Anal. Chem. 1990, 62, 1010-1015
1010
(7) Fuoss. R. M.; Hsia, K. L. J . Am. Chem. SOC. 1968, 90, 3055-3060. (8) Lichtin, N. N.; Bartlett, P. D. J . Am. Chem. SOC. 1951, 73, 5530-5536. (9) Harned, H. S.;Embree, N. D. J . Am. Chem. SOC. 1934, 56. 1042-1044. (IO) Harned, H. S.; Ehlers, R . W. J . Am. Chem. SOC. 1933, 55, 652-656. (11) Nims, L. F. J . Am. Chem. SOC.1933, 55, 1946-1951. (12) Handbook of ChemistryandPhysics. 63rd ed.; Weast, R. C., Ed.; The Chemical Rubber Co.: Cleveland, OH, 1982; p D-174. (13) Pimentei, G. C.; Spratiey, R. D. Chemical Bonding Clsrified through Quantum Mechanics ; Holden-Day, Inc.: San Francisco, CA, 1969. (14) Salem, L. Electrons in Chemical Reactions: First Principles; Wiley-Interscience: New York, 1982. (15) Coplan. M. A.; Fuoss, R. M. J . fhys. Chem. 1964, 68. 1181-1185. (16) D'Aprano, A.; Goffredi, M.; Triolo, R. J . Chem. Soc.. Faraday Trans. 7 1976, 72,79-84. (17) Debye, P. Polar Molecules; Dover Publication: New York, 1965. (18) Charlot, G.; Tremillon, 6. Les reactions chimiques dans les solvants et les sels fondus; Gauthier-Villars Editeur: Paris, 1963. (19) Froiich. H. Theory of Dielectrics; Clarendon Press: Oxford, U.K., 1958. (20) Hyne, J. B. J . Am. Chem. SOC. 1963, 85. 304-306.
(21) Bodenseh. H. K.; Rarnsey, J. B. J . fhys. Chem. 1963, 67,140-143. (22) Accascina, F.; Petrucci, S.; Fuoss, R. M. J . Am. Chem. SOC. 1959, 87, 1301-1305. (23) Kolthoff, I.M.; Bruckenstein, S. J . Am. Chem. SOC. 1956, 78, 1-9. (24) Lichtin, N. N.; Leftin, H. P. J . fhys. Chem. 1956, 6 0 , 160-163. (25) Lichtin, N. N.: Rao, K. N. J . Am. Chem. SOC. 1961, 83, 2417-2424. (26) Time Series frocessor - T S f - User's Guide; Hall, Bronwyn H., Ed.; TSP International: Stanford, CA. 1983. (27) Preti, C.; Tosi, G. Anal. Chem. 1981, 5 3 , 48-51. (28) Preti, C.; Tassi, L.; Tosi, G. Anal. Chem. 1962, 5 4 , 796-799. (29) Franchini. G. C.; Preti, C.; Tassi, L.; Tosi, G Anal. Chem. 1968, 6 0 , 2358-2364. (30) Franchini. G. C.; Marchetti, A.; Preti, C.; Tassi, L.; Tosi, G. Anal. Chem. 1989, 67, 177-184. (31) Van Meurs. N.; Dahmen, E. A . M. F. Anal. Chim. Acta 1959, 27, 10-16.
RECEIVED for review August 24, 1989. Accepted February 2, 1990. The Minister0 della Pubblica Istruzione (M.P.1,) of Italy is acknowledged for the financial support.
Electrolyte Dropping Electrode Polarographic Studies. Solvent Effect on Stability of Crown Ether Complexes of Alkali-Metal Cations ZdenBk Samec*,' and Paolo Papoff Istituto d i Chimica Analitica Strumentale del CNR, Via Risorgimento 35, 56100 Pisa, Italy
Polarography with the electrolyte dropping electrode (EDE) was pioneered by Koryta et al. (1, 2 ) . In the further development, the four-electrode assembly was designed ( 3 ) and employed without a substantial modification in double layer studies ( 4 , 5 )or in measurements of the ion transport across liquid-liquid interfaces (6-14). Fundamental factors in polarography with the EDE were examined (11)for both possible configurations, i.e. for the electrolyte solution dropping upward, called by some authors (6, 7) the ascending solvent (e.g. water) electrode, and for the electrolyte solution dropping downward. In addition to the classical potential-scan polarography ( I , 2 ) , current-scan polarography has been intro-
duced (6) and used in a number of studies (7-14). The advantage of this method is an easy compensation of the ohmic potential drop. On the other hand, the enormous current density a t the very beginning of the drop life can cause irreversible changes in the boundary state and conditions (15). In either case, the cell and apparatus deserve an experienced and skillful experimental approach. The observation (16) that macrocyclic polyethers form stable complexes with alkali and alkaline earth metal cations has stimulated a great deal of interest in these compounds for their possible applications in various branches of chemistry and biology (17). Extensive thermodynamic data (18, 19) suggest that the stability of macrocyclic complexes depends on the relative cation and ligand cavity size, the number and spacial arrangements of the ligand binding sites, the substitution on the macrocyclic ring, and the solvent effects. While these data refer mostly to studies in homogeneous media, an analysis of the interfacial complex formation is obviously of a closely related significance. It has been shown (9, 10,20-25) that a powerful insight into its mechanism, kinetics, and thermodynamics can be gained from measurements of the faradaic ion transport across a liquid-liquid interface in the presence of the ligand. However, systematic electrochemical studies of the factors listed above are lacking. In the present paper we intend to demonstrate the use of a simple three-electrode assembly with the EDE for the study of the interfacial ion transport. In particular, wer focus on the transfer of alkali-metal cations from water to nitrobenzene or 1,2-dichloroethane facilitated by dibenzo crown ethers. We wish to throw some light on the stability of crown ether complexes and the selectivity of ion-ligand interactions in various organic solvents.
