute has no preference between the solvent molecules, the solvent composition in the cybotactic zone, in the immediate neighbourhood of the solute, is the same as in the bulk. For such cases: where pH(s), and pH(s), represent the standard pH(s) values in solvents 1 and 2, respectively. The deviation from the ideal dependence on the composi- tion of the mixture, pH(s) us.x1 plot (Fig. 4) indicates that the solvent composition in the neighbourhood of the solute may be different from that in the bulk. Fig. 4 shows the pH(s) values as a function of xl, the bulk mole fraction of water. pH(s) us. x1 plots are not linear but a dotted line indicating linearity over the entire range (henceforth referred to as the ideal line) is also shown. The results from the QLQC method at 25 "Cfor the prefer- ential solvation by water, a,, and for the local mole fraction of water, xi, around a hydrogen ion are shown in Table 3 and plotted in Fig. 5 and 6, respectively, as functions of the composition of the binary mixtures of water and acetonitrile. The preferential solvation of hydrogen ions by water is posi- tive, i.e. water molecules show a greater tendency to be in the immediate vicinity of a given hydrogen ion than acetonitrile molecules.This preference is a maximum at x, z 0.25. The results, in terms of preferential solvation by water, d,, Fig. 5 and Table 3, are similar to those also obtained for chloride ions in aqueous acetonitrile" with a similar At G*(Cl-, W +AN)kJ mol-' = 42.1.31 Therefore, as the preferential solvation of hydrogen ions by water is positive, the standard pH(s) values in these mixtures are more similar to the standard pH(s) values in water than to the pH(s) in acetonitrile (Fig.4). This is different from com- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 \. :\.,.'\.' .\. .\.'; \ .'<'.. \*' ,\. c.\.\ \ i. \ I I I I 1 I I I I 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 xw Fig. 4 pH(s) us. mole fraction of water, x,, in acetonitrile-water mixtures. +, KH tartrate; 0,KH, citrate; f, KH phthalate. The dashed straight lines correspond to the ideal variation of the pH(s) values for the three buffer solutions. Table 3 Results from the QLQC method at 25 "C 0.05 0.95 0.03 0.98 0.10 0.90 0.07 0.97 0.15 0.85 0.10 0.95 0.20 0.80 0.14 0.94 0.25 0.75 0.17 0.92 0.30 0.70 0.21 0.91 0.35 0.65 0.24 0.89 0.40 0.60 0.27 0.87 0.45 0.55 0.30 0.85 0.50 0.50 0.33 0.83 0.55 0.45 0.35 0.80 0.60 0.40 0.37 0.77 0.65 0.35 0.39 0.74 0.70 0.30 0.40 0.70 0.75 0.25 0.40 0.65 0.80 0.20 0.39 0.59 0.85 0.15 0.37 0.52 0.90 0.10 0.31 0.41 0.9.5 0.05 0.21 0.26 0.0 0.1 0.2 0.3 0.4 0:5 0:s 0.7 0:8 0:9 1.0 xw Fig.5 Preferential solvation of hydrogen ions by water in acetonitrile-water mixtures, a,, as a function of the solvent composi- tion. The dashed straight line corresponds to total preference for water. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 f / / , xw Fig. 6 Local mole fraction of water, xk, near hydrogen ion as a function of its bulk mole fraction, xw, according to the QLQC method positions close to the acetonitrile pure where x, < 0.25 and the preferential solvation by water decreases quickly, Fig.5. The pH(s) values obtained could be explained in terms of the structural features of the acetonitrile-water mixtures. The structure of mixtures of water and acetonitrile was explored by Marcus and Migron" using the QLQC and the inverse Kirkwood-Buff integral (IKBI) methods. These authors con- cluded that these methods indicate strong microheterogeneity in a middle range of compositions in mixtures of water and acetonitrile, i.e. preference for neighbours of the same kind, which extends over several concentric shells around a given molecule. Thus, there is a preference of a given water mol- ecule for water molecules rather than acetonitrile molecules.This preference is a maximum at xAN x 0.75. The same applies to the preference of acetonitrile molecules to be in the vicinity of a given acetonitrile molecule. Thus, the plot of pH(s) values us. x,, Fig. 4, can be explained by taking into account that in acetonitrile-water mixtures there are three regions.7.1 1.33.34 On the water-rich side there is a region in which the water structure remains more or less intact and the acetonitrile molecules gradually occupy the cavities between water molecules without disrupting the water struct~re.~' The limit of xANbeyond which the acetonitrile can no longer be accommodated within the cavities of the structure of water is CQ. 0.15." The variation of the pH(s) values for KH phthal- ate over the whole composition range is shown in Fig.4, where the dotted line represents the expected varation of pH(s) between xAN = 0.5 and pure acetonitrile solvent. pH(s) values vary linearly with x, with an inflection point at xAN x 0.15. The slope of the pH(s) us. x, plot is greater in the water- rich region, xAN < 0.15, than in the regions where water- acetonitrile mixtures show microheterogeneity, which is in Table 4 Linear solvation energy relationships for pH(s) values reference standard LSER Y KH tartrate pH(s) = 9.19 -8.97n* + 3.35a + 1.718 0.994 KH, citrate pH(s) = 9.34 -7.16n* + 1.89a + 0.978 0.999 KH phthalate pH(s) = 12.79 -8.25n* + 0.49a + 0.158 0.999 acetate buffer pH(s) = 12.87 -8.43n* + l.lla + 0.268 0.999 phosphate buffer pH(s) = 11.18 -4.92n* + 0.62~~+ 1.258 0.998 329 1 accordance with the low values of 6, in acetonitrile-water mixtures with xAN < 0.15, Fig.5. In the range 0.15 d xAN < 0.75 there are clusters of molecules of the same kind sur- rounded by regions where molecules of the two kinds are near each other. In this middle range of compositions prefer- ential solvation of hydrogen ions by water is high, Fig. 5, this could explain the low slope of the linear variations of the pH(s) us. x, plots (Fig. 4). At xAN 2 0.75 the number of water clusters is low, and water-acetonitrile interactions that could be discounted in the middle range now become important. One may consider this region as 6, decreases, Fig. 5, and then a concave variation of the pH(s) us.x, plot as represent- ed by the dotted line for KH phthalate in Fig. 4 could be expected. The boundaries of the regions are, of course, not sharp.' ' On the other hand, it is not self-evident that solvatochro- mic parameters would be valid to represent generalized solutes in binary solvent mixtures with regard to the proper- ties they are supposed to measure. Preferential solvation in such mixtures may interfere more seriously with the ability of indicators to represent generalized solutes than in the case of single solvents. Progress has been made' 1*36 and although this problem has not been solved unequivocally, these investi- gations, provide significant evidence that the solvatochromic parameters seem to have general validity.It is therefore, of interest to examine the LSER which explain any solute pro- perty varying with solvent composition as a linear com-bination of the microscopic parameters of the solvent responsible. The Kamlet-Taft ' expression states : XYZ = (XYZ), + aa + b/?+ s7r* (21) where a, /? and n* are the microscopic parameters previously described, XYZ is the solute property, XYZ, the value of this property for the same solute in a hypothetical solvent for which a = /? = n* = 0 and a, b and c are the susceptibilities to changes in a, /3 and n*,respectively, of the solute property studied. This equation can include additional terms or some of its terms can become equal to zero, depending on the pro- perty of the solute to be described.37 Values of the Kamlet- Taft solvatochromic parameters n*,' 1.38 a' 1*39 and /?"7'5 for acetonitrile-water mixtures over the entire range of composi- tion are known.Table 2 gives the relevant solvatochromic parameter values for the mixtures studied. Several attempts were made to find the best form of the Kamlet-Taft equation to describe the variation of stan-dard pH(s) values in acetonitrile-water mixtures. Multiple regression analysis was applied to our pH(s) data. All pos- sible combinations of solvatochromic parameters were checked. The best fit was obtained when the three solvatoch- romic parameters a, #? and 7r* are used, yielding the general equations shown in Table 4. From a practical point of view these equations enable us to know the standard pH(s) values of the buffer reference solutions studied in any binary solvent acetonitrile-water mixture up to 70 wt.% acetonitrile, and thus permit the standardization of potentiometric sensors in these mixtures.References 1 T. Mussini and F. Mazza, Electrochim. Acta, 1987,32, 855. 2 S. Rondinini, P. R. Mussini and T. Mussini, Pure Appl. Chem., 1987,59,1549. 3 A. K. Covington, R. G. Bates and R. A. Durst, Pure Appl. Chem., 1985,57, 531. 4 J. G. Dawber, J. Ward and R. A. Williams, J. Chem. SOC., Faraday Trans. I, 1988,84,713. 5 P. Chatterjee, A. K. Laha and S. Bagchi, J. Chem. SOC.,Faraday Trans., 1992,88, 1675. 6 N. Papadopoulus and A. Avranas, J. Solution Chem., 1991, 20, 293. 3292 J. CHEM. SOC.FARADAY TRANS., 1994, VOL. 90 7 H. Kovacs and A. Laaksonen, J. Am. Chem. SOC., 1991, 113, 5596. 24 25 R. G. Bates, Crit.Rev. Anal. Chem., 1981, 10, 247. J. R. Haak and J. B. F. N. Engberts, Recl. Trau. Chim. Pays Bas, 8 9 10 11 12 13 14 15 J. G. Dawber, J. Chem. SOC., Faraday Trans., 1990,86287. Y. Marcus, Aust. J. Chem., 1983,36, 1719. Y. Marcus, J. Chem. SOC., Faruday Trans. I, 1988,84,1465. Y. Marcus and Y.Migron, J. Phys. Chem., 1991,%, 400. J. Kamlet, R. M. Doherty, M. H. Abraham, Y.Marcus and R. W. Taft, J. Phys. Chem., 1988,92,5244. M. J. Kamlet, J. L. M. Abboud, M. H.Abraham and R. W. Taft, J. Org. Chem., 1983,48,2877. K. Dimroth and C. Reichardt, 2. Anal. Chem., 1966,215,344. T. M. Krygowski, P. K. Wrona, U. Zielkowska and C. Rei- chardt, Tetrahedron, 1985,41,4519. 26 27 28 29 30 31 32 33 1986,105,307.J. G. Dorsey and B. P. Johnson, Chimicaoggi, 1986,23. S.Balakrishnan and A. J. Easteal, Aust. J. Chem., 1981,34,943. M. C. R. Symons and S. E. Jackson, J. Chem. SOC., Faruday Trans. 1,1979,75,1919. C. H.Langford and T. R. Stengle, J. Am. Chem. SOC., 1969, 91, 4014. Y.Marcus, J. Chem. SOC., Faruday Trans. 1,1989,85,381. Y.Marcus, Pure Appl. Chem., 1986,58,1721. J. Barbosa, S. Buti and V. Sanz-Nebot, Talanta. 1994,41,825. M. I. Davis, Thermochim. Acta, 77,421. 16 P. Chatterjee and S. Bagchi, J. Chem. SOC., Faruday Trans., 1991,87, 587. 34 35 A. J. Easteal, Aust. J. Chem., 1979,32, 1379. A. J. Easteal and L. A. J. Woolf, J. Chem. Thermodyn., 1988, 20, 17 18 19 20 J. Barbosa and V. Sanz-Nebot, Anal. Chim. Acta, 1991,244,183. J. Barbosa and V. Sanz-Nebot, Mikrochim. Acta, 1994, 111, in the press. T. Mussini, A. K. Covington, P. Longhi and S. Rondinini, Pure Appl. Chem., 1985,57,865. J. Barbosa, J. L. Beltran and V. Sanz-Nebot, Anal. Chim. Acta, 1994,288,271. 36 37 38 39 693. Y.Migron and Y.Marcus, J. Chem. SOC., Faraday Trans., 1991, 87, 1339. M. J. Kamlet and R. W. Taft, Acta Chem. Scand., Ser B, 1985, 39, 611. W. J. Cheong and P. Carr. Anal. Chem., 1988,60,820. J. H. Park, M. D. Jan& D. S.Kim and P.W. Carr. J. Chromato- 21 22 J. Barbosa and V. Sanz-Nebot, Talanta, 1989,36,837. J. Barbosa, V. Sanz-Nebot and M. E. Torrero, Talanta, 1991,38, gr., 1990,513, 107. 425. 23 J. Barbosa, Encyclopedia of Analytical Science, in the press. Paper 4/01396D;Received 9th March, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003287

