This device was designed specifically for pesticide cleanup, but we feel that many other chromatographic techniques can be adapted to this system. Any technique that uses only one solvent throughout sample addition and elution could be adapted to this system without significant modifications other than possible changes in column and sample loop dimensions. With the addition of valves and timing circuits, such techniques as ion exchange chromatography and multiple eluate chromatographic techniques could be performed by this system.
The system described here is not suitable for high pump pressures, and special equipment may be necessary where high pressures are required. The basic system design, however, should be satisfactory if appropriate high pressure equipment is available.
RECEIVED for review September 29, 1971. Accepted May 2, 1972.
Response of a Calcium-Selective Electrode in Acid Solutions John Bagg and Robert Vinen Department of Industrial Science, University of Melbourne, 35 Royal Parade, Victoria 3052, Australia The interference by H + ions with a liquid-membrane electrode selective toward Caz+(Orion Research Inc.), which has been in contact with aqueous solutions for 4 2 hr, is very similar to the interference by alkalimetal ions. The interference potential, defined as the difference between the observed potential and the potential calculated assuming no interference, is given by: RT
E = - In (1 F
+ K~a~/ac,”’) *
where the selectivity parameter Ka = (1.8 0.3) X 101. If the electrode has had only a brief contact with aqueous solutions, the accepted limit on pH for absence of significant H interference may be extended from 5.5 down to 4.0 when [Caz+] 2 0.1. When electrodes which have been in prolonged contact with aqueous solutions are immersed in solutions of pH 3-5, and calcium concentration in the range 10-1-10-4 molal, an anomalous negative change in potential is observed. Associated with this negative change is the entry of water and calcium salts into the membrane, and a slow process probably due to the formation of a barrier layer. +
THEBEHAVIOR of a liquid-membrane selective toward Ca2+, first described by Ross ( I ) , in mixed CaClz-alkali-metal chloride solutions (without significant interference by H+) has been studied previously ( I , 2) and some progress has been made toward an understanding of the processes involved (2, 3). The emf of the cell, SCEl’CaCl2-MC11 liquid-membrane electrode (Orion Research Inc. type 92-20) was found by the authors to fit the following equation over the range of total ionic strength, I = 0.006-0.600, and for pH 2 5.5 (3): E
=
E,
+ RT In F -
+
KhIa2)
where KM is the selectivity parameter and a1,a2are the singleion activities of Ca2+ and the alkali-metal, respectively [calculated using the convention proposed by Bates (4)]. This equation can be derived by simple theory ( 2 , 3 )and this theory predicts that the emf with a mixed electrolyte should always ( 1 ) J. W. Ross, Jr., Science, 156,1378(1967). (2) J. Bagg, R. Vinen, and 0. Nicholson, J. Plzys. Cliem., 75, 2138 (197 1). (3) J. Sandblom, G. Eisenrnan, and J. L. Walker, ibid., 71, 3862 (1967). (4) R. G. Bates, “Ion-Selective Electrodes,” R. A. Durst, Ed., Nor. Bur. S t u d . (U.S.)Spec. Publ. No. 314(1969).