'Present address: The J. Heyrovskg Institute of Physical Chemistry and Electrochemistry, Czechoslovak Academy of Sciences, DolejSkova 3, 182 23 Prague 8, Czechoslovakia.
Chemicals. Reagent grade chemicals from Carlo Erba, 1,2dichloroethane (1,2-DCE),nitrobenzene (NB),LiC1, NaC1, KCl,
A simple three-electrode polarographic assembly with the electrolyte dropplng electrode for the study of the Ion transport across a IIquM-Hquid Interface was presented. Transfer of alkall-metal catlons from water to nltrobenzene (NB) or 1,2-dlddwoethane (1,2-DCE) facimatd by complex formation with dlbenzo-18trown-6, dlbenzo-24trown-8, or dlbenzo30-crown-10 was studied. Experimental current vs potential data were used to clarlfy the mechanism of the ion transport and to evaluate the stablllty constants of crown ether complexes. The stabNlty constants are 105-107 larger in NB and 10'o-10'2 larger In 1,P-DCE than those In water. The selectlvlty sequence changes with the solvent, comparing water (K' > Na' > Rb' > Cs') wlth NB (Na' > K' > Rb' > Cs', )'lL and 1,2-DCE (Ll' > Na' > K' > Rb' > Cs'). Solvent effects can be understood In terms of dlfferences in cation solvatlon, which plays a dominant role In low polar media.
0003-2700/90/0362-10 10$02.50/0
EXPERIMENTAL SECTION
IC 1990 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 62, NO. 10, MAY 15, 1990
1011
A
06
PTFE
B
Flguro 1. Three-electrode polarographic cell with the electrolyte
dropping electrode: (WE) largaarea Ag/AgCI working electrode, (RE) Ag/AgCI reference electrode, (CE) platinum wire counter electrode. RbCl, and CsC1, or Fluka AG, dibenzo-18-crown-6 (B218C6), dibenzo-24-crown-8 (B224C8), tetraphenylarsonium chloride (TPAsCl), or tetramethylammonium bromide (TMABr), were used as received. Tetraphenylarsonium 3,3'-corno-bis(undecahydro-l,2-dicarba-3-cobalta-closo-dodecabor)ate (TPAsDCC) and dibenzo-30-crown-10(B,30ClO) were generously presented by Dr. K. Bage, Institute of Inorganic Chemistry, and Dr. I. Stibor, Institute of Organic Chemistry and Biochemistry, Czechoslovak Academy of Sciences, respectively. Water was triply distilled. Electrolytic Cell. Figure 1shows the scheme of the electrolytic cell. The aqueous phase (density 1.051 g cm") was dropping upward into 1,2-DCE (density 1.242 g cm") or NB (density 1.205 g ~ m - from ~ ) a PTFE capillary (inner diameter 0.02 cm) with the flow rate u = (1.3-1.6) X lop3g s-l and the drop time t l = 4-6 s. The column height of the aqueous electrolyte solution (42 cm to the orifice of the capillary) and the volumes of the organic and the aqueous electrolyte solutions in the cell were held constant. The couple of the reference and the counter electrodes for the aqueous phase in the original four-electrode assembly with the EDE (3) was replaced by a single reference Ag/AgCl electrode with an area of about 5 cm2. The electrode was connected to the input for the working electrode (WE) of the conventional three-electrode potentiostat, i.e. to the input of the current follower which was held at the virtual ground. Owing to the low current density, the polarization of this reference electrode was negligible (less than 5 mV). The reference Ag/AgCl electrode for the organic solvent phase (RE) was dipped into the 0.01 M TPAsCl aqueous agar-agar gel solution in a Pasteur pipet, the tip of which was about 0.2 cm far from the drop. The counter electrode (CE) was a platinum wire wound round the drop as shown in Figure 1. Apparatus. The potential difference E = @AB - @' across the galvanic cell (l), where MI = Li, Na, K, Rb, or Cs ( x = 0.01-1.0)
(WE)
(RE)
and L = B218C6,B224C8,or B230C10(y = 0.1-1.0), was controlled by means of a conventional three-electrode potentiostat. The cell potential difference E can be expressed as (2) E = Ab - Abref where A@ = @" - 9" is the Galvani potential difference between water (w) and the organic solvent (0)and A@refis the sum of all other potential differences involved in the cell (1). Since the latter potential difference is a constant, A@ is controlled in the defined way. The polarographic currents were recorded at a scan rate of 2 mV s-l. All measurements were carried out at ambient temperature, i.e. 25 f 2 "C.