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3293-3299 3293 Effect of Solvent on the Reaction of Coordination Complexes Part 19.T-Base Hydrolysisof (ap)-(0-Methoxy benzoato)(tetraethylenepentamine)-cobalt(iii) in Aquo-organic Solvent Media Achyuta N. Acharya and Anadi C. Dash* Department of Chemistry, Utkal University, Bhubanes war-751 004, India The kinetics of base hydrolysis of (afl)-(o-methoxy benzoato)(tetren)cobalt(itt) have been investigated in aquo- organic solvent media [O-70% (v/v) cosolvents] at 10 d t/"C d 40 (I = 0.01 mol dm-3) using methanol, ethanol, propan-1-01, propan-2-01, butan-1-01, tert-butyl alcohol, ethylene glycol, 2-methoxyethanol, acetone, acetonitrile, 1,4-dioxane and dimethyl sulfoxide as cosolvents. The second-order base hydrolysis rate constant increased non-linearly with increasing mole fraction (XOrs)of all cosolvents, except for the ethylene glycol-water system ; ethylene glycol had a rate-retarding effect.The transfer Gibbs energy of the transition state (TS) relative to that of the initial state (IS),for transfer of species from water to mixed solvent varied non-linearly with 0;'andX,,, , reflecting the individuality of the cosolvents and thereby suggesting that the relative stabilities of the transition state and the initial state were governed by the preferential solvation effect. The solvent stabilisation of the initial state and the transition state has been assessed for the methanol-water and ethanol-water systems by combining the solubility data of the dithionate salt of the complex with the transfer Gibbs energy data for S,0,2-.The thermodynamic parameters (AH* and ASt)were sensitive to the structural changes in the bulk soIven t phase. The importance of hydrophobic effects,'V2 preferential Experimental s~lvation~-~and solvent structural effects7-' on the kinetics (agS)-(o-Methoxybenzoato)(tetren)cobalt(rn) diperchlorateand energetics of the ligand-substitution reaction of was prepared as described earlier.17 The analytical data for transition-metal complexes has been emphasized in several the complex (C, H, N and Co) agreed with those predicted by recent studies.' I-' Recently, we reported the base hydrolysis its molecular formula. (aBS)-(0-MethoxybenzoatoXtetren)inof (a~S)-(salicylato~tetraethylenepentamine)cobalt(~~~) cobalt(II1) dithionate was prepared by repeated rec-mixed-solvent media4 using methanol and dimethyl sulfoxide rystallization of the perchlorate complex salt with sodium as cosolvents.It was evident that solvation of the initial state di thiona te, washed successively with water, absolute alcohol, and the transition state was strongly influenced by these diethyl ether and stored over fused calcium chloride. Calc. for cosolvents. However, (salicylato)(tetren)cobalt(IIr) (where [Co(tetren)0,CC,H,(OCH3)]S206 :Co, 10.6. Found : Co,tetren = tetraethylenepentamine) was capable of generating a 10.4. The dithionate complex exhibits AJnm (&/dm3 mol- ' reactive amid-imido conjugate base via intramolecular cm-')at 490 (199) in 0.1 mol dm-3 HClO, which closely acid-base equilibrium, -0--* -NH-7HO.* .N--.16 agreed with the data of the corresponding perchlorateSuch an effect will be absent in the corresponding o-complex.' methoxybenzoato complex. Hence the solvent perturbation Organic solvents? (analytical grade) were distilled after effects of the tetren ligand and the aromatic moiety on drying over 4A molecular sieve. 1,4-Dioxane and 2-the base hydrolysis of (apS)-(o-methoxybenzoato)(tetren)-methoxyethanol were made peroxide-free by repeatedlycobalt(II1) (I) can be examined much more clearly. With this refluxing them over sodium hydroxide and metallic sodium, aim we made a thorough study of this reaction in aquo- respectively, and distilling before use. Density and boiling organic solvent media which differed in their structural, elec- points were taken as a check of the purity of the solvents.trostatic and hydrophobic interaction propensities. Fresh solvent mixtures were prepared by mixing a known volume of organic solvent with water before recording any kinetic run. Doubly distilled water was used; the second distillation was made from alkaline KMnO, using an all- glass distillation still. All other chemicals used were of ana- lytical grade. Solubilities of the dithionate complex in MeOH-H,O and EtOH-H,O were determined by equilibrating an excess of the solid complex in the acidified mixed solvent ([HClOJ x lo-, mol dm-3) for 3-4 h in stoppered flasks, thermostatted at 25.0 +_ 0.2"C. Samples of the solutions were withdrawn b""' and the concentrations were determined spectrophotometri- cally at 490 nm. All the spectral measurements were made I using a Jasco 7800 UV-VIS recording spectrophotometer.(cr/?S)-(o-methoxybenzoato)(tetren)cobalt(ltl) ion ~~ ~ ~ Part 18: A. C. Dash and G. C. Pradham, Znd.J. Chem., Sect. A, 7 Abbreviations : methanol, MeOH ;ethanol, EtOH ; propan-1-01, 1994, 33, 661. This work was partly presented at the 23rd IUPAC PrOH ; propan-2-01, Pr'OH ; butan-1-01, BuOH ; tert-butyl alcohol, Conference on Solution Chemistry, held at the University of Leices- Bu'OH ; ethylene glycol, EG; 2-methoxyethanol, ME; acetone, AC; ter, UK, 16-23 August, 1993. acetonitrile, AN; dimethyl sulfoxide, DMSO; 1,4-dioxane, D. Kinetics The base hydrolysis of the title complex (perchlorate form) was studied under pseudo-first-order conditions ([OH-]~/[complex]~2 25) at 320 nm using an automated Hi-Tech (UK) SF 51 stopped-flow spectrophotometer con- trolled by an Apple IIGS PC.The decay of the complex was found to be single-exponential for any run under the experi- mental conditions, even after extending the reaction over an expanded timescale (0.01-50 s). The rate measurements were made at [NaOH], = 0.01 mol dm-3 and 10.0 < t/"C < 40.0, with At/"C 2 20 (at least four temperatures) for any solvent composition. The other experimental details were the same as described earlier.' The pseudo-first-order rate constant and its standard deviation for any run was calculated from at least seven replicate measurements. Results and Discussion The rate data (kobs) in the fully aqueous medium strictly obey a second-order relationship, kobs= ko,[OH -]T [see eqn.(I)] (tetren)CO02CC6H,(o-oMe2+ + OH-kon-(tetren)CoOH'+ + -O,CC,H,(o-OMe) (I) in the range 0.010 < [OH-],/mol dm-3 < 0.10 at 20.0 Q t/ "C < 40.0 (I= 0.10 mol dm-3) (see Table 1). The activation parameters (AHS and ASS)agree well with those reported by us earlier' '(see Table 1) at I = 1.0 mol dm -'. The large posi- tive value of AS: is consistent with the S,1 conjugate-base (CB) mechanism '' involving dissociatively activated conju- gate base of the complex generated by the deprotonation of the N-H group of the coordinated tetren. Variation of Rate Constant with Solvent Composition The second-order base hydrolysis rate constants, koH ,in dif- ferent aquo-organic solvent media at I = 0.01 mol dm-' are collected in Table 2.All aquo-organic solvent systems except EG-H,O exerted a marked accelerating effect over the com- position range studied. Since ko, is subject to ionic strength effects, its correction to I = 0 (k&) was made by using the relationship in eqn. (1) log (koH)= log kgH + 22, 2B AI1',/( 1 + (1) where Z, and zB denote the charges of the reacting species, A = 1.8246 x 106/(DsT)3/2,B = (50.29 x 108)/(D,T)lI2,D,is the bulk relative permitti~ity'~ and Q was chosen to be 5 A. log kgH at 25 "C increased non-linearly (except for MeOH and DMSO) with increasing mole fraction of the organic solvent, Xorg (see Fig.1); such plots are linear for MeOH-H,O and DMSO-H,O media. In contrast, non-J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 4.6 4.0 3.4 I 00 Y 0-2.8 2.2 Y Y n -I I I I I1.6 0.10 0.20 0.30 0.40 0.50 0. 0 XO,, Fig. 1 log k&., us. X, for different aquo-organic solvent media at 25 "C: (a) EG, (b)MeOH), (c) DMSO, (6)EtOH), (e) AN, (f)PrOH, (9)Pr'OH, (h) AC, (i) ME, (j) Bu'OH, (k)D the EG-H,O system (see Fig. 1). The rate effect is modestly sensitive to the nature of the cosolvents, being maximum for 1,4-dioxane. The variation of log kgH with 0,' is generally non-linear ; linearity is observed for DMSO-H,O. Plots of log kg, us. E; and Grunwald Winstein's solvent ionising parameter, Y for which E; (MeOH-H,O, Pr'OH-H,O, AC-H,O, DMSO-H,O, AN-H,O),' and Y (MeOH, Pr'OH, Bu'OH and AC)2' data are available, are not linear.These facts reflect the inadequacy of the structureless dielectric con- tinuum model, which depicts the solvent and solute as uni- formly charged rigid spheres. The importance of hydrogen bonding and other non-electrostatic medium effects in con- trolling the reactivity of the substrate is apparent. The observed solvent effect may be interpreted in terms of the conjugate base mechanism: kca (CB+)' -H2O Products (11) linear rate retardation with increasing X,, was observed for fast Table 1 Rate data for base hydrolysis of I in aqueous medium" [OH-]/10-2 mol dmT3 20.0 * 0.01"C 25.0 f0.1 "C 35.0 f0.1 "C 40.0_+ 0.1 "C 1.o 0.22 f0.01 0.44 f0.02 1.54 f0.09 2.89 f 0.09 2.0 0.46 f0.01 0.86 f0.02 2.95 f 0.07 5.38 0.27 5.0 1.20 * 0.04 2.28 f0.01 7.85 k0.36 13.2 & 0.7 7.50 1.90 k 0.08 3.60 f0.14 11.9 f0.2 23.3 f1.2 10.0 2.54 f0.10 4.99 0.14 17.4 f1.2 30.6 f2.1 k,,/dm3 mol s-l 23.4 f0.6 44.9 k1.3 155 f3 284 f7 AH:fiJ mol -' 92.5 f 0.3(93 f2)bAS/J K-' mol-' 97 f l(90 f5)b [complex], = (3-4) x I = 0.1 mol dm-3, 1 = 320 nm.I = 1.0 mol dm-3.'7 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3295 Table 2 Rate data for base hydrolysis of I in mixed-solvent media with various amounts of organic solvent (in vol.%) at different temperatures k,,"/dm mol -s -t/T 5.2 vol.% 10 vol.% 20 vol.% 30 vol.% 40 vol.% 50 VO~.%O 60 vol.% 70 VO~.% MeOH 20.0 56f 1 60f 1 67 f2 75 f2 87 f1 105 f2 136 k 3 188 f7 25.0 108 f3 117 f 4 134 f3 159 f2 188 & 7 230 k 2 301 f13 386 f13 35.0 393 f9 425 f8 489 f4 632 k 2 814 f8 lo00 f20 1295 f34 1733 f98 40.0 676 f24 749 f20 958 f 15 1177 k 32 1495 f 44 1903 k 76 2769 k 87 3618 f162 (0.024)b (0.047) (0.099) (0.159) (0.228) (0.306) (0.398) (0.507) EtOH 20.0 60f 1 73 f 1 102 f2 149 f3 217 f9 318 k 12 448f 15 680 f16 25.0 114 f4 137 f3 197 f2 309f7 440f9 605 f14 950 f11 1412 f30 35.0 387 f7 475 f4 740f 17 1106 f27 1629 f50 2127 k 80 3357 f 53 5100 f 118 40.0 724f 11 900f9 1291 f 22 1996 f44 3132 f68 4180 90 6043 f210 9060 f300 (0.017)b (0.033) (0.07 1) (0.116) (0.170) (0.235) (0.3 15) (0.416) PrOH 20.0 54f 1 67 f2 95 f1 136 f7 202 f 5 265 k6 362 f 15 514 f5 25.0 102 _+ 1 125 k 8 190 f3 285 _+ 4 404f6 541 f16 754 f 14 985 f20 35.0 356 f 8 446f 10 659 f 13 1029 f15 1494 f34 1916 f40 2549 f 86 4085 k 100 40.0 639 f18 836 f6 1261 f18 1918 _+ 56 2732 f 39 3684 f81 5005 f 180 6700 f320 (0.01 3)b (0.026) (0.057) (0.093) (0.127) (0.179) (0.246) (0.336) Pr'OH 20.0 62 f2 82 f2 128 f3 218 _+ 3 356 f9 483 f 11 900f 11 1533 f 60 25.0 120 f3 152 f6 244f5 426 f 10 701 & 10 1111 f50 1818 f50 3079 f 174 35.0 396 f5 524 k 7 891 f5 1467 k31 2401 f39 3658 f60 5795 f56 9504 f294 40.0 733 f17 1OOOf 18 1580 f19 2417 f52 3965 f 103 6296 f318 loo00 f500 16500 f900 (0.01 3)b (0.025) (0.055) (0.091) (0.135) (0.190) (0.260) (0.3 5 2) Bu'OH' 20.0 62f 1 82 f1 134 f3 178 f 3 344f7 529 f11 858 f5 1429 f18 25.0 121 f 3 154 f3 251 f2 343 f5 631 f7 1014 f 21 1646 f35 2878 _+ 76 35.0 408f8 536 f5 883 f4 1197 f13 2322 f112 3346 f88 4970 i-39 8660 f300 40.0 759 f14 990f 15 1635 f31 2101 f51 4004 f32 4963 f220 9591 f476 15330 f1032 (0.010)b (0.021) (0.045) (0.075) (0.1 12) (0.159) (0.221) (0.306) ME 20.0 56 & 1 63 f2 73 f1 97 f2 130 f3 196 f1 258 f6 427 f12 25.0 109f6 123 f3 151 f4 206f9 283 f5 395 f9 568 f6 904 k 24 35.0 346 f16 418 f9 544 f22 703 f23 1064 f28 1493 f31 2141 k 83 3444 f102 40.0 664 f 16 794 f21 1047 f28 1442 f 26 2104 f41 2895 f106 4498 f69 7259 f198 (0.012)b (0.025) (0.054) (0.089) (0.132) (0.185) (0.254) (0.346) EG 20.0 39 f1 34 f1 24 f1 19 f1 16f 1 12 k 1 10.3 f0.3 9.9 f0.4 25.0 76f 1 66 f2 49 f 1 38 f 1 32f 1 28 f1 26 & 2 24 f 1 35.0 267 f1 246f6 189 f5 157 f2 136 k 3 125 f2 118 * 2 115 f6 40.0 499 f8 466 f 11 377 f3 330 f8 291 i-8 265 f5 257 f5 248 f4 (0.017)b (0.035) (0.075) (0.121) (0.177) (0.243) (0.3 25) (0.428) AC 20.0 64f2 77 f1 129 f3 231 f 4 446 f10 855 i-11 1702 f42 3232 f142 25.0 123 f3 159 f3 252 f 10 402 f10 820 f20 1538 f93 3273 f84 6200 f200 35.0 413 f8 531 f 18 937 f24 1657 f29 3024 f122 5608 f335 10800 k 300 19300 f1300 -40.0 761 f29 976 f 25 1685 f37 2980 f67 5597 f192 9988 k 583 19200 f1400 (0.01 3)b (0.026) (0.057) (0.094) (0.139) (0.195) (0.267) (0.361) AN 20.0 56+ 1 65 f1 97 f2 156 f2 262 f6 448 5 819 f6 1657 f39 25.0 104f2 128 f 2 189 f4 316 8 537 f5 926 f13 1720 44 3264 160 35.0 374 f 8 442 f 14 717 f13 1136 f38 2045 f27 3527 f95 6275 f155 12400 f600 40.0 659 f11 807 f 16 1312 & 14 2185 f89 3838 f118 6510 _+ 158 12300 k 700 25800 f1100 (0.01 8)b (0.036) (0.079) (0.127) (0.185) (0.254) (0.3 38) (0.442) DMSO 10.0 16f 1 17 f1 18 f1 27 _+ 1 33 f1 47 * 3 20.0 58 f1 65 f1 83 f2 119 f4 159 f5 261 f7 25.0 118 f5 137 f2 165 f5 214 k9 320 & 9 463 f23 35.0 387 k8 422 f 14 547 f4 712 f31 1221 f31 2116 f 74 40.0 765 + 18 860f 15 1066 f29 1421 f69 2287 f79 3754 f141 (0.014)b (0.027) (0.059) (0.68) (0.124) (0.201) 3296 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Continued t/OC 5.2 VO~.