be more positive than with a pure CaCh solution of the same Caz+ activity. To study the effect of H+, similar emf measurements have been made in CaC12-HCl solutions ( I , 5, 6) but the results show some inconsistencies. In the original paper by Ross ( I ) , emf of cells containing CaCkHC1 mixtures, [Ca2+l, 10-1-10-4m,pH 2-9, were reported and selectivity parameters SH = lo5, given, with SHdefined by the following equation:
E
=
E,
+ RT - In (al + SHa32) 2F
with a3 the single-ion activity of H+. In later work, however, both Ross (5) and Orme (6), using the same mixed solutions, observed more complex behavior than predicted by either Equation 1 or 2. With solutions of pH 3-5, and [Caz+] > 10-3m, the emf became more negutioe with increasing [H+] at a given constant [Ca2+]. The magnitude of this change did not appear to be reproducible and might depend upon the particular electrode construction (6). At lower pH, the emf increased in an approximately Nernstian manner with [H+], as might be expected from Equation 2 when S H>> al. ~ ~ Widely ~ varying selectivity parameters of S H = 2 x lo5(6) and lo7(5)were reported. The only other liquid-membrane electrode which has been studied in solutions of varying pH appears to be the leadselective electrode (7); however, no negative change was detected in solutions of [Pb2+], 10-2-10-4m,pH 2.5-8.0. It is of interest to note that a solid-state electrode selective toward calcium has been reported to show a negative change with pH in a solution [Caz+l = 1.4 x lo-%, pH 3-5, but not at lower [Ca2+] (8). The exact composition of this electrode has not been revealed but may be an organophosphorus compound immobilized in a thin matrix. At present the only explanation offered for the negative change is by Orme, who suggested that some impurity in the liquid-ion exchanger might be responsible (6). In view of the expanding applications of the calcium-selective electrode for measurements in multi-ionic solutions, more information about the influence of pH is clearly desirable. The aim of ( 5 ) J. W. Ross, Jr., ibid., p 57. (6) F. W. Orme, “Glass Micro-Electrodes,” M. Lavallee, Ed., John Wiley & Sons, New York, N.Y., 1969, p 376. (7) S. La1 and G. D. Christian, Anal. Chin?.Acra., 52,41(1970). (8) G. A. Rechnitz and T. M. Hseu, ANAL.CHEM., 41,111 (1969).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 11, SEPTEMBER 1972
1773
-I .
70
40
'1
3 20
E
*-: E,
jOl
a.
20
Y
0
10
b.
-101
C.
Figure 3. Emf us. pH for CaCL-HCI-LiCI mixed solutions, (a) I = 0.6, (6) I = 0.3, (c) I = 0.06. All all cases R = LiCl/ CaClzand curves are labelled with [Ca] from distilled water and analytical grade chemicals. The ratio MX/CaX2 varied from 0.5-150 and the total ionic strength of the solutions from 0.003-0.6. The pH varied from 8-2, governed by the addition of HX or NaOH, and was measured with a glass electrode (E.I.L. Model 23A pH meter). The liquid-membrane, reference electrodes, and solution were contained in a glass cell, electrically-shielded and, normally, maintained at 25 i 0.1 "C. If the effect of stirring was to be studied a small "Teflon" (Registered Trade Name of E.I. du Pont de Nemours) -coated magnetic stirrer at 250 rpm was used, otherwise the solutions were unstirred. The emf of the cell was measured with a Keithley 610C electrometer. Previous work had shown that for pH 2 5.5, stable values of emf were attained within 1-2 min after immersion in fresh solutions and were then reproducible to within 1 rn V over a period of several weeks (2).
----
Figure 1. Emf us. pH, (a)-CaClz, (b) CB(NO~)~, at salt concentrations from 10-e 10-'m
*
RESULTS
a.
20 4
6
8
C.