10
-
1
-
I 0
I
I
1 0.6
I 20,
C
l/@P
0
+
Flgure 2. Potential-scan polarograms: (A) 0.01 M L i i l 0.5 M MgSO, in water and 0.01 M TPAsDCC in 1,P-DCE;(B) 0.01 M LiCl 0.5 M MgSO, + 0.825 mM TMABr in water and 0.01 M TPAsDCC in 1,P-DCE; and (C) 0.01 M NaCl 0.5 M MgSO, in water and 0.01 M TPAsDCC + 1 mM B224C8 in 1,2-DCE.
+
+
Convention. By definition, the electrical current connected with the transfer of a cation from water to the organic solvent was regarded as positive. RESULTS AND DISCUSSION Potential-Scan Polarography. Figure 2 illustrates the polarographic behavior of various ion transfer systems (base electrolytes, TMA+, Na+ in the presence of B224C8). Faradaic currents of base electrolytes can yield information on thermodynamics of base electrolyte ion transfer. In particular, data on the alkali-metal cations M+ are of significance for the thermodynamic analysis of the interfacial transfer involving M+ and a ligand L. TMA+ was examined here as a test ion, the behavior of which is well understood (26-29). The electrical current at the positive potential limit of polarograms of base electrolytes (Figure 2, curve A) is con-
1012
ANALYTICAL CHEMISTRY, VOL. 62, NO. 10, MAY 15, 1990 I
I
I
I
I
I = rnw(i) Acw(i) = 402 Dw(i)1/2v2/3t1/6p-2/3 Acw(i) (7)
I
I
where i denotes TMA+, rnw(i)is the mass-transport coefficient, Acw(i)= c"(i) - c(i) is the difference between the bulk and surface ion concentration in the aqueous phase (in mol L-'), Dw(i)is the ion diffusion coefficient in the aqueous phase (in cm2 s-l), u is the aqueous solution flow rate (in g s-l), t is the time (in s), and p is the density of the flowing i.e. the aqueous phase (in g ~ m - ~ The ) . value of the mass-transport coefficient rnw(i)= 19.2 mA M-' was calculated from eq 7 for Dw(i)= 9.5 X lo4 cm2 s-l (26), u = 1.36 X g s-', t = 5 s, and p = 1.051 g ~ m - The ~ . result agrees well with the value of 19.4 mA M-' calculated from the experimental limiting current Ilim = mw(i)co(i). The experimental log Z vs log t plots at various potentials have slopes that are close to 1/6, as expected from eq I. E l j 2in eq 6 can be expressed as a sum of the standard potential Eo(i) for the ion transfer reaction
/'
Cs' I
I
Rb' I 1
K'
I
Li
Na' I
I
1
1
1
1
1.5
2.0 ( r , +d,)-':
nm"
Flgure 3. Plot of the standard Gibbs energy of transfer from water to NB (0)and 1,PDCE (0)for alkalhtal cations vs the reciprocal radius of the hydrated ion estimated as the sum of the Pauling radius r , and the diameter of water molecule d , = 0.31 nm.