% 10 VO~.% 20 vol.O/o 30 VO~.% 40 vol.O/o 50 vol.% 60 vol.% 70 vol.% D 15.0 33 f1 43 f1 74 f2 146 f4 278 f8 420 f3 810 f10 1491 f32 20.0 62 f2 84 + 2 151 f3 304f3 527 f12 lo00 f50[235 f7Id [397 f31' 1878 f43 [872 f131" 3346 f122 30.0 35.0 25.0 228 + 6 419 f5 123 f2 300 f5 556 + 11 163 +3 513 f 11 1027 f25 285f2 1066 f 22 2022 k107 581 f6 2084 f49 3835 f92 1050 f27 3272 f35 5900 f265 1900 f58 6849 250 10300 f400 3735 f59 6100 f140 12100 f 400 16700 f1200 (0.01 l)b (0.023) (0.050) (0.083) (0.123) [11400 f5001' (0.173) C16600 f14001' (0.239) [32700 f 12001' (0.328) a k;,/dm3 mol-' s-l = 52.4 f 1.2, 98.2 f 1.7, 338 f 7 and 603 f15 at 20.0, 25.0, 35.0 and 4O.O"C, respectively, I = 0.01 mol dm-3, 1 = 320 nm.Values in parentheses denote mole fraction of the organic component. 'For BuOH, kOH/dm3 mol-' s-' (vol.% BuOH): 54 1 (2.0), 59 f 1 (4.0), 63 f1 (6.0), 67 f 2 (8.0) at 20.0"C; 101 f1 (2.0), 109 f3 (4.0), 118 + 1 (6.0), 131 f4 (8.0) at 250°C; 340 f12 (2.0),377 f10 (4.0), 409 f8 (6.0), 438 f11 (8.0) at 35.0"C;659 f9 (2.0), 704 f 14 (4.0), 778 f16 (6.0), 844 f 8 (8.0) at 40°C.'At 10°C. At 40°C. Table 3 Calculated values of [A,G"(CB+)* -A, G"(C2+)],,,JkJ mol-for different aquo-organic solvent systems at 25 "C" [A, G"(CB+): -Al G"(C2')]~,+w,/kJmol-' organicsolvent 10 VOl.Yob 20 vol.% 30 vol.% 40 vol.% 50 ~01.96 60 vol.% 70 vol.% ~ MeOH' -0.50 & 0.10 -0.92 f0.08 -1.57 f0.07 -1.98 & 0.10 -2.15 f0.06 -1.94 f0.12 -1.42 f0.10 -EtOH -0.32 f0.08 -0.44 f0.06 -0.60 f0.08 0.05 f 0.07 1.32 0.08 2.90 f0.06 Pr'OH 0.88 f0.11 0.76 f0.07 0.60 f0.08 0.79 -t 0.08 1.40 f0.12 2.06 f0.08 -Bu'OH -0.33 f0.07 0.16 f0.05 1.61 f0.06 2.11 f0.06 2.68 0.07 --EG -0.07 f 0.09 0.03 f0.09 -0.06 f0.06 -0.19 f0.11 -0.48 f 0.09 -0.18 f0.17 ME 0.17 f0.08 1.27 f0.08 2.13 f0.12 2.55 & 0.07 2.66 f0.08 2.54 f0.06 0.61 f0.08 -D 1.56 f0.07 3.65 f0.05 5.05 f0.06 4.80 f0.08 4.22 f0.09 --AC 2.02 & 0.07 3.96 f0.1 1 6.65 f0.08 9.40 0.08 12.48 f0.16 -DMSO 1.26 f0.07 3.49 f0.09 6.44 f0.12 10.02 f0.09 15.46 f0.13 Based on the TATB assumption.Amount of organic solvent (vol.%) 'Based on the TPTB assumption. a where R = (tetren)cobalt(w) and X = (o-OCH,)C,H,CO, -. where GS(CB+)$ denotes the transition-state Gibbs energy of ko, takes the form : the monopositive conjugate base [(CB'): = (tetran-H) Co2'. * * -O,CC,H,(o-OCH,XTs)], and G,"(C2+) and kOH = kCB KCB (2) G,"(OH-) the standard Gibbs energies of the dipositive sub- where KcB and kc, are as defined in eqn. (11).The activation strate (c2') (Is) and OH-, respectively, in the medium s. Gibbs energy for a given solvent composition s (see Fig 2) Eqn. (4) yields the transfer Gibbs energy of the transition can be given by state (CB+): relative to that of the initial state (C2') of the -complex when transfer of species occurs from water (w) to AGK = G,(CB+)' -G:(C2+) -G,"(OH-) (3) mixed solvent (s). -[A, G"(OH-)I,+w, (4) The relative transfer Gibbs energy term, A[At G*llstw), at zero ionic strength, could be calculated from the rate data using the relationship: A[A, Gx](stwl= RT In (k&/kgH). The differences between the transfer Gibbs energies of the tran- sition state and the initial state of the complex ion, [A, G"(CB+): -A, Go(C2+)](s+w),were calculated [see eqn.(4)] using available data for [At G(OH -)Istw,.22tThese values (see Table 3), based on the TATB-TPTB scale, reflect the effects of differential solvation of the transition and initial states. Note that with increasing hydrophobicity of the alco- hols or the aprotic cosolvents, [A, G"(CB+)t],,tw, tends to be more positive than [At Go(C2+)](stw) except for EG-H,O. Plots of [A, G"(CB+): -At Go(C2')](s+w)vs. 102/D, are gener- ally non-linear and also solvent specific (see Fig. 3); linearity is, however, maintained for DMSO-H,O and AC-H,O products r,.FR~.1 Gibbs energy profile of the reaction, RX2+f= t [A,G"(OH-)],,+,JkJ mol-' for MeOH-H,O is based on the ji:otren~('cjC),CC.h',(o-OCH3)I2+) + OH-+products; AG: = TPTB assumption and the TATB assumption is used for all other I c': i (CB') solvent systems reported here.J. CHEM. iOC. FARADAY TRANS., 1994, VOL. 90 3297 16.0 c 12.0 I-E3 ! 8.0 +N 0,& d I ll* 4.0 m ?.2-u h 0 I-2.5 ' ICL,1 1 2.5 4.0 5.0f 102/D, Fig. 3 [At G"(CB+)'-A, Go(C2+)](,,,JkJ mol-' us. 102/D, at 25 "Cfor different aquo-organic solvent media: (a) MeOH, (b)EG, (c) EtOH), (4Pr'OH), (e) Bu'OH), (f)ME, (9)D, (h)AC, (i)DMSO. ( x ) water. media. The calculated values of [At G"(CB+)* -A, Go(C2+)](2tw)for different solvents at D,= 62.5 (25 "C) (see Fig.3) follow the order:AC > D > ME > Pr'OH > B Bu'OH x EG > EtOH > MeOH. If the electrostatic effect is constant, this result would mean that the solvation of the transition and initial states might be sensitive to solvent structure and hydrophobic interaction effects. c I 7Y -3 !---2.0 --5.5I I I 1 0 20 40 60 0 0.10 0.20 0.30 0.40 0.50 organic solvent (vol.%) XOW Fig. 4 [A,G"(i)],,+,JkJ mol-' us. vol.% organic solvent: (a), (b), Fig. 5 A.HS/kJ mol-' us. Xorl for different aquo-organic solvent MeOH; (c), (4, EtOH; open symbols, initial state; closed symbols, media: A, (a) MeOH, (b)EtOH, (c) PrOH, (a) Pr'OH), (e) Bu'OH, (f)transition state ME. B, (a)AC, (b)AN, (c) DMSO, (4D, (e) EG. J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 170 I A 150 130 110 90 T 9 0 9 0 0 0.10 0.20 0.30 0.40 0.50 0 0.10 0.20 0.30 0.40 0.50 XO rg XOW Fig. 6 AS*/J K-' mol-' us. Xorgfor different aquo-organic solvent media. Notation as for Fig. 5. Assessment of the transfer Gibbs energy of both the initial state and the transition state could be made for the MeOH-H,O and EtOH-H,O systems as follows. The trans- fer Gibbs energies of the dithionate salts of (afiS)-(o-methoxy- benzoatoXte tren)co bal t(m) in MeOH-H ,0 and Et OH-H ,0 media were determined from the measured solubilities of the complex using the relationship given in eqn. (5) A,G"(complex) = 2RT In (Sw/S,)(f~+lf~+) (5) where S, and S, denote the solubilities of the complex salt in aqueous and mixed-solvent media, respectively, and f;+ and f;+ denote the activity coefficients of the complex ion in aqueous and aquo-organic solvent media, respectively.The activity coefficient, f,+ (based on the assumption that f;+ = fi+ = 1 at I = 0), was calculated using the relationship: logf,, = -4AI'/'/(l + BuI"~)with A, B and a as described earlier. The Gibbs energy of transfer of the complex ion, [A, Go(C2+)](s+w), where A, G"(comp1ex) = A, Go(,'+)+ A,G"(S,062-), was calculated using the available data for [A, Go(S2062-)](stw).f-The transfer Gibbs energy of the t [A,G(S,0,2-)]F?;~/kJ mol-': 2.9, 2.1, 4.4, 7.5, 11.3, 14.0 at 10, 20, 30, 40,50, 60 vol.% MeOH (TPTB scale); 2.3, 5.1, 8.4, 12.9 at 20, 40, 60,80 vol.% EtOH (TATB scale).We are grateful to Dr. J. Burgess, University of Leicester, Leicester, UK, for communicating these data to us. complex ion, A, Go(C2+), along with A[A, GSlbtw) and [A,G"(OH-)],+,, [see eqn. (4)] enabled us to dissect the solvent effects associated with the initial and transition states of the complex ion for MeOH-H,O and EtOH-H,O media. Fig. 4 suggests that [A, Go(CB+)t],,+,, < [A, Go(Czf)](stw) is valid for all compositions of MeOH-H,O and EtOH-H,O (<40% v/v EtOH), thereby indicating that the stabilising influence of the mixed solvent is greater on the relatively more expanded transition state than on the initial state. Note that extrema are observed for solvation of both the initial state and the transition state (see Fig. 4), which demonstrates the importance of solvent structure.Thus, this work high- lights the solvent structure mediated solvation of the complex ions of varying charge and hydrophobic Variation of Activation Parameters (A@ and AS) with Solvent Composition The activation parameters (AH: and ASs) (see Fig. 5 and 6) exhibit non-linear variation with solvent composition for all mixed-solvent systems. Extrema in such plots are discernible for EtOH-H,O, Pr'OH-H,O, Bu'OH-H,O, Me-H,O, AC-H,O, DMSO-H,O and WH,O media. Note that the extrema observed for these mixed-solvent media are at rela- tively low values of Xorg.This might have a direct bearing on the solvent structural effects as it is known that a small addi- J.CHEM. 3OC. FARADAY TRANS., 1994, VOL. 90 120 110 "0r I-OQO W22 100 3O 3 0 90 I 1 I I 1808 100 120 140 160 180 2 A9/J K-' mol-' Fig. 7 AH*/kJ mol-' us. AS'/J K-' mol-' for all aquo-organic solvent media investigated. The sizes of the circles vary in order to accommodate a number of data points with their errors within a circle for different solvent compositions of the various solvent systems studied. tion of cosolvents modifies the structure of liquid water. The positions of the maxima in the AH* (or ASS)us. XPriOHplot (at XPriOHx0.05 and 0.18) match those of the sharp minimum in the relative partial molar volume of Pr'OH (v2 -VO,) 30-32 and the maximum in the ultrasonic absorption, respectively, for the Pr'OH-H,O ~ystem.~~.~~AHS (or AS*), however, tends to a maximum at X,,,, x0.20, at which point the relative partial molar volume of DMSO (v2-VO,) 30-32 in the DMSO-H,O medium tends to a minimum.If AHS is split into a reaction component and a solvation componentg (AH*= AHh + AS:), it is reasonable to expect that solvent structural perturbations causing hydrogen-bond formation or breakage will affect the solvation component of AHS,as this effect in the bulk phase will be passed on to the reaction site via the solvation shells of the transition and initial states. The fact that the solvent effects on AHSand AS* are essentially mutually compensatory (see Fig. 7) demon-strates this point. A.N.A. thanks the CSIR, New Delhi for the award of a Senior Research Fellowship. References 1 Y.Cheng, M. Page and C. Jolicoeur, J. Phys. Chem., 1993, 97, 7359. 2 N. Muller, Acc. Chem. Res., 1990, 23, 23. 3 Y. Marcus, J. Chem. SOC.,Faraday Trans., 1990,86,2215. 4 A. C. Dash and P. K. Das, Int. J. Chem. Kinet., 1992,24, 165. 5 A. C. Dash, N. Dash, P. K. Das and J. Pradhan, J. Chem. SOC., Faraday Trans., 1991,87,3753. 6 Y. Marcus, Ion Solvation, Wiley, New York, 1985, ch. 7. 7 A. Al-Alousy, S. Alshehri, M. J. Blandamer, N. J. Blundell, J. Burgess, H. J. Cowles, S. Radulovic, P. Guardado and C. D. Hubbard, J. Chem. SOC.,Faraday Trans., 1993,89, 1041. 8 F. Armand, H. Sakuragi and K. Tokumaru, J. Chem. SOC., Faraday Trans., 1993,89, 1021. 9 M. J. Blandamer, Ado.Phys. Org. Chem., 1977,14,203. 10 A. C. Dash and J. Pradhan, J. Chem. SOC., Faraday Trans. I, 1989,85,2797. 11 A. C. Dash and P. K. Das, J. Chem. Soc., Faraday Trans. I, 1989,85,2405. 12 A. C. Dash, N. Dash and J. Pradhan, Ind. J. Chem., Sect. A, 1992,310,824. 13 A. C. Dash and N. Dash, Int. J. Chem. Kinet., 1990,22, 1237. 14 J. G. Dawber, J. Chem. SOC.,Faraday Trans., 1990,86,287. 15 I. Tejera, A. Rodriquez, F. Sanchez and M. L. Moya, J. Chem. SOC., Faraday Trans., 1991,87,2573. 16 A. C. Dash and G. M. Harris, Inorg. Chem., 1981,20,4011. 17 A. N. Acharya and A. C. Dash, Int. J. Chem. Kinet., 1994, 26, in the press. 18 M. L. Tobe, Acc. Chem. Rex, 1970,3, 370. 19 Ref. 6, p. 186. 20 Ref. 6, p. 188. 21 (a)T. W. Bentley and G. E.Carter, J. Am. Chem. SOC., 1982, 104, 5741;(b)J. E. Leffler and E. Grunwald, Rate and Equilibrium of Organic Reactions, Wiley, New York, 1963, p. 298; (c) G. S. Groves and C. F. Wells, J. Chem. SOC., Faraday Trans. 1, 1985, 81, 2479; (d)R. E. Robertson and S. E. Sugamori, J. Am. Chem. SOC., 1969, 91, 7254; (e) E. Akhtar and R. A. Begum, J. Bangla-desh Acad. Sci., 1978, 2,9. 22 M. J. Blandamer and J. Burgess, Transition Met. Chem., 1988, 13, 1. 23 A. C. Dash and N. Dash, J. Chem. SOC., Faraday Trans. 1, 1987, 83,2505; 1988,84,75. 24 A. C. Dash and J. Pradhan, J. Chem. SOC., Faraday Trans. I, 1988,84,2387. 25 G. M. El-Subruiti, I. M. Sidhamed and C. F. Wells, Znt. J. Chem. Kinet., 1990,22, 891. 26 K. H. Halawani and C. F. Wells, J. Chem. SOC., Faraday Trans. I, 1989,85, 2999. 27 A. C. Dash and P. K. Das, Int. J. Chem. Kinet., 1990,22, 307. 28 A. C. Dash, P. K. Das and J. Pradhan, Transition Met. Chem., 1991, 16, 358. 29 G. M. El-Subruiti and C. F. Wells, Int. J. Chem. Kinet., 1991, 23, 161. 30 K. Nakanishi, Bull. Chem. SOC. Jpn., 1960, 33, 793. 31 J. Kenttamaa, E. Tommila and M. Martti, Ann. Acad. Sci. Fenn., Ser. A, 1959, no. 93. 32 J. Kenttamaa and J. J. Lindberg, Suom. Kemistil. B, 1960,33, 32. 33 M. J. Blandamer, Introduction to Chemical Ultrasonics, Aca-demic Press, London, 1973, ch. 11. 34 D. E. Bowen, M. A. Priesand and M. P. Eastmann, J. Phys. Chem., 1974,78,2611. Paver 4/02159B; Received 12th April, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003293