IO '2
4
P
6
8
'2
4
6
8
Figure 2. Emf 1;s.pH for CaCI2-HCI-NaCI mixed solutions, ( a ) I = 0.6,(6) I = 0.3, (c) I = 0.06. In all cases R = NaCI/ CaCh and curves are labelled with [Ca] the present work is to attempt to gain some of that information by making emf measurements in mixed Ca-H-alkalimetal salt solutions. EXPERIMENTAL
The calcium-selective electrode immersed in a mixed electrolyte may be represented by the following solution chain: SCEI ~CaXz mlP liquid-membrane,CaClz MX mzPl HX m361 ~
~
m?,Ag, AgCl
I
where mla is the molal concentratior, of CaClZin the reference compartment and m?, m t , and m? are the concentrations of the indicated species in the outer solution. X = C1- or NO3- and M = Li' or Na+. The liquid-membrane was an Orion Research, Inc. Type 92-20 (reported to be O.lm calcium salt of a long-chain dialkylphosphoric acid in dioctylphenylphosphonate (DOPP) (5, 9). The exterior saturated calomel reference electrode (SCE) was an E.I.L. Type RJ-23 (porous-plug liquid-junction). All solutions were prepared (9) M. Frant, Orion Research, Inc., Cambridge, Mass., private communication, 1971. 1774
Typical graphs of emf (with SCE) us. pH for pure CaClz and Ca(N03)2solutions of different concentrations are shown in Figure 1. The electrode was fist immersed in 10-*m CaClz at pH = 5.5 for a period of not less than 48 h r ; and then the pH was adjusted to 8.0 by the addition of NaOH, the emf measured after 3-min immersion; HC1 added to reduce pH, the emf measured again after the same period; and this procedure was repeated until pH = 2.0. The shape of the curves was very similar to that observed by previous workers (5, 6) and confirmed the negative change. The magnitude of the change was greatest in O.lm [Ca], smaller at lower concentrations, and was absent at 0.0001rn. Although the magnitude of the change was not very reproducible, it seemed to be greater in nitrate solution than in the chloride solution. Similar emf us. pH measurements were made in CaC12-HC1 solutions but using an Ag,AgCl reference without liquid-junction. The shape of the resulting curves was almost identical to that using SCE with liquid junction. It must, therefore, be concluded that the changes of liquidjunction potential with pH were not great enough to account for the changes in emf. All subsequent measurements reported in this paper were made with SCE. If equilibrium was attained at the solution-membrane interface, the ratio of H/Ca in the membrane surface would be governed by a3/a112. Examination of Figure 1 shows that the negative change is not governed by a3/aI1'*; for example, the minimum in the emf us. pH curve at [CaCl,] = 0 . 1 ~OC1 curs at a 3 / a 1 1 ~= z 5.95 X and at [CaCL] = 0.0001m, this value of ~ 3 / ~ l l is" attained at pH 4.3 but no minimum is detectable, The negative change, therefore, is dependent upon [Ca*+]and not upon an equilibrium ratio.
ANALYTICAL CHEMISTRY, VOL. 44, NO. 11, SEPTEMBER 1972
'4
I
C.
t ""0
'60
'120
'I80
'240
'300
'360
time min.
'420
'
40
d. '0
60
I20
100
240
300
360
420
time min.
Figure 4. Emf US. time for electrode immersed in O.lm CaCI,, (a) stirred solution at 15 "C, ( 6 ) stirred solution at 25 "C, (c) stirred solution at 35 "C, (6) unstirred solution at 25 "C. 4 pH changed 7-3 and t pH changed 3-7
The effect of ionic strength upon the interference by H+ was investigated by measurements upon CaC12-HC1-NaC1 mixed solutions with a varying ratio, R = NaC1/CaCl2, but with I maintained constant for a given series of solutions. Using a previously determined value of K N ( ~ 2 ) , R was chosen to keep the interference by Na+ at a very low level at all ionic strengths. Figures 2a, 26, and 2c show emf us. p H at Z = 0.6, 0.3, and 0 . 0 6 ~ . The values of emf at pH = 6.0 show that, except at Z = 0.6, R = 100, no significant interference by Na+ or H+ occurred, thus confirming previous estimates of Kxa and SH. The results may then be said to have achieved the initial aim of showing the effects of total ionic strength and [Caz+] without the complication of interference by Na+. Figures 2a, 26, and 2c show the greatest negative change occurs at the highest [Ca2+land is very small at the lowest [Ca2+]. irrespective of the total ionic strength. These results show, in addition, that the total chloride concentration is not an important factor. For example, the negative change is very similar at Z = 0.3, R = 10, [Cl-] = 0.28, [Caz?] = 0.023m, and at Z = 0.06, R = 0, [Cl-] = 0.04, [Ca2+] = 0.02m despite the difference in [Cl-]. No negative change was observed at I = 0.6, R = 100, although [Ca2+]was sufficiently great that such a change would be expected from the observations at lower [NajCa] (see Figure 1 and Figure 2c). The absence of negative change would seem to be associated with the entry of Na+ into the membrane. The value of the emf at pH 6.0 under these conditions did show appreciable interference by Na+, in agreement with earlier work (2). To study the effect of interference by alkali-metal ions further, measurements were made in CaC12-HC1-LiC1 solutions. Li+ has been shown to interfere more strongly than any of the other alkali metals (2, IO). Figures 3a, 36, and 3c show emf US. pH for Z = 0.6, 0.3, and 0.06 and, except under the condition of least interference, R = 15, no negative change was observed even at [Ca*+] = 4.6 x 10-2m. Evidently the concurrent entry of Li+ and H+ ions prevents the negative change. (10) J. Bagg and W. P.Chaung, Aust. J. Clzem., 24,1963 (1971).