nected with the transfer of either the cation M+ to the organic phase or the DCC- anion to water. The current at the negative limit is due to the transfer of either the anion Cl- to the organic phase or the cation TPAs+ to water. In either case, the onset of the faradaic current is governed by the standard Gibbs energy of transfer AG,,O of this ion from water to the organic solvent (2). This quantity determines the standard potential difference A + O of the ion transfer reaction
A4' = AGtro/zF
(3)
where z is the ionic charge, or the standard potential Eo in the present potential scale (cf. eq 2) Eo = AcPO - A&f (4) Standard Gibbs energies AGotr(M)for the transfer of alkali-metal cations from water to 1,2-DCE or NB were derived from the positive polarization limit of polarograms (27). The plots of ACtr0(M) vs the reciprocal radius of hydrated ions are displayed in Figure 3. The Li+ and Na+ ions are too hydrophilic and their transfer functions to 1,2-DCE cannot be derived from polarographic measurements (27). The values AGtro(M)= 83 kJ mol-' and 65 kJ mol-', respectively, were estimated by a linear extrapolation along the plot shown in Figure 3. The analysis of the TMA+ ion transfer from water to 1,2DCE (Figure 2, curve B)suggests that the process is reversible and controlled by diffusion of the TMA+ ion. Owing to the ion association in the organic solvent, the overall reaction can be written as (27-29) TMA+(w)
+ DCC-(o)
= TMADCC(o)
(5)
In such a case, the equation for the reversible anodic polarographic wave should apply (28)
E = E l j 2+ ( R T / F ) In I(Zli, -
(6)
where El,* is the reversible half-wave potential, Z is the electrical current, and Zlim is the limiting electrical current. In fact, the plot of log Z(Zlim - Z)-l vs E is a straight line with the slope of 60 f 2 mV, giving El12= 0.515 or 0.413 V for the transfer to 1,2-DCE or NB, respectively. The instantaneous diffusion current should be governed by the IlkoviE equation (30)
TMA+(w) = TMA+(o) (8) and terms reflecting the ion association in the organic solvent and other contributions (28). On this basis, cf. also ref 27, the standard potentials Eo(i) = 0.581 and 0.405 V were derived from the present experimental data for the TMA+ transfer to 1,2-DCE or NB, respectively. Equation 4 was then used to evaluate the reference potential difference Adref= -0.421 and -0.370 V from the values of Eo(i) above and from the known standard potential differences A4"(i) = 0.160 (29) and 0.035 V (2) for the TMA+ ion transfer to 1,2-DCE and NB, respectively. Once the reference potential difference is known, A+' for any charge transfer reaction can be given in the present potential scale. In particular, the standard potentials Eo(M) for the alkali-metal cations were calculated from the transfer Gibbs energies AGtro(M)by using eqs 3 and 4; cf. Table I. The single polarographic wave WBS observed in the presence of a crown ether in the organic solvent phase (Figure 2, curve C). Faradaic currents can be ascribed to the overall reaction (22) r M+(w) + s L(o) = MrLsr+(o) (9) the product of which is the complex of the ion M+ and the neutral ligand L, having the stoichiometry r:s (cation to ligand). Obviously, the case with the 1:l stoichiometry is analogous to the interfacial ion association 5. When the metal ion is present in an excess, the charge transfer process should be controlled by diffusion of the ligand in the organic phase, and the instantaneous current Z should be governed by an equation analogous to eq 7 (22)
Z = rno(L)Aco(L) = 402(r/s) Do(L)'iZu2/3t1/6p-2/3 Aco(L) (10) where Aco(L) = co(L)- c(L) is the difference between the bulk and surface ligand concentration in the organic phase. The equation for the anodic wave reads (22) E = E'(M) - ( R T / r F ) x In [~oaw(M)ryo(L)sy~ML)~lrno(ML)mO~L)"] + (RT/rF) In I(Zli, - 1)" (11) where Ko is the stability constant of the complex in the organic solvent, aw(M)is the activity of the alkali-metal cation in the aqueous phase, ys are the activity coefficients, and rno(ML) is the mass-transfer coefficient of the complex in the organic solvent phase rno(ML) = 402r D0(ML)'I2u2/3t1/6p-2/3
(12)
The polarographic data are summarized in Table I. As expected from eq 10, the limiting current Zh was proportional to the bulk concentration co(L) of the crown ether in the
ANALYTICAL CHEMISTRY, VOL. 62, NO. 10, MAY 15, 1990
1013
Table I. Polarographic Data for the Facilitated Ion Transfer from Water to 1,2-Dichloroethane and Nitrobenzene (in Parentheses)"
B218C6 ion
Eo(M),V
mo(L), mA M-'
slope, mV
B224C8 Ellz, V
mo(L), mA M-'
slope, mV
B230C10 .ElIz, V
mo(L), mA M-'
slope, mV
1.28d 19.4 58b 0.61b 18.1 65 0.62 17.3 (0.77) (61) (65)b*c (0.44)b,c Na+ 1.10d 19.1 58 0.54 18.0 60 0.42 15.7 (0.40) (0.72) (61) (0.43) (60) (14.2) (14.0) (12.5) 0.43 18.2 65 K+ 0.90 19.4 58 0.40 15.6 (0.61) (55) (0.31) (66) (0.35) (14.5) (12.6) (14.5) 58 0.45 18.5 58 0.40 16.4 Rb+ 0.83 19.5 (0.57) (14.5) (65) (0.39) (68)' (0.32)c cs+ 0.77 18.7 65 0.37 17.9 (0.53) 'Metal ion concentration 0.01 M unless otherwise indicated. bMetal ion concentration 1 Prl. cEstimated from graphic wave. dAn extrapolated value.