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3301-3307 Thermodynamics and Kinetics of the Reaction of Copper(l1) and lron(ll1) with Ultra-small Colloidal Chalcopyrite (CuFeS,) Ewen Silvester, Franz Grieser* and Thomas W. Healy School of Chemistry, University of Melbourne, Parkville, Victoria, 3052 Australia Dan Meisel" and James C. Sullivan Chemistry Division, Argonne National Laboratory, Argonne, lL 60439, USA The kinetics of oxidation of colloidal chalcopyrite (CuFeS,) have been examined in the presence of aqueous Fe"' and Cu" at pH 2.3 and in the presence of radiolytically generated Fe(OH),+ at neutral pH. The reaction of Cu" with CuFeS, has been ascribed to an ion-exchange reaction of copper for iron in the CuFeS, lattice. The initial copper sulfide phase that forms, as a result of the exhange, is rich in Cu' and thermally converts to the known covellite (CuS) phase.Oxidation of CuFeS, by Fe"' at low pH leads to the dissolution of the chalcopyrite phase and the formation of Fe2+(aq), Cu"(aq) and So(s). The formation of an unstable intermediate CuS, phase is suggested from thermodynamic calculations. Under neutral pH conditions, CuFeS, is oxidized by Fe(OH),+ adsorbed on the particle surface. Oxidation is restricted to the first monolayer of the particles due to the forma- tion of an Fe'" hydroxide layer at the particle surface. The oxidation of CuFeS, by Fe(OH)'+ is slower than by Fe3+(aq), presumably due to the lower redox potential of the former species. The oxidation of chalcopyrite in acidic Fe"' solutions has received a great deal of attention'-' owing to the applica- tions in commerical hydrometallurgy of this mineral.2 Under acidic conditions the reaction between CuFeS, and Fe"' gen- erally follows CuFeS2(s)+ 4Fe3+(aq)+Cu2+(aq)+ 5Fe2+(aq)+ 2S0(s) (1) although it is not clear whether this equation accurately describes the process relating to the initial reaction kinetics.X-Ray photoelectron spectroscopy (XPS) of air oxidized CuFeS, surfaces under acidic conditions6 has shown that, in the early stages of the reaction, iron migrates into solution while copper remains in what is essentially an oxidized chal- copyrite crystal lattice as Cu', with a composition approach- ing CuS,. Similar findings have been made with electrochemically oxidized CuFeS, electrode^.^ The use of an 'ultra-small '-particle-sized chalcopyrite sol allows a more detailed investigation of the initial reaction kinetics due to the higher proportion of material resident at the solid/aqueous solution interface.This paper deals with the spectroscopic, thermodynamic and kinetic aspects of the reaction of 'ultra-small' chalcopyrite with both Cu" and Fe"' at low pH. In addition the pH dependence of the reaction of Fen' with this mineral has been investigated using pulse radiolysis techniques to generate a transient aqueous Fe"' species [Fe(OH)'+] at neutral pH. Experimental The preparation and dialysis of chalcopyrite sols have been described previously.8 These sols were saturated with argon in gas syringes and stored until required; usually less than 5 days.In all experiments the CuFeS, sol concentration was determined from the sol absorbance at 480 nm (&480 = 7100 dm3 mol- cm-'). Reaction of Chalcopyrite with CU" and Fe"' at pH 2.3 Cu" and Fe"' solutions were prepared from perchlorate salts in lop2 mol dmp3 perchloric acid. Chalcopyrite sols pre- pared at (3.0 k0.3) x mol dm-3 were reacted with Cu" or Fe"' solutions in equal volumes such that, upon mixing, the chalcopyrite concentration was (1.5 & 0.2) x mol dm-3 and the pH ca. 2.3. All the solutions were saturated with argon gas prior to reaction and all experiments were conducted at a temperature of 25 "C. In a separate set of experiments the long-term reaction products were determined by analysis of the metal ion concentrations/speciation in solution. Separation of the aqueous phase from the particles was by filtration of the sol through an Amicon PM30 membrane (pore diameter ca.20 A)in an Amicon 8010 stirred cell assembly. Total copper and iron concentrations in the filtrate were determined by a stan- dard atomic absorption (AAS) technique. Fe" was determined spectrophotometrically using the o-phenanthroline method.' Stopped-flow Measurements Reaction kinetics were studied by the stopped-flow technique using a Hi-tech SF-51 stopped-flow unit. In this system, time- resolved absorbances were recorded at single wavelengths over a pathlength of 0.3 cm. Chalcopyrite sols were reacted with iron(1Ir) solutions in the range 5 x lo-' to mol dm-3 and the sol bleaching monitored at 480 nm.Under moderately acidic conditions Fen' speciation is dominated by Fe3+, Fe(OH)2+ and Fe2(OH)24+, with Fe3+ being the prin- ciple form (> 70%). Both Fe(OH)*+ and Fe,(OH),4+ absorb light at 480 nm;" however, the concentrations of these species were sufficiently low to allow this contribution to the measured absorbance to be ignored. The solubility of crystalline haematite (Fe20,) is exceeded at the higher concentrations of Fe"' reacted ; however, the amorphous Fe(OH), phase, with a precipitation edge at a higher pH, is more likely to form initially." The reduction of Fe"' by the sol results in the formation of Fe", of which the dominant species under acidic conditions is Fe(H,0),2+(aq).'o This species absorbs in the UV region below 300 nm and very weakly in the near IR (NIR)12 and so does not contribute to the measured absorbance at 480 nm.At low pH direct acid dissolution of the sol is possible. The contribution of this pathway to the overall dissolution is dis- cussed in the subsequent sections. Pulse Radiolysis of CuFeS,-Fe" Solutions at Neutral pH A series of sol diluents containing Fe" (prepared from FeS04.7H,0) were mixed with the CuFeS, sol in varying ratios to give a series of mixtures, all containing 5 x Table 1 Initial solution concentrations used in the pulse radiolysis experiments CuFeSJmol dm - Fe"/mol dm - PH 1.0 x 10-4 5 x 10-4 6.63 5.0 x 10-5 5 x 10-4 6.70 2.6 x 10-5 5 x 10-4 6.57 1.3 x 10-5 5 x 10-4 6.15 moI dm-3 Fe", but with varying sol concentrations.The CuFeS, sol and Fe" concentrations studied and the corre- sponding pH of the mixed solutions are given in Table 1. All the sols and diluents were saturated with N,O prior to mixing and the mixing was carried out shortly before irradia- tion. Irradiation of Solutions The irradiation of aqueous solutions with ionizing radiation has been described previously. l3 The Argonne National Laboratory EINAC facility and spectrophotometric detection system used have been described el~ewhere.'~*'~ Solutions and sols were irradiated in quartz cells with optical paths between 0.5 and 2.0 em. Radiation pulses were multiples of 40 ns pulses (up to a maximum of 10 pulses) with a repetition rate of 60 Hz.In this way radiation doses between 2 and 17 krad were applied to the sols. Dosimetry was achieved by irradiation of mol dm-3 KSCN solutions saturated with N2O.I3 In N,O saturated Fe" solutions the following series of reactions occur after primary irradiation H,O-+e-(aq), OH', H,O,, H, ,H30+ [predominantly e(aq)- and OH'] (2) e-(as) + N,O + H,O --+ N, + OH' + OH-(3) Fez+ + OH'-+ Fe(OH),+ (4) yielding an Fe'" species which can then react with the sol. In the absence of the chalcopyrite sol the absorbance spectrum immediately after irradiation (100 ms) was found to corre- spond closely to a previously measured spectrum for Fe(OH)'+.'* This first hydrolysis product of iron" was observed to be kinetically stable in the second before further hydrolysis and the eventual formation of Feu' hydroxide.After ca. 10 s, the absorption spectrum resembled that of iron'" hydroxide particles in the <10 nm size range.16 Changes in the sol absorbance were monitored at 480 nm. At longer times and at higher radiation doses, the formation of Fe"' hydroxide contributed significantly to the measured absorption at this wavelength. Results and Discussion Reaction of Chalcopyrite with Cu" at Low pH Fig. 1 shows the UV-VIS absorption spectrum of a 1.5 x mol ~im-~ CuFeS, sol (a) before and (b) after reaction with mol dmm3 Cu'. Reaction of the chalcopy- rite sol with aqueous Cu" results in significant bleaching of the sol absorption at 480 nm indicating loss of this mineral phase.Accompanying the observed bleaching is a pro-nounced broadening of the absorption band. This was attrib- uted to an increased scattering component in the sol extinction. When heated to >70°C for ca. 1 min the reacted sol developed a strong absorption band in the NIR region [Fig. l(c)] which, for the reasons outlined below, can be assigned to the presence of crystalline covellite" (nominally CuS, although more accurately described as Cu,S * CuS,"). J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1.o 0.8 0.6 0 m -e8 % 0.4 0.2 I I I I I I I 10.0 300 400 500 600 700 800 900 wavelength/nm Fig. 1 UV-VIS absorption spectrum of a 1.5 x rnol dm-j CuFeS, sol (a) before and (b) after reaction with Cu" (pH 2.3) at an initial concentration of rnol dm-'.Heating the reaction product yielded a sol with spectrum (c), corresponding to that of crystalline covellite (CuS). Copper sulfide sols when first precipitated from solution are 'golden-brown' in colour. Electron diffraction studies indicate that the particles are a poorly crystalline form of covellite while XPS analysis shows that copper is exclusively in the form of Cu'.'' With time, and at an enhanced rate at higher temperature, these sols change to a green colour and concomitantly become more crystalline. The NIR band that forms is attributed to a transformation of a third of the copper in the particle to Cu", consistent with the known structure of crystalline covellite.The intensity of the NIR absorption band gives a good quantitative measure of the crystalline covellite concentration. In the case shown in Fig. l(c) it is readily calculated that greater than 90% of the copper in the system [both from Cu"(aq) and CuFeS,] has been converted into covellite, indicating a nearly quantitative replacement of chalcopyrite by covellite under these condi- tions. The absence of the NIR band before heating suggests