Figure 5. Emf cs. time for electrode immersed in, (a) O.lm Ca(NO& -, ( 6 ) O.lm CaCI, -,solutions stirred at 25 "C. pH changed 7-3 and t pH changed 3-7
4
----
Under all conditions which produced a negative change, the emf us. pH curve was not retraced upon increasing the pH back to the original high value. After making measurements at decreasing pH, upon increasing the pH a smaller negative value was observed. This hysteresis suggested that equilibrium had not been attained and some slow process might be operating. The change of emf with time was, therefore, investigated after removing an electrode (previously immersed for not less than 48 hr in 10-2rn CaC12, pH 5 . 5 ) from 15-min contact with 0 . 1 ~ CaCl,, pH 7.0, and transferring to a similar solution at pH 3.0. Graphs of emf US. time are shown in Figures 4a, 46, 4c, and 4d at three temperatures, 15,25, and 35 "C, with stirring, and at 25 "C without stirring. The curves in Figures 4a, 40, and 4c show an initial rapid decrease in emf during the first 10 min, and the rate of this decrease is almost independent of temperature. This decrease is followed by a slow rise in emf which is dependent upon temperature. In absence of stirring, the initial decrease takes place more slowly than with stirring, but the subsequent rise in emf is unaffected by stirring. After 2-3 hr immersion at pH 3.0, the electrode was transferred back to pH 7.0 and a rapid rise in emf was observed followed by a slower decrease toward the original value. Again the rapid increase was almost independent of temperature while the slower decrease was temperature-dependent. Emf 1's. time measurements were also made in 0 . 1 Ca(N03)* ~ and are shown in Figure 5. These results confirmed the previous indication that the initial decrease was greater in Ca(N03)* than in CaCh solutions, but showed that the rate of recovery toward an equilibrium value is very similar in both solutions. In addition to emf, the resistance of the liquid-membrane was measured at intervals after transferring from pH 7.0 to 3.0. At temperatures from 15-35 "C, the resistance increased monotonically to a value approximately twice the initial value. Upon returning the pH to 7.0, the resistance decreased close to the initial value with only slight hysteresis. In the final set of experiments, the effect of time of contact with aqueous solutions on the magnitude of the negative change was studied. Measurements were made first on a liquid-membrane prepared from a batch of ion-exchanger which had only a brief contact with CaC12 solution. After 12 hr or less immersion, almost no negative change was detected, but after longer periods of immersion, the magnitude of the negative change increased (Figure 6). The appearance
ANALYTICAL CHEMISTRY, VOL. 44, NO. 11, SEPTEMBER 1972
1775
b.
‘0
60
120
time min.