Li+
V
64
0.56
60 (64) 59 (56) 64
0.45 (0.42) 0.28 (0.26) 0.29 (0.28)c 0.32
bvc
organic solvent (0.1-1.0 mM) but independent of the bulk metal ion concentration cw(M)in the aqueous phase (0.01-1.0 M). With the same flow rate and the complex stoichiometry, the ratio of the mass-transfer coefficients rno(L)in 1,2-DCE and NB reflects mainly the ratio of the diffusion coefficients Do(L),e.g. for B218C6 Do(L) = 5.9 x lo4 (10) and 2.3 x lo4 cm2 s- (25),respectively. By comparing the mass-transfer coefficients for the TMA+ and the ligands (Table I), one can estimate the stoichiometric ratio as r / s = [rno(L)/rnW(i)][Dw(i)/Do(L)]1/2 = 1.3 for Dw(i) = 9.5 X lo4 cm2 (26) and the diffusion coefficient of B218C6 in 1,2-DCE Do(L) = 5.9 X lo4 cm2 s-l (IO). Owing to 10-20% uncertainty in D values and variation in the flow rate u from experiment to experiment, cf. the Experimental Section, we conclude that the stoichiometric ratio r / s probably equals unity. The half-wave potential El,: was independent of the ligand concentration, but it varied with cW(M),as illustrated for K+ ion in Figure 4. Equation 11 predicts that the plot of vs log aw(M) should have the slope of -59 mV, which was actually the case. The plots shown in Figure 4 have slopes of -58, -55, or -54 mV for B218C6,B224C8, or B230C10, respectively. Since the logarithmic analyses, i.e. the E vs log Z(Zli, - Z)-l plots, were straight lines with slopes of about 59 mV (Table I), the stoichiometric coefficient s should equal unity. I t can be concluded that a complex with the stoichiometry 1:l is formed in all cases a t the water-1,2-DCE interface. The same stoichiometry has been already evidenced for crown ether complexes of alkali-metal cations a t the water-NB interface (20, 25). Mechanism of Ion Transfer. In the studies of the cation transfer facilitated by a macrocyclic ligand, Koryta et al. (20, 21,23) have suggested the EC mechanism for the ion transfer M+(w) = M+(o)
(134
+
M+(o) L(o) = ML+(o) (13b) On the other hand, Freiser et al. (IO, 24) presented several arguments in favor of two different mechanisms, depending on the concentration of the metal cation. For lower concentrations of the cation, a CE type of mechanism has been proposed
L(0) = L(w)
(14a)
+ L(w) = ML+(w)
(14b)
M+(w)
(14~) ML+(w) = ML+(o) while a t higher concentrations the EC mechansim of eq 13 should prevail, with an intermediate occurrence of both (IO). However, neither of these proposals seems to be well founded. In contrast to the current-scan measurements, which indicated the occurrence of two electrochemical processes (IO), we have observed only single polarographic waves under potentiostatic
63
the foot of the polaro-
Figure 4. Dependence of the reversible half-wave potential on the concentration c wof the K+ cation in the aqueous phase in the pres-
ence of B218C6 (O),B224C8 (O), and B230c10 (A)in 1,P-DCE: aqueous phase, KCI + 0.5 M MgSO,; organic solvent phase, 0.01 M TPAsDCC 1 mM crown ether.
+
conditions over the broad range of metal ion concentrations (0.01-1.0 M). Another argument (IO) for the CE mechanism was that the transition time of chronopotentiograms increased with the period of time the two liquid phases had been held in zero-current contact. But no analogous effect has been observed under the potentiostatic conditions (25). Senda et al. (25) pointed out that the question about the mechanism can only be answered with the help of a kinetic analysis. The conclusion was then made (25) that the most probable mechanism consists in a single electrochemical step described by eq 9. Still there exists a possibility that the single electrochemical step involves the adsorbed ligand (IO),e.g.