that the initial copper sulfide product is exclusively in the Cu' phase. Fig. 2 shows the aqueous concentrations of iron and copper as well as the concentration of adsorbed copper after reaction of CuFeS, sols with Cu" over a range of initial con- centrations.Virtually all (>95%) the aqueous iron was found to be in the divalent form. Feu is significantly leached from the particle even in the absence of Cu', indicating that acid dissolution is a major reaction pathway at low pH. The absence of any measurable aqueous copper under these con- ditions suggests that acid dissolution takes the form CuFeS,(s) + 2H+(aq)+CuS(s) + Fe2'(aq) + H,S(aq) (5) J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 IL E0 n h Y + Na k n "0 2 I I 1 2[C~~+(aq)]/lO-~rnol dm-3 Fig. 2 Concentrations of aqueous Fe" (a),aqueous Cu" (W) and adsorbed Cu" (O),10 min after reaction of CuFeS, with Cu" at initial concentrations of up to 2 x rnol dm-3. Initial [CuFeS,] = 1.5 x rnol dm-3; pH 2.3.Insert: Aqueous Cu" concentration as a function of initial [Cu2+(aq)], on a less sensitive scale (x axis, Cu2+(aq)/10-3 rnol dm-j; y axis, CuAaq)/lO-' mol dm-3). For systems in which Cu" was added, the more than quan- titative replacement of copper for iron indicates that very little sulfide departs the system in the form of H2S(g), suggest- ing a direct ion-exchange reaction, i.e. CuFeS,(s) + Cu2+(aq)+2CuS(s) + Fe2+(aq) (6) Thermodynamic calculations show that, in the presence of Cu" significant transformation of CuFeS,(s) to CuS(s) should occur, the extent of which depends on the H2S(aq) concentra- tion in the solution. While this concentration is unknown, it is useful to consider an extreme case. It has been noted pre- viously that CuFeS, sols contain a slight excess of adsorbed S2-(ca.3%)." This excess sulfide would correspond to an H,S(aq) concentration of cu. mol dm-3, and, in turn, to solution concentrations of Cu2 '(as) and Fe2+(aq) in equi- librium with CuFeS,(s) and CuS(s) of the order of 10-l4 and mol dm-3, respectively. This concentration of Fe2'(aq) is very close to the value observed experimentally. At higher concentrations of added Cu" the level of H2S(aq) would be considerably lower, due in part to direct precipitation of CuS(s). Accordingly, the equilibrium level of Fe2 '(aq) should be higher. It is surprising that, even at the highest concentra- tions of Cu", the release of Fe" is limited to mol dm-j. Given that this concentration corresponds to two thirds of the total iron in the particle it seems likely that the observed level is a result of an inhibited reaction with the particle core rather than thermodynamic constraints.The adsorption of Cu" above the level at which Fe" is liberated into solution indicates that a further mechanism of copper retention must be operating; this is presumably the adsorption of Cu" onto the modified particle surface. An attempt was made to study the kinetics of the reaction of Cu" with chalcopyrite by the stopped-flow technique; however, the bulk of the reaction occurred within the mixing time of the apparatus (ca. 20 ms). Acid dissolution was observed to be a much slower process, confirming that the reaction of Cu" with the particle is not preceded by acid dis- solution.While the reaction of Cu" with CuFeS, does not outwardly appear to be a redox process, electron transfer must be involved. Copper is added as Cu" although no Cu" is evident in the UV-VIS absorption spectrum of the reacted colloid, i.e. the NIR band of 'green' copper sulfide is initially absent. Similarly, while iron is believed to be present in chalcopyrite as Fe"', it enters the solution as Fe". These redox changes can be reconciled in terms of the reaction Cu'Fe"'S:'-(s) + Cu"(aq) -+ Cu$"-S'(s) + Fe2+(aq) (7) where Cu\S"S' represents a Cu' sulfide mineral of overall stoichiometry CuS. Reaction of Chalcopyrite with Fe"' at Low pH The reaction of CuFeS, sol with Fe"' resulted in bleaching of the sol absorbance at 480 nm and the formation of Fe" in solution, indicating that oxidation of the sol had occurred. In Fig.3 is shown the concentration of copper that appears in solution with increasing initial Fe"' concentration. For initial Fe"' concentrations of < rnol drn-j, copper is not detected in solution, even though oxidation of the sol occurs. Above this concentration, copper is detected in the aqueous phase, increasing to a level close to that expected for the complete oxidation of the chalcopyrite phase. Since neither Cu' nor Cu" hydrolyse at this pH, the copper retention that is observed can only be due to association with sulfide. The observed behaviour can be accounted for by the sequential reactions, CuFeS,(s) + 2Fe3+(aq) -,CuS,,(s)+ 3Fe2+(aq)+ (2 -y)S'(s) (8) CuS,(s) + 2Fe3+(aq)+Cu2&(aq)+ 2Fe2+(aq)+ ySo(s) (9) where the parameter (2 -y) represents the degree to which the partially oxidized lattice has nucleated to form S'(s).This interpretation is consistent with previous studies on chal-copyrite electrodes2' and chalcopyrite slurries' which have found preferential release of iron in the initial stages of reac- tion. XPS studies indicate that the copper remains as a copper(r) sulfide phase most accurately described as an iron deficient chalcopyrite. 6,7 If no adsorption of Fe" occurs the number of electrons transferred per CuFeSz unit N,, can be obtained from the equation, At low initial concentrations of Fe"' the sol bleaching is a reliable measure of the concentration of CuFeS, sol reacted, with little contribution from light scattering to the sol absorption spectrum.At higher initial Concentrations of Fe"' (>5 x mol drn-j) light scattering is evident, indicative 41 t? Ne 3-2-I 1 ' . '.--C 0I ' ' """I ' ' """I -5 -4 -3 -2 log aFeo+ Fig. 3 Concentration of Cu" (0)produced in solution as a function of initial Fe"' concentration, and the number of electrons transferred per CuFeS, unit (0)calculated as described in the text. Initial [CuFeS,] = 1.5 x mol d~n-~;pH 2.3. 3304 of either elemental sulfur formation or sol aggregation. For these systems the Cu" concentration has been used as an approximate measure of chalcopyrite oxidation.In Fig. 3 is shown the reaction stoichiometry, calculated as described above, over the range of initial Fern' concentrations studied. At high excess of oxidant the stoichiometry has the expected upper limit of four electrons per CuFeS, ,consistent with that found by Linge' using electrochemical techniques and consis- tent with reaction (1). The data do not exhibit the lower limit of two electrons per CuFeS, expected from reaction (8) which is presumably due to the concurrent acid dissolution reac- tion, the effect of which is greater at lower initial Fe"' concen- trations and the stoichiometry of which is zero is terms of eqn. (I). In the thermodynamic interpretation of this system, it is instructive to construct a predominance area diagram in terms of the solution species involved in the reaction dynamics. Such a diagram is shown in Fig.4,where the sta- bility regions have been calculated using the log absorption coefficients, log aFe3+ and log aFez+,as variables. This figure has been constructed for constant Cu2'(aq) activities of and lo-, mol drn-,, total dissolved sulfur concentra- tion of mol dm-, and constant pH (2.3). The assump- tion of constant H,S(aq) activity is inaccurate but it allows a qualitative description of the system. As an initial approx- imation, the Gibbs energy of formation (AfGe) of the CuS, phase has been taken to be the same as that for covellite. The oxidation of CuFeS, can be considered in terms of its reaction path, as is shown in Fig. 4,starting at aFe3+= and uFe2+= lo-'' mol dm-,.Oxidation of CuFeS, under these conditions results in the formation of Fe2+(aq), both from CuFeS, and as a reduction product of Fe" and Cu2'(aq). The increasing concentration of Cu'+(aq) leads to a broadening of the CuS(s) stability field, as shown by the successive CuS(s)/Cu2 +(aq) boundary lines on this figure. An equilibrium point is not marked on this diagram; under the conditions considered it would be the intersection point between the three predominance regions: CuFeS, ; CuS, Fe2+, H,S; and CuS, Fe2+, So, the position of which is con- trolled by the activity of H,S(aq) which will change during the reaction. A copper sulfide product with a Af G" close to that of covellite should not form to a significant extent since the bulk of the reaction occurs outside the stability field of log aFeZ+ -1 5 -1 0 -5 0 0 -5 0)Y (D m --1 0 // '/ / , -1 5 Fig.4 Phase diagram for the Cu-Fe-S system in terms of the master variables log uFe3+and log uFez+showing the reaction path for the oxidation of CuFeS, starting at initial conditions of aFe3+= rnol dm-, and uFez+= lo-" mol dm-'. Diagram prepared for conditions of constant activity of Cu2+ lod4 or lo-' mol drn-,), constant pH (2.3) and constant total dissolved sulfur of mol dm-3. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 CuS(s). The retention of copper that is observed in these systems must therefore be due to either the formation of a CuS,(s) phase of significantly lower Af G * than that of covel- lite, or to kinetic restrictions on the release of copper from the oxidized chalcopyrite lattice.In Fig. 4 is shown where the CuS,(s)-Cu2'(aq) boundary would need to be in order for a copper sulfide to be the dominant oxidation product. The position of this line corresponds to a Gibbs energy of formation of the order of -130 kJ mol-', compared to that for crys- talline covellite of -48.9 kJ mol- '. This value would seem to be unreasonably low and kinetic restrictions on the release of copper from the oxidized chalcopyrite lattice must be favoured as the reason for the observed retention of copper. Kinetics of Chalcopyrite Oxidation of Fe"' at Low pH Bleaching of the sol absorbance occurred on the millisecond to second timescale.From the preceding discussion it can be noted that a considerable excess of Fe"' is necessary for the reaction to be described by reaction (1). Even with a high excess of Fe"' it appears unlikely that this reaction would adequately describe the sol bleaching on short timescales. Acid Dissolution of Chalcopyrite The effects of perchloric, nitric and hydrochloric acids upon the absorbance of chalcopyrite at 480 nm were investigated to determine the extent of direct acid dissolution in these systems. In Table 2 are shown the initial sol absorbances at 480 nm [A(O)],the apparent first-order rate constant fit to the bleaching of the sol absorbance at this wavelength, and the sol absorbance at longer times (20 s) after mixing [A(oo)] for sols reacted with these acids.The measured A(0) values compare well with the expected value (A = 0.31; [CuFeS,] = 1.5 x mol drn-,) indicating that very little bleaching of the sol occurs within the mixing time of the stopped-flow system. After mixing, some bleaching of the sol absorbance is observed, although as will be shown in the fol- lowing section, both the extent of bleaching and the fitted rate constants are small compared to that observed in the presence of Fe"'. Also shown in Table 2 is the effect of 5 x lo-, mol dm-, Al"' (in 5 x lo-, mol dm-, HNO,) upon the sol absorbance. Sol bleaching under these conditions is less than that observed for an equilvalent concentration of nitric acid alone. It appears that aluminium has the effect of protecting the sol from dissolution, presumably via specific binding to the sulfide surface thus preventing interaction of protons with the surface ligands.Chalcopyrite Oxidation by Fe"' at pH 2.3 Fig. 5(a) shows the measured initial sol absorbances for the range of Fe"' concentrations studied. At all concentrations partial oxidation of the chalcopyrite sol occurred within the mixing time of the stopped-flow apparatus (20 ms). At low initial Fe"' concentrations, the measured A(0)values decrease Table 2 Initial absorbance [A(O)],apparent first-order rate constant (k),and final absorbance [A(oo)],for the acid dissolution of chalcopy-rite" HCI 0.306 0.32 0.277 HC10, 0.294 0.978 0.277 HNO, 0.303 0.7 12 0.288 HNO,b + Al(NO,), 0.311 0.07 0.294 ~~ " Concentration of all acids was 0.005 mol drn-,.Optical path length was 0.3 an. 5 x lo-, mol dm-, Al(NO,), in 0.005 mol dm-, HNO,. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 -0.28 0.26 0 0.24 1 0 0 I.,.,.,.,,0.22 01 23456 5 4 -3 k.I s2 1 B e 0 0' ' ' ' ' I " ' ' 1 0 1 2 3 4 5 6 0.10' ' ' . ' ' ' . ' . J' I 01 23456 [Fe"']/10-3 mol dm-3 Fig. 5 (a) Initial absorbances at 480 nm for a CuFeS, sol oxidized by Fe"' at pH 2.3, using the stopped-flow technique. The dashed line shows the expected A(0)value. (b) Fitted first-order rate constants for CuFeS, sol bleaching at 480 nm. The decays were recorded over a 2 s duration.(c) A(m) (sol absorbance after 20 s) values for kinetic measurements made over a 2 s duration. with increasing Fe"' concentration suggesting that the rate of this initial process is dependent on oxidant concentration. At higher Fe"' concentrations the measured A(0)values increase, presumably due to the formation of light absorbing or scat- tering products. Previous studies have shown that CQ. 35% of the CuFeS, in colloidal chalcopyrite resides in the first mono- layer." At all but the lowest concentrations studied, ca. 30% of the sol absorbance is lost during the mixing time. It appears therefore that very little of the oxidation of the first monolayer of the CuFeS, particles is captured by the stopped-flow technique. Bleaching of the sol absorption was observed over several orders of magnitude in time without any distinct separation between different reaction steps. Single-exponential fits were performed on the measured sol bleaching over a several periods of time (2,20 and 200 s) in order to obtain an indica- tion of the rates involved.At all Fe'" concentrations the fitted rate constant value decreased with increasing measurement time, indicating that the reaction is not a single-exponential process. Fig. 5(b)shows the fitted rate constants for measure- ments made over a duration of 2 s. The initial increase in the fitted rate constant with increasing Fe"' concentration further confirms the dependence of the rate of oxidation of the first monolayer on oxidant concentration.In terms of the rate- determining step, this dependence could be interpreted as either a collisional encounter in solution or adsorption of Fe"' onto the particle surface. At the higher Fe"' concentra-tions, the CuFeS, sol bleaching rate is independent of Fe"' concentration. In this domain, which corresponds to particle oxidation beyond the first monolayer, the rate-determining step is possibly either the release of ions into solution from the particle surface or electron transfer from a species in the solid particle to an adsorbed Fe"' species. Fig. 5(c) shows A(co) values corresponding to the data shown in Fig. 5(a) and (b).Again, the presence of an absorb- ing or light scattering product is evident at higher Fe"' con-centrations. Similar observations were made on all the timescales studied.Pulse Radiolysis of Fe"-CuFeS, Solutions at Neutral pH To overcome the mixing time restrictions of the stopped-flow technique, pulse radiolysis experiments were undertaken with the oxidant created in situ by the radiation pulse. Initial con- centrations of Fe(0H)' +,generated according to reactions (2)-(4), were varied by altering the radiation dose applied to CuFeS, sol-Fe" mixtures. In this way initial concentrations of Fe(OH)'+ could be generated in the range 1.2 x to 1.0 x lop4mol dm-3. Unlike the low pH systems, CuFeS, sol bleaching was adequately described by first-order kinetics at all doses and it is these rate constants which are con- sidered in the following discussion. Fig.5 shows the first-order rate constant (@-I) against dose [or initial Fe(OH), concentration], for several sol con- + centrations. Only at the lowest sol concentration does the measured first-order rate constant vary linearly with the oxidant concentration, as might be expected from the relative concentrations of sol particles and the oxidant present. Using the results from the lowest sol concentration, the bimolecular rate constant can be calculated to be 8.5 x lo4 dm3 mol-' s-' in terms of CuFeS, molecules, or 1 x lo8 dm3 mo1-l s-' in terms of CuFeS, particles. It is clear that the reaction observed at near-neutral pH is substantially slower than the reaction at low pH (see previous section). This is probably due to the lower redox potential of the Fe(OH),+/Fe2+ couple compared to that for Fe3+/FeZ+.Interestingly, the fitted rate constant value decreases with increasing sol con- centration. It seems that the rate of oxidation is determined by the number of oxidant species per particle and not by the total concentration of surface sites. Hence, the oxidation occurs by adsorbed Fe"' species and not solution species. Fig. 7 shows the total sol bleaching [derived from A(w) values] against initial Fe(0H)' concentration for several sol+ concentrations. Also shown is the expected bleaching for the 61 1 0 10 000 20 000 dose/rad I 0 4 8 12 [Fe(0H) *+I ,,,,/10 -mol dm -' Fig. 6 First-order rate constants for CuFeS, sol bleaching at 480 nm as a function of dose, or initial Fe(OH)'+ concentration (lower scale), for sol concentrations of (0)5 x lo5and (m) mol dm-3 1.3 x lo-', (0)2.6 x (0) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0.00 T,, 1 dose/rad I I 0 4 8 12 [Fe(OH)2+]iniJ10-5 mol dm-3 Fig. 7 Total CuFeS, sol bleaching at 480 nm [as derived from mea- sured A(m) values] us. applied dose, or initial Fe(OH)'+ concentra- tion (lower scale), for sol concentrations of (0)1.3 x lo-', (0) 2.6 x (0) lo-' and (I)mol dm-3. The expected sol5 x bleaching for the first monolayer of reaction for the three lowest sol concentrations is shown by the dashed lines. The monolayer absorb- ance of the highest sol concentration is off scale, as indicated. first monolayer of the particles (dashed lines).For all sol con- centrations, sol bleaching is limited to the first monolayer and is apparently diminished at higher doses. As was noted in the experimental section, the generation of Fe"' at pH 6.5 will result in the formation of solid iron(@ hydroxide after several seconds, the absorption spectrum of which will overlap with that of the sol at 480 nm. At higher doses the contribution of iron(@ oxide to the measured absorbance is clearly significant. It is likely that the limitation of the reac- tion to the first monolayer is also a result of the formation of iron(@ hydroxide given that it would probably form as a surface coating on the CuFeS, particles. It can be readily shown that layers of greater than 10 A in thickness could form on the particles at the higher doses employed.For elec- tron transfer to occur across such a layer, charge carriers must be injected into one of the bands of the iron(u1) hydrox- ide film. The rate of reaction then becomes a function of both the relative positions of the conduction and valence bands of the two solids and the conductivity of the film.,' Fig. 8 shows schematic diagram of the relevant energy levels for the CuFeS,/Fe(OH),/water interface at pH 7. It is clear that the concentration of valence band holes in iron(rI1) hydroxide, as controlled by the Fe(OH),+/Fe2+ couple, would be low under these conditions. Given that the mechanism of electron transfer in this system is hole transfer from adsorbed Fe(OH),+ to iron(@ hydroxide and then to CuFeS, ,the rate of electron transfer would be significantly diminished by the presence of the iron(m) hydroxide layer.In addition to the restrictions on electron transfer, there would also be restrictions on the accompanying charge-balancing steps. A thin product layer on the particles would be sufficient to provide a physical barrier to ion movement into solution and limit reaction to the first monolayer of the CuFeS, particles. Conclusions In the low-pH reactions of both Cu" and Fe"' with CuFeS, particles, Fe" is preferentially released from the colloid during the initial stages of the reaction. While the reaction of Cu" with CuFeS, appears to be an ion-exchange process, of I 0.6 I -Fe(OH)2'/Fe2' v) W I 0.8 I __t r5 1.0 ' Lu 1.2 CuFeS2 particle core aqueous 1.4 1.6 1.8 2.22.0 1 2.4 1 Fig.8 Redox levels for a chalcopyrite/iron(III) hydroxide interface, in contact with an aqueous solution at pH 7 (us. SHE). Chalcopyrite band positions are estimates from ref. 19, while iron(iI1) hydroxide bands are estimates based on the positions of these bands for haema- tite (Fe,O,). Also shown are the positions of the H+/H, and Fe(0H)' +/Fez-+ redox couples in solution and the relevant corrosion potentials (Edccomp)for CuFeS, . copper for iron in the CuFeS, structure, an intra-lattice elec- tron-transfer step must be involved. The initial copper sulfide phase which forms is rich in Cu' and is thermally converted into the known covellite phase.For the reaction of Fe"' with chalcopyrite the release of Cu" appears to be kinetically rather than thermodynamically controlled. The low-pH reac- tion of Fe"' [dominantly Fe3 '(as)] with colloidal chalcopy- rite is much faster than that at near-neutral pH [Fe(OH)2+(aq)], presumably due to the lower redox poten- tial of the latter. Further, the reaction at neutral pH is restricted to the first monolayer of the chalcopyrite particles due the formation of an iron(rr1) hydroxide layer at the parti- cle surface. The dynamics of the oxidation reaction by Fe(OH), + are determined by intra-particle transfer from the lattice to Fe(OH)2 + adsorbed at the particle surface. This work was supported in part by the Australian Research Council (ARC) in the form of a Special Research Centre grant to the Advanced Mineral Products Centre.Work at the Argonne National Laboratory is performed under the auspices of the ofice of Basic Energy Sciences, Division of Chemical Sciences, US-DOE under contract no. W-31-109- ENG-38. E.S. acknowledges the receipt of a Commonwealth Post-Graduate Resarch Award. References 1 H.G. Linge, Hydrometallurgy, 1976,2,51. 2 J. E. Dutrizac, Metall. Trans. B, 1978,9B, 431. 3 P. B. Munoz, J. D. Miller and M. E. Wadsworth, Metall. Trans. B,1979, lOB, 149. 4 D. L. Jones and E. Peters, Extractiue Metallurgy of Copper ZI, The Metallurgical Society of AIME, New York, 1976. 5 J. E. Dutizac, Metall. Trans. B, 1981, 12B, 371. 6 A. N. Buckley and R. Woods, Aust.J. Chem., 1984,37,2403. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3307 7 A. N. Buckley, I. C. Hamilton and R. Woods, in Flotation of 15 S. Gordon, K. H. Schmidt and E. J. Hart, J. Phys. Chem., 1977, Suljide Minerals, Developments in Mineral Processing, ed. K. S. 81,104. E. Forssberg, Elsevier, Amsterdam, 1985, vol. 6. 16 P. Mulvaney, R. Cooper, F. Grieser and D. Meisel, Langmuir, 8 E. J. Silvester, T. W. Healy, F. Grieser and B. Sexton, Langmuir, 1988,4, 1206. 1991, 7, 19. 17 E. J. Silvester, F. Grieser, B. Sexton and T. W. Healy, Langmuir, 9 A. I. Vogel, Quantitative Inorganic Analysis, Longmans, London, 1991,7,2917. 1953. 18 D. J. Vaughan and J. R. Craig, Mineral Chemistry of Metal Sul-10 C. F. Baes and R. E. Mesmer, The Hydrolysis of Cations, Wiley, fides, Cambridge University Press, Cambridge, 1978, p. 180. New York, 1976. 19 E. J. Silvester, F. Grieser, D. Meisel, T. W. Healy and J. C. Sulli-11 R. M. Garrels and C. L. Christ, Solutions, Minerals and Equi- van, J. Phys. Chem., 1992,%, 4386. libria, Freeman, Cooper and Company, San Francisco, 1965. 20 T. Biegler and M. D. Home, J. Electrochem. Soc., 1985, 132, 12 M. Ehrenfreund and J-L. Leibenguth, Bull. SOC. Chim. Fr., 1970, 1363. 7,2494. 21 S. R. Morrison, Electrochemistry at Semiconductor and Oxidized 13 I. G. Draganic and Z. D. Draganic, The Radiation Chemistry of Metal Electrodes, Plenum Press, New York, 1980. Water, Academic Press, New York, 1971. 14 D. Meisel, W. A. Mulac and M. S. Matheson, J. Phys. Chem., 1981,85,179. Paper 4/04335I; Received 15th July, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003301