180
240
Figure 6. Emf us. time for electrode immersed in O.lm CaCI,, (a) 12-hr previous immersion in O.lm %aClzat pH 5.5, ( b ) 30 hr, (c) 80 hr, (6)120 hr. 3. pH changed 5.5-3 of the membrane changed from a clear yellow initially, to a milky opaqueness after long periods of immersion. The resistance of the membrane fell from 6 X lo8 R after 12-hr immersion to 9 X lo7 R at the end of 100 hr. The membrane could be restored to its original clear yellow appearance and high resistance by removal from solution and standing for 24 hr. At the end of this period, upon reimmersion, the small negative change, typical of short immersion times, was also observed. DISCUSSION
KH was calculated at several pH values with [CaC12] = lO-4m, from Equation 1, and, judging by the constancy of the calculated values, shown in Table I, the equation was obeyed satisfactorily with an average KH = (1.8 0.3) X 10’. The results of Ross and Orme at higher concentrations, omitting the regions of negative change, were treated in the above manner and also gave reasonable fit to Equation 1. Table I shows that KH from the present work is in good agreement with the value from Orme but was considerably lower than that of Ross. There have been changes in composition of the ion-exchanger since its inception (9), and from the above results the current electrode appears to be less sensitive to interference by H+ than the earlier version. The results from the CaCl2-HC1-NaC1 mixed solutions at Z = 0.3 and 0.06, R = 50, could also be fitted to Equation 1 giving KH = (0.9 f 0.3) x 10’. This value and the results in Table I show that KH is only weakly-dependent upon total ionic strength in the range, Z = 3 X 10-e3 X 10-‘m. In CaCI2-HC1-LiC1 mixed solutions, interference by both Li+ and H+ occurs and the simple theory leading to Equation 1 predicts the following equation for multiple interference (293):
*
E
=
E”
RT +In F
+ KLiaz f K ~ a o )
( ~ 1 ’ ’ ~
(3)
where the selectivity parameters have the same values as in binary mixed solutions. Substituting the previously determined values, K L = ~ 0.33 (2) and KH = 1.8 X lo’, Equation 3 fitted the results of Figures 3b and 3c to within 1 2 mV. 1776
At I = 0.6, Equation 3 was not obeyed with the above selectivity parameters but if the assumption is made that K L ~= 0.33 then a value KH = (1.7 i 0.3) X 102 was obtained. A similar increase of selectivity parameters with ionic strength has been observed previously for alkali-metals in binary mixtures (2,11). Equation 1 held over the whole of the higher concentration range, 10-3-10-1m CaCL, provided that the electrodes had only been immersed for a moderate period in aqueous solutions, e.g. < 12 hr. In the absence of the negative change, therefore, the interference by H+ seems governed by the same processes as the interference by alkali ions. The present results suggest that although the limit pH 3 5.5 for no significant interference by Hf is very satisfactory, by ensuring only short contact times before measurement, the pH limit may be extended to 4 at high [Ca*+]. The negative change is manifested only after prolonged contact with aqueous solutions. After such contact, the membrane is known to take up water. Direct evidence comes from IR spectra (2) and solubility measurements in the solvent DOPP (0.5 mole H20/mole DOPP); and indirectly from the decrease of membrane resistance with time of contact (2) and the appearance of a milky layer. Such milky layers are often observed during water transfer between two phases (12) and may be due to the expulsion of a third phase (13) or the formation of micelles surrounding the water molecules (14). In addition, metal and hydrogen ions almost certainly exchange in a hydrated form, thus providing another means of introducing water into the membrane (15). The presence of water is necessary for the electrode to function in a reproducible manner. The potentiometric response is erratic and the impedance of the membrane is irreproducible (16) until the membrane has been in contact with water or aqueous solution for several hours. The presence of water is likely to increase the solubility of CaC12or Ca(N0J2 in the membrane, possibly in the form of complexes of the type C a C l z ~ x H 2 0 ~ y D O or P P CaClz. (alkylphosphoric acid).yH20 (17). The solubility of Ca salts in organophosphorus solvents can be appreciable, e.g., CaC12 6H20 and Ca(NO&. 4H20 have solubilities of 1.90 and 0.99 molejliter, respectively, in tributylphosphate (18). If the negative change in potential is associated with the solubility of Ca salt, then its variation with [Ca2+]is consistent with the greater amount entering the membrane at higher [Ca2+], and the difference in response to CaCL and Ca(NO& is due to the different solubilities of these salts. When Ca-selective electrodes have been used to determine activities in solutions of high ionic strengths, reliable results were obtained only if the interior reference and exterior sample solutions were matched in [