M+(w)+ Lads = ML+(o) In fact, the drop-time measurements at various ligand concentrations (Figure 5) point to the strong adsorption of crown ethers at the liquid-liquid interface in the potential range close to the zero-charge potential difference (electrocapillary maximum). At a constant potential, the drop time and, hence, the surface tension decrease as the crown ether concentration increases, which would correspond to an increased surface activity of the ligand. We conclude that the potential-scan polarographic measurements do not permit distinguishing among various mechanisms outlined above. However, these measurements point clearly to the existence of a single electrochemical process, which is faradaic by nature and can be ascribed to the overall reaction 9. The successive occurrence of the faradaic and nonfaradaic (adsorption) processes, which has been
1014
ANALYTICAL CHEMISTRY, VOL. 62, NO. 10, MAY 15, 1990 0 2 A VrV
0
.02
I
r
Table 11. Stability Constants of Crown Ether Complexes in 1,Z-Dichloroethane and Nitrobenzene (in Parentheses) log KO
I
Bz18C6
ion 50-
Li+
11.3 (4.8)
Na+
11.2 (7.3, 6.3O)
~
B224C8
B230C10
13.2 (5.5) 13.3
14.2
(7.7)
(7.5, 6.6") 12.4
12.8
400 10.4 (6.6) 9.4 (6.3) 8.5 (5.4")
9.9
K+
(7.2, 6.0°)
Rb+
8.3
(5.2)
cs+
30
E V
04
02
O b
(8.3) 11.1
(7.3)
9.4
"From ref 20.
Figure 5. Dependence of the drop time t , on the potential E or
+
potential difference, Ap, for 0.01 M LiCl 0.5 M MgSO, in water and 0.01 M TPAsDCC B224C8 at various concentrations of crown ether (mM): 0 (O), 0.25 (W), 0.50 (A),and 1.0 (V).
+
Table 111. Relative Stabilities of B218C6Complexes of K ' and Na+ (in Parentheses) Ions in Various Solvents
solventa DNb DMSO DMF MeOH water
PC MeCN
NB 1,2-DCE
AGwo(M),
kJ mol-'
log K c
29.8 26.6 19.0 18.0
-13.0 (-13.4) -10.3 (-9.6) 9.6 (8.2)
15.1 14.1 4.4 0
5.3 (14.6) 8.1 (15.1) 23 (34) 49 (6sd)
3.43 (3.31) 3.55 (3.34) 5.00 (4.36) 1.6 (1.1) 5.08 (3.87) 4.80 (5.01) 7.2 (7.3) 9.9 (11.2)
0 (0)
JAGo + AG,,O (M)+
kJ mol-'
-23.4 (-26.0) -21.4 (-22.4)
-9.8 (-9.0) 0 (0)
-14.6 (-1.2) -10.2 (-7.2) -9.0 (-1.4) 1.6 (7.4)
Abbreviations: DMSO, dimethylsulfoxide; DMF, N,N-dimethylformamide; MeOH, methanol; PC, 1,2-propylenecarbonate; MeCN, acetonitrile; NB, nitrobenzene; 1,2-DCE, 1,2-dichloroethane. "The Gutmann donor numbers from ref 31. cFrom ref 18 or Table I1 (NB, 1,2-DCE). d A n extrapolated value.
LI' 1
1
cs'
K * Rb*
Na' 1
of Li+, Na+, K+, Rb+, and Cs+, respectively (19). Obviously, KO values are 105-107 larger in NB and 1010-1012larger in 1,Z-DCE than those in water. Second, a remarkable change in the selectivity sequence takes place when NB is replaced by 1,2-DCE (Figure 6). Both effects can be understood in terms of differences in ion solvation. The ratio of stability constants in water and an organic solvent K w / K ocan be related to the standard Gibbs energies of transfer for the ion AG,O(M), ligand AG,O(L), and complex LG,,O(ML) through a thermodynamic cycle (18)
1
4
r, rnm
Figure 6. Selective complexation of alkalimetal cations by B218C6 (0, O), B224C8 (0, W), or B230C10 (A,A)in 1,P-DCE (empty points) or NB (full points): r , is the Pauling ionic radius.
indicated by the current-scan measurements (IO), results probably from different initial conditions of the galvanostatic experiment. In particular, the high current density during the early stage of the drop life can cause an excessive accumulation of an adsorbed intermediate a t the liquid-liquid interface, which is then stripped off from the interface at a potential different from that for the faradaic process. Solvent Effects. Equation 11 was used to evaluate the stability constants KO of crown ether complexes in 1,2-DCE and NB from reversible half-wave potentials El12(Table I). The ion activity coefficients y were estimated with the help of the Debye-Huckel theory as described in a previous communication (27). The stability constants are given in Table I1 and displayed in Figure 6. Two features should be noted. First, the enhancement of the complex stability in low polar solvents is considerable. In water, log P = 0, 1.1,1.6, 1.08, and 0.83 for B218C6complexes
M+(o)
+
L(o)
=
ML+(o)
In this way
-RT In ( K o / K w )= AAGO = AGtro(ML)- AGtro(L)- AGtro(M) (17) A comparison (27) with data for other organic solvents has indicated that Gibbs energies of transfer of alkali-metal cations correlate with the Gutmann donor numbers (31), i.e. the highest transfer energies are found for the transfer to solvents having the lowest donating strength; cf. Table 111. Large and positive values of ion transfer energies seem to be responsible for the high stability of complexed ions in these solvents. In fact, the difference AG,O(ML) - AG,O(L) calculated from the present data for B218C6complexes of Na+ and K+ (Table 111) drops to a low value in 1,2-DCE, so that the last term on the right-hand side of eq 17 dominates. In contrast to NB or l,Z-DCE, the solubility of dibenzo crown ethers in protic solvents like water or methanol is low (32)and the partition C much ~ ~ ~ higher ( L ) than unity (25). coefficient P = C ~ ~ ~ ( L ) / is
Anal. Chern. 1990, 62, 1015-1019
Consequently, the standard Gibbs energies of transfer of these ligands from water to NB or 1,ZDCE are negative, AGtro(L) = -RT In P < 0, and those for complexed ions must be negative too, as one would expect for bulky hydrophobic ions. On the other hand, it is clear from Table I11 that the variations in transfer energies for both the complexed ions and the ligand must be considered in solvents of high polarity. A comparison of the cavity size of B218C6and the diameters of unsolvated ions shows that the optimal spatial fit is reached for the potassium cation (18,19). Actually, B218C6 exhibits the selectivity for K+ in polar solvents like water, methanol, dimethylformamide, or dimethyl sulfoxide, the selectivity sequence being K+ > Na+ > Rb+ > Cs+. However, the positions of Na+ and K+ are reversed in acetonitrile, and the same is true for nitrobenzene (Figure 6). In 1,2-DCE the selectivity sequence is quite different, Li+ > Na+ > K+ > Rb+, and it seems to follow mainly the change in the cation solvation; i.e. the cavity size effect no longer dominates. A similar behavior was found for the other two crown ethers studied (Figure 6). Thus, the results of the present study confirm the trends previously observed (18, 19) but show that in low polar solvents, inclusive of perhaps the hydrocarbon interior of biological membranes and their models, the desolvation and solvation processes play a key role.
LITERATURE CITED (1) Koryta, J.; Vanqsek, P.; Biezina, M. J. Elecfroanal. Chem. Interfaclal Nectrochem. 1078, 6 7 , 263. (2) Koryta, J.; Vanqsek, P.; Biezina. M. J. Electroanal. Chem. Interfacial Electrochem. 1077,75, 211. (3) Samec. 2.; MareEek, V.; Weber, J.; Homolka, D. J. Nectroanal. Chem. Interfacial Nectrochem. 1070,9 9 , 385. (4) Kakiuchi, T.; Senda, M. Bull. Chem. SOC.Jpn. 1083,56, 1322. (5) Kakiuchi, T.; Senda, M. Bull. Chem. SOC.Jpn. 1083,56, 1753. (6) Kihara, S.; Yoshitia, 2.; Fujinaga, T. Bunsekl Kagaku 1082,3 1 , 297. (7) Yoshida, 2.; Freiser, H. J. Nectroanal. Chem. Interfacial Nectrochem. 1084, 162, 307.
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(8) Yoshida, Z.; Freiser, H. Inorg. Chem. 1084,2 3 , 3931. (9) Lin, S.; Freiser, H. J. Electroanal. Chem. Interfacial Elechochem. 1085, 191, 437. (10) Lin, S.; Zhao, 2.; Freiser, H. J. flectroanal. Chem. Interfacial Electrochem. 1088, 210, 137. (11) Kihara, S.; Suzuki, M.; Maeda, K.; Ogura, K.; Umetani, S.; Matsui, M. Anal. Chem. 1088, 58, 2954. (12) Lin. S.; Freiser, H. Anal. Chem. 1087,59, 2834. (13) Yoshida, 2.; Kihara, S. J. flectroanal. Chem. Interfackl flectrochem. 1087,227, 171. (14) Kihara, S ; Suzuki, M.; Sugiyama, M.; Matsui, M. J. Nectroanal. Chem. Interfacial Electrochem. 1088,249, 109. (15) Koryta, J. Nectrochim. Acta 1088,33, 189. (16) Pedersen, C. J. J. Am. Chem. SOC.1067,8 9 , 2495, 7017. (17) Lehn, J. M. Angew. Chem., Int. Ed. Engl. 1088, 27, 89. (18) de Jong, F.; Reinhoudt, D. N. A&. Phys. Urg. Chem. 1080, 17, 279. (19) Izatt, R. M.; Bradshaw, J. S.: Nielsen, S. A,; Lamb, J. D.; Sen, D.; Christensen, J. J. Chem. Rev. lg85,85,271. (20) Hofmanovi, A.; Hung, Le Q.; Khalil, W. J. flectroanal. Chem. Interfacial Electrochem. 1082, 135, 257. (21) Homolka, D.; Hung, le Q.; Hofmanovi, A,; Khaiil, M. W.; Koryta, J.; MareEek, V.; Samec, 2.; Sen, S. K.; Vanfsek, P.; Weber, J.; Biezina, M.; Janda, M.; Stibor, 1. Anal. Chem. 1080,5 2 , 1606. (22) Samec, 2.; Homolka. D.; MareEek, V. J. Electroanal. Chem. Interfacial Electrochem. 1082, 135, 265. (23) Vanqsek, P.; Ruth, W.; Koryta, J. J. Nectroanal. Chem. Interfacial Electrochem. 1083, 148. 117. (24) Yoshida, 2.; Freiser, H. J. Nectroanal. Chem. Interfacial €/echochem. 1084, 179, 31. (25) Kakutani, T.;Nishiwaki, Y.; Osakai, T.;Senda, M. Bull. Chem. Soc. Jpn. 1088,59, 781. (26) Wandiowski, T.; MareEek, V.; Holub, K.; Samec, 2. J. Phys. Chem. 1080,93, 8204. (27) Samec, 2.; MareEek, V.; Colombini, M. P. J. €/echoanal. Chem. Interfacial Electrochem. 1088,257, 147. (28) Makriik, E.; Hung, Le Q. J. flectroanal. Chem. Interfacial flectrochem. 1083, 158, 277. (29) Wandiowski, T.; MareEek, V.; Samec, 2. Nectrochlm. Acta, in press. (30) Heyrovskq, J.; Kiita, J. principles of Poiarography; Publishing House of the Czechoslovak Academy of Sclences: Prague, 1965. (31) Marcus, Y. J. Solution Chem. 1084, 13, 599. (32) Takeda, Y. Bull. Chem. SOC.Jpn. 1083.56, 3600.
RECEIVED for review August 8, 1989. Revised manuscript received January 18, 1990. Accepted February 1, 1990.
Multiple Sensor Response in Segmented Flow Analysis with Ion-Selective Electrodes D. Brynn Hibbert,* Peter W. Alexander,* Sri Rachmawati, and Sylvia A. Caruana Department of Analytical Chemistry, University of New South Wales, P.O. Box I , Kensington, New South Wales, Australia 2033
Improved sensitivity in potentiometric analysis is achieved wlth multlple cells In a continuous flow system. Cells connected In series have been applied previously for Improved sensltlvHy In potentlometric detection, but only when the cell solutions are phydcally Isolated from one another. The novel use of air-segmented flow wlth appropriate cell design Is shown here to glve addltlve cell response, even though the cell electrolyte solutions are connected. A simple theory for two cells In serles Is developed to show that the total cell potentlal Is expected to be Ar$[(2a 2 ) / ( a 2 ) ] ,where Ar$ is the slngle cell potentlal and CY is the ratio of the resistance between sells to that within each cell. Experimentally,the sensltlvlty of a three-cell sensor system for detection of chloride ion is shown to glve up to 3 times the single Nemstlan slope. The detectlon limit In the sub-Nernstian region of the cell response Is 0.7 pM, an improvement of approximately 10 over the single cell system, and peak height measurements show a relative standard deviation of 1.5 YO for determination of a chloride sample of 20 pM concentration.
+
+
The response characteristics of ion selective electrodes in rapid flow air segmented analysis have been reported (I),with flow rates up to 9 mL/min required to produce very rapid electrode response. However, the sensitivity of the electrode detection method is limited by the Nernstian response characteristics of the potentiometric detectors sensitive to ionic species. A monovalent anion, for example, is theoretically detectable with a slope of 59.1 mV a t 25 "C, and an error of 1 mV in reading gives an approximate 4% error in ion concentration due to the logarithmic relationship with measured electrode potential. In flow analysis systems with ion selective electrode sensors, the sensor response is subject to a noise level created by the flow across the sensor surface where the noise level is often more than 1 mV. Hence, there is significant error expected in flow analysis with potentiometric detectors unless the noise level can be reduced or the sensitivity of the sensor response can be improved. Improvement in the sensitivity of potentiometric electrode sensors is therefore desirable in future developments on
0003-2700/90/0362-1015$02.50/00 1990 American Chemical Society