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3309-3313 Reaction of Peroxornonosulfate Radical with Manganese(![) in Acidic Aqueous Solution A Pulse Radiolysis Study J. Berglund and L. 1. Elding Inorganic Chemistry 1, Chemical Center, P.O. Box 124,22100 Lund, Sweden G. V. Buxton, S. McGowan and G. A. Salmon" Cookridge Radiation Research Centre, The University of Leeds, Cookridge Hospital, Leeds, UK LS16 6QB The reaction between the SO5-radical and Mn" has been proposed to be the most important process for regeneration of Mn"' in the Mn'"/''-catalysed autoxidation of S" in acidic aqueous solution. In the present study, the second-order rate constant for this reaction has been determined at pH 3.0 and 10 mmol dm-, ionic strength by use of pulse radiolysis. The study was performed in the presence of excess S".Under these conditions Mn" is distributed among the complexes Mn2'(aq), [Mn(HSO,)]+ and [Mn(S03)Mn12+. The rate of reaction decreases as a function of increasing [Mn' which is rationalized qualitatively by a mechanism involving three parallel reactions between SO5-and the Mn" complexes, with rate constantsk,, , k,, and k,, , respectively. H+Mn2++ SO5--Mn3++ HSO,-(16) H+ [Mn(HSO,)]+ + SO,--[Mn(HSO3)l2+ + HSO,-(17) H+[Mn(S0,)Mnl2+ + SO,--[Mn(S03)Mn13' + HS0,-(18) For increasing total concentrations of Mn", formation of the sulfito-bridged complex is favoured which implies that k,, < k,, , k,, . Values of the second-order rate constant in the range 2 x lo8-2 x 10,' dm3 mol-' s-' have been determined, depending on which Mn" species is predominant.Subsequent slow processes are observed following the formation of Mn"'. These reactions have been attributed to the disproportionation of Mn"' and reactions between the Mn"' species and excess S". The implications of the present results for the Mn'l'/l' catalysed autoxidation of S'' are discussed. Recent results indicate that aqueous phase oxidation of mechanism of the reaction between hexaaquamanganese(I1) sulfur(rv)(SO,. nH,O, HSO,-, SO,,-) by molecular oxygen, and the peroxomonosulfate radical with pulse radiolysis in catalysed by metal ions, uiz. MnIrl/I1, Fel*I/I1, Co"'/" and CulI1/I1, order to determine the rate constant for this reaction and takes place by a common free-radical chain mechanism. gain a better understanding of the manganese-catalysed The chain is initiated by reaction between the trivalent metal autoxidation of SIV.ion and SIV according to reaction (1) and is propagated by SO,-, SO5-and SO4-radicals. Experimental M"' + HSO,--M" + SO,-+ H+; Chemicals and Solutions (M = Mn, Fe, Co, Cu) (1) A stock solution of 2.72 mmol dm-, manganese@) perchlor- ate was prepared by dissolving an accurately weighed The reduced metal ion is re-oxidized by the strong oxysulfur amount of Mn(ClO,), -H20 (Johnson Matthey GmbH) in radicals SO5-and SO4-as well as by hydrogen per-100 ml water. The pH of the secondary manganese@) solu- oxornonosulfate, HS05-, which is also generated in the tions was adjusted to pH 3.0 by use of perchloric acid (BDH,chain, reactions (2)-(4).60%)and the ionic strength was kept at 10 mmol dmP3 using sodium perchlorate (BDH, 99.9%) as supporting electrolyte. H+M" + SO,--M"' + HS05-(2) The solutions were saturated with oxygen by use of the stan- dard syringe-bubbling technique as described previo~sly.~ M" + SO4--M"' + MI' + HS05--M"' + SO4-+ OH-(3) An oxygen-free stock solution of ca. 20 mmol dm-, (4) sulfur(1v) was freshly prepared before each set of experiments by dissolving ca. 0.26 g Na,SO, (Merck p.a.) in 100 ml 20 Consequently, these reactions are very important in order to mmol dm- perchloric acid, which was continuously flushed close the catalytic cycle and to rationalize the complicated with argon. The concentration was checked by use of stand-processes governing metal ion catalysed autoxidation of S".ard ion chromatography and found to be constant to within However, the kinetics and mechanisms for most of these reac- 97% of the prepared value. The pH and ionic strengths of the tions have not been studied. secondary sulfur(1v) solutions were adjusted to 3.0 and 10 Recently, it has been proposed that reaction (2) is mmol dm-,, respectively, and they were saturated with responsible for the regeneration of Mn"' in the manganese- nitrous oxide using the syringe-bubbling technique. catalysed process and that reactions (3) and (4)can be All solutions were prepared by use of deionized (Millipore, neglected, provided reaction (2) is fast, i.e. k, 9 lo4 dm3 Milli-Q) water and all gases were used as received from the mol-' s-l.' Therefore, we have studied the kinetics and suppliers without additional purification.J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Apparatus and Measurements Pulse radiolysis experiments were carried out using a pulsed electron beam generated by a 3 MV Van de Graaff acceler- ator. Peroxomonosulfate radicals were produced by irradiat- ing, with 17-41 Gy, 0.2 ps pulses, a solution containing Mn" and SIV,half saturated with N,O and oxygen, and contained in a triple-pass cell of optical length 6.8 cm,see below. In order to minimize thermal oxidation of SIV before irradiation, equal volumes of separate solutions of N,O-saturated S" and 0,-saturated Mn" were mixed in a rapid-mix apparatus less than a second before the pulse was delivered.Thus, in all the experiments, the irradiated solutions were half-saturated with both N,O and O,, i.e. [N,O] = 0.0125 mol dm-3 and [O,] = 675 pmol dm-3. A xenon arc lamp, operated in either continuous or flashed mode, was used as the analysing light source. The dose per pulse was measured using the standard thiocyanate dosimeter (lo-, mol dm- 3, saturated ~with oxygen, for which GE= 2.48 ~x ~ m2 J-'.' The experiments were carried out at ambient temperature, i.e. 20 & 1 "C. Kinetic data were evaluated using either the TREAT program,6 developed to analyse simple pulse radiolysis traces, or the FACSIMILE kinetic modelling ~rogram.~ Generation of SOs-The radiolysis of an aqueous solution results in the gener- ation of hydrated electrons, eaq-, hydroxyl radicals, a small yield of hydrogen atoms and small amounts of hydrogen per- oxide and hydrogen, reaction (5).In the presence of N,O, the hydrated electrons are converted to OH radicals by reactions (6) and (7) with k, = 9.1 x lo9 dm3 mol-'s-' and k, = 9.4 x 107 s-,.~ ksN,O + eaq--N, + 0-0-+ H,O OH + OH-(7) The reaction of OH with Mn" is a hundred times slower than that of OH with hydrogen sulfite and conditions were chosen such that 99% of the OH radicals reacted with HS03-, gen- erating sulfite radicals according to reaction (8) with k, = 4.5 x 10' dm3 mol-'~-~ ksHS03-+ OH -SO3-+ H20 (8) Sulfite radicals react rapidly with dissolved molecular oxygen in a subsequent step forming peroxomonosulfate radicals, reaction (9), with kg = 2.5 x lo9dm3 mol-I s-l.' k9SO3-+ 0, -SO5-(9) SO5-radicals are known to be involved in the chain propa- gation steps of the free-radical chain oxidation of SIV by molecular oxygen, the mechanism of which has been dis- cussed in detail by McE1roy.l' The relevant reactions at pH 3.0 are reactions (10) and (ll), the rate constants of which have been determined to be:" k,, = (1.0 0.1) x lo3 and k,, = (1.1 & 0.1) x lo4dm3 mol-' s-'.An earlier estimate of these rate constants gave only the upper limit, i.e. (klo+ kll) < 3 x lo5dm3 mol-' s-l.', SO5-+ HS0,-+ HS0,-+ SO3-(10) SO5-+ HS03--+ SO4-+ HS04-(1 1) These values indicate that these reactions are unimportant under the conditions and timescales used in the present experiments and that reactions of SO4-may, therefore, be neglected. In addition to the small yield of H generated in the primary radiolytic event, reaction (5), a further small yield is formed by reaction (12) in competition with reaction (6).e-(aq) + H+ 4H (12) At pH 3.0 the total yield of H is calculated to be 21% of the yield of radicals and in the presence of 0, at the concentra- tion used in this study H will be converted to HOz with t1,, x 50 ns. HO, reacts with Mn2+(aq) with a rate constant of 6.0 x lo6 dm3 mol-' s-l.13 Thus at the highest concentra- tion of Mn2+ used in this study (500 pmol dm-3) this reac- tion is 100 times slower than reaction (2). Results and Discussion Preliminary Observations Fig.1 shows a typical kinetic trace at 470 nm after irradiation of Mn"/SIV solution. The rapid increase in absorbance fol- lowed by a somewhat slower decrease is probably due to the formation of 0,-and its protonated form HO, generated in reactions (13) and (14), where reaction (13) is in competition with reaction (Is)? 0-+ 0, -,0,- (13) 0,-+ H+(H,O)eHO,(+OH-) (14) 0-+ H+(H,O) -+ OH(+OH-) (15) We have observed similar long-lived absorptions at both pH 6.0 and 3.0 in pulse-irradiated water containing both N,O and O,, but they are not present in the absence of 0,.14 The rate constants for reactions (12) and (14) indicate that the yield of 0,-is expected to be only 2.9% of that of OH, but 0,-absorbs strongly at 470 nm.1s.16 Further evidence that 0-can react with solutes in systems at low pH is provided by the work of Zehavi and Rabani.l7 time Fig. 1 Typical kinetic trace at 470 nm obtained by pulse radiolysis of an aqueous solution of Mn" and S'", half-saturated with molecular oxygen and nitrous oxide. Conditions: [Mn"] = 100 pmol dm-3, [S'"] = 800 pmol dmP3, pH 3.0, I = 10 mmol dm-3, dose = 17 Gy. t Note added in proof: In recent experiments we have shown that protonation of 0,-at pH 3.0 is too rapid to account for the 'spike' in Fig. 1 which, at present, we attribute to transient absorption induced in the cell. J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 200 ps H IIIIIIII time Fig. 2 Decay of the Mn"' product at 440 nm.Conditions: [Mn"] = 100 pmol dm-3, [StV] = 200 pmol dm-3, pH 3.0, I = 10 mmol dm-3, dose = 30 Gy. The first-order increase in absorbance to a plateau value is attributed to formation of Mn"'. Under certain experimental conditions, see below, a slow decrease in absorbance could be observed following the build-up of Mn'", Fig. 2. Both reactions were studied in the wavelength range 440 to 470 nm at pH 3.0 under the experimental conditions 50 < [Mn"]/pmol dmw3< 500, 200 < [SIV]/pmol dm-j < 800 and I = 10 mmol dm-3. Doses in the range 17-41 Gy were used, corresponding to an initial concentration of SO, -in the range 10-25 pmol dm -'. Spectra Fig. 3 shows spectra recorded at 2, 20 and 170 ys after irra- diation of a mol dm-' Mn" and 2 x mol dm-3 S" solution and presented as GE-values, i.e. the product of 61 I I I I I I 5 -,., 4-1, c 0 I -I 250 350 450 550 wavelength/nm Fig.3 Spectra recorded at (-) 2, (---) 20 and (. . -.) 170 ps after irradiation and, for comparison, (-. the spectrum of SO, -adjust-0-) ed to G(SO,-)= 5.0 x rnol J-'. Conditions: [Mn"] = 100 pmol dm-3, [S"] = 200 pmol dm-j, pH 3.0, I = 10 mmol dmb3, do= = 19 Gy. the radiation chemical yield, G, of the light absorbing species in units of mol J-and E, the species' absorptivity in units of m2 mol-'. After 20 ,us the formation of manganese(II1) is com- plete and the spectrum shows a maximum at ca. 450 nm. The spectrum recorded after 170 ps shows that the absorbance has decreased over the spectral region from 420 to 600 nm owing to the subsequent, slower reaction.Also shown in Fig. 3 is the spectrum of SO, -expressed in terms of GEbased on G(S0,-) = 5.0 x lo-' mol J-I, the expected yield of SO,-in this system. Additional spectra of the manganese(@ product were recorded at the plateau of the kinetic traces, cf: Fig. 1, for solutions containing different [Mn"] and [S"]. All exhibit a maximum at 260 nm and most of the spectra have an addi- tional, but smaller, absorption in the wavelength range 350-600 nm with a maximum in the region 440-470 nm which is characteristic of many Mn"' absorption spectra previously reported." The extent of absorption in the region 440-470 nm is dependent on the [S'"] : [Mn"] ratio.The absorption decreases as the ratio increases, Table 1, and assuming G(Mn"') = 5.0 x lo-' rnol J-', the effective molar absorp- tivity varies from 1.0to 4.2 mz mol-'. This effect may be due to formation of different Mn"' complexes, e.g. hydrolysed manganese(m), [Mn(HS03)lZ and [Mn(S03)Mn13 +,under+ different experimental conditions, see discussion below and ref. 1. These complexes most likely have different absorp- tivities in this wavelength region, but the values of &:Ao lie in the range previously observed for complexes of Mn"'. 19,20 The apparent decrease in absorbance change on the forma- tion of manganese(In), when the [S"] : [Mn"] ratio increases, may also be explained by an increase in the rates of the sub- sequent reactions when the SrV concentration is increased.Kinetics The growth of absorbance due to formation of Mn"' was analysed according to the pseudo-first-order rate law (I). d[Mn"']jdt = k,b,[SO,-] (1) A first-order correction for the subsequent slow decay was applied when necessary. The value of the observed rate con- stant decreases with increasing concentration of Mn" using both 19 and 41 Gy pulses, Table 2 and Fig. 4. There is also a small dependence on the S" concentration. The values of kobs are slightly larger using [S"'] = 800 pmol dm-3 than using [S"] = 200 ymol dm-3, Table 2. kobs is slightly dependent on dose but is independent of wavelength, within experimental error. Expressed as bimolecular rate constants the values range from 2 x lo8 to 2 x 10" dm3 mol-' s-'.It is sug-gested that the variation in these values is due to a change in the nature of the Mn" species with [S"] and [Mn"] (see below). Table 1 GE values at 470 nm for the Mn"' product determined at different concentrations of S"' and Mn" cs'vl /pol dm-3 [Mn"]/pmol dm-3 [S"] : [Mn"] GE470 mz J-' 200 500 0.4 : 1 2.1" 200 200 1:l 1.Sb 800 500 1.6 : 1 1.5" 200 100 2:l 1.4" 800 200 4: 1 1.2b 800 100 8:l 0.90" 5000 500 10: 1 1.o" 5000 100 50 : 1 0.50" " 19 Gy pulses were used at pH 3.0. 41 Gy pulses were used at pH 3.0. I H i I I I I I I 12 r I ma 0 4 0' I I I I I I0 200 400 600 [Mn "]/pmol dm-3 Fig. 4 Observed pseudo-first-order rate constant as a function of excess concentration of Mn".Conditions: see Table 2. [S"] = (a)200 and (b) 800 pmol dm-3, dose = 19 Gy (closed symbols) and 41 Gy (open symbols). The errors are given as the standard deviations of the mean calculated from four to six measurements. Mechanism Recently, it was shown that a complex between Mn" and hydrogen sulfite with stability constant bl = 3 x lo4 dm3 mol-' is formed in the interval 2.4 < pH < 4.0.' At pH 4.0, Table 2 Observed pseudo-first-order rate constants at different con- centrations of S" and Mn" 4,,~105 s-1 CS'"1 [Mn"] /pmol dm-j /pmol dm-3 dose = 19 Gy dose = 41 Gy 200 50 3.4 k 0.5 4.1 & 0.8 200 100 3.0 f 0.3 3.6 k0.7 200 200 2.0 f 0.1 2.1 & 0.4 200 300 1.6 f 0.1 1.5 & 0.1 200 400 1.0 f 0.1 1.3 & 0.1 200 500 1.1 * 0.1 1.5 f0.1 800 50 9.4 f0.1 11.0 f2.5 800 100 9.6 f0.1 8.0 f0.7 800 200 7.7 & 0.3 7.3 f1.0 800 300 6.3 f0.1 5.8 f0.6 8W 400 1.9 & 0.1 2.1 f 0.1 80% 500 2.8 _+ 0.1 The errors are grven as the standard deviations of the mean calcu- lated from four io six measurements.Conditions: pH 3, Z = 0.1 mol dm-3. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 formation of a bridged Mn" complex, [Mn(S03)Mn12+ was also indicated.' Consequently, the distribution of Mn" between various complex species in the present study is gov- erned by the total concentrations of SrV and Mn". The total concentration of Mn" can be written according to eqn. (11) where #Il2 denotes the stability constant of [MnS03Mn12+.[Mnllltotai= [Mn2'1(1 #Ii[HsO3-1 + B12CMn2+ICHS03-I/CH+1) (11) The dependence of kobs on the concentrations of Mn" and STv may be rationalized qualitatively by use of the following mechanism, reactions (16)-( 18). Mn2+ + SO5----+ H+ Mn3+ + HS0,-(16) H+[Mn(HSO,)]+ + SO,--[Mn(HSO3)l2++ HSO,-(17) [Mn(S03)Mn12++ SO,-H+ [Mn(S03)Mn13---+ + + HS05-(18) The observed rate constant of eqn. (I) will be the sum of the contributions from these three parallel reactions with rate constants k16, k17 and k18 and can be written as eqn. (111). kobs = k16[Mn2+] 4-k,,[Mn(HSO,)+] 4-k18[Mn(SO3)Mn2+] (111) As the concentration of Mn" is increased, the ratio [Mn(S03)Mn2 +]/[Mn"],,,,, increases according to eqn.(11) and thus, the rate of reaction decreases provided that k18 < k16, k17. The small increase in the value of kobs observed for increasing S'" concentrations indicates that k, 7 is slightly larger than k16 since a change of the total s" concentration only influences the ratio between [Mn2'] and [Mn(HSO,)+]. Based on a mechanism proposed by Jayson et al. for the reaction between Fe2+ and H0, Cabelli and BielskiI3 have suggested the following mechanism for the reaction between Mn" and the perhydroxyl radical: Mn2+(aq)+ HO, -+ [Mn(OOH)I2+ (19) [Mn(OOH)]2' + Mn2+(aq)+ [Mn(0OH)Mnl4+ (20) [Mn(OOH)I2+-+ Mn3+(aq)+ H0,-(21) [Mn(OOH)MnI4+ + Mn3+(aq)+ products (22) However, in the mechanism proposed by Jayson et al. for the Fe2+/H02 system, Fe2+ is oxidized to Fe3+ by HO,.After electron transfer, the reaction products persist together as an outer-sphere successor, reaction (23). Fe2+ + HO, -+Fe3+, H0,- (23) The successor complex is in equilibrium with a bridged species, [Fe11'H0,Fe"]4+. Both complexes decompose to give Fe3+ and H02-and Fe3+, Fe2+ and HO,-, respec-tively. Thus, reaction (23) is the only redox reaction in the mechanism, contrary to what is claimed by Cabelli and Bielski. Generation of [Mn(00H)I2 + followed by formation of [Mn(00H)MnI4+ according to reactions (19) and (20) does not seem likely since intramolecular electron transfer in [Mn(OOH)I2' is expected to be very fast. Note that in [Mn(OOH)]'+ a radical is coordinated to Mn", while in Fe3+, H0,-it is a molecular ligand that interacts with the oxidized metal centre.The mechanism proposed in the present study rationalizes qualitatively the experimental results without introducing J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 any complex formation between Mn2+ and SO,-. It should be emphasized that a mechanism analogous to the one postu- lated by Cabelli and Bielski cannot explain why the value of the observed rate constant decreases for increasing concen- trations of Mn" unless [Mn(S0,)Mn13 decomposes more + slowly than [Mn(SO,)] +.This implication would be in sharp contrast to the conclusion made by Cabelli and Bielski that [Mn(00H)MnI4+ decomposes considerably faster than [Mn(00H)] ' .+ Subsequent Slow Processes By use of 41 Gy pulses, the decrease in absorbance following the formation of Mn"' was observed using 50 < [Mn"]/pmol dm-' < 500 and [S"] = 200 and 800 pmol dm-3.However, at lower doses, CQ. 19 Gy, this decay was observed only at low Mn" concentrations, i.e. 50 and 100 pmol dm-3. At a constant Mn" concentration (1 x mol dm-3) the decay rate slows down with increasing sulfur(1v) concentration. At low doses, the absorption decays according to simple first- order kinetics while at higher doses, the decays are best described by a second-order process. It was not feasible to rationalize these slow processes quantitatively. However, it should be emphasized that generation of Mn"' according to reactions (16)-( 18) triggers the manganese-catalysed autoxi- dation of S*.The Mn'" species all react further with different rates and form sulfite radicals, reactions (24)-(26). These rad- icals react with molecular oxygen reproducing SO -radicals, reaction (9). Mn3+(aq)+ [Mn(HSO,)]+ 2Mn2+(aq)+ SO3-+ Hf (24) [Mn(HSO3)l2+-+ Mn2+(aq)+ SO3-+ H+ (25) [Mn(S03)Mn13+-+ 2Mn2+(aq)+ SO3-(26) The peroxomonosulfate radical then reoxidizes Mn" to Mn"' and so on. On the short timescales used in the present study, steady-state conditions have not been established. The pro- cesses following the initial formation of Mn"' are therefore complex. In addition to reactions (24)-(26) dispro-portionation of Mn"' according to reaction (27) is also pos- sible. 2Mn3+(aq)+ 2H20eMn02+ 4H+ + Mn'+(aq) (27) Changing the dose and the initial concentrations of Mn" and S'" may therefore favour different pathways for the decay of M n"'.Conclusion The present study clearly demonstrates that Mn" is oxidized to Mn"' by SO,-radicals and that this is a very fast reaction. The second-order rate constants are approximately 2 x lo8-2 x 10" dm3 mol-' s-' depending on which of the Mn" complexes is predominant in solution. This result gives con- clusive support to the suggested mechanism for the manganese-catalysed autoxidation of SIV. The system is complicated owing to formation of sulfito complexes of Mn" and the initiation of the autoxidation of SIV.Further informa- tion about the kinetics and mechanisms for reactions with peroxomonosulfate radicals and Mn" may be provided by a laser flash photolysis study currently planned at Cookridge Radiation Research Centre.This technique allows the rad- icals to be generated in the absence of StVwhich simplifies the observed kinetics. Sulfite radicals are produced by homolytic cleavage of dithionate and the SO,-radicals are formed by subsequent reaction of the sulfite radicals with molecular oxygen according to reaction (9). Financial support from the Commission of the European Communities within the STEP research program (contract STEP-005-C), from the Swedish Natural Science Research Council and from the Royal Physiographic Society of Lund is gratefully acknowledged. Prof. Sture Fronaeus is acknow- ledged for valuable comments.References 1 J. Berglund, S. Fronaeus and L. I. Elding, inorg. Chem., 1993,32, 4527. 2 R. van Eldik, N. Coichev, K. Bal Reddy and A. Gerhard, Ber. Bunsenges. Phys. Chem., 1992, %, 478. 3 C. Brandt, I. Fabian and R. van Eldik, Znorg. Chem., 1994, 33, 687. 4 G. A. Salmon and A. G. Sykes, Methods Enzymol., 1993, 227, 522. 5 E. M. Fielden, The Study of Fast Processes and Transient Species by Electron Pulse Radiolysis, ed. J. H. Baxendale and F. Busi, Reidel, Dordrecht, 1982, pp. 49-62. 6 F. Wilkinson, CRRC internal report, 1988. 7 A. R. Curtis and W. P. Sweetenham, FACSZMILEICHEKMAT User's Manual, Harwell Laboratory 1988. 8 G. V. Buxton, C. L. Greenstock, W. P. Helman and A. B. Ross, J. Phys. Chem. Ref: Data, 1988, 17, 513. 9 G. V. Buxton, G. A. Salmon and N. D. Wood, in Physico-Chemical Behauiour of Atmospheric Pollutants, ed. G. Restelli and G. Angeletti, Kluwer Academic, Dordrecht, 1990, pp. 245-250. 10 W. J. McElroy, Atmos. Environ., 1986,ZO 323. 11 S. McGowan, Ph.D. Thesis, Leeds, 1994. 12 R. E. Huie and P. Neta, Atmos. Environ., 1987,21, 1743. 13 D. E. Cabelli and B. H. J. Bielski, J. Phys. Chem., 1984,88,6291. 14 G. V. Buxton, G. A. Salmon and J. E. Williams, to be published. 15 W. D. Felix, B. L. Gall and L. M. Dorfman, J. Phys. Chem., 1967, 71, 384. 16 R. E. Buhler, J. Staehelin and J. Hoigne, J. Phys. Chem., 1984, 88,2560. 17 D. Zehavi and J. Rabani, J. Phys. Chem., 1971,67,701. 18 G. Davies, Coord. Chem. Rev., 1969,4, 199. 19 V. P. Goncharik, L. P. Tikhonova and K. B. Yatsimirskii, Russ. J. Znorg. Chem., 1973,18,658. 20 J. P. Fackler and I. D. Chawla, Znorg. Chem., 1964,3, 1130. 21 G. G. Jayson, B. J. Parsons and A. J. Swallow, J. Chem. Soc., Faraday Trans. i,1973,69,236. Paper 4/01866D; Received 29th March, 1994

ISSN:0956-5000    

DOI:10.1039/FT9949003309

出版商:RSC    

年代:1994

数据来源: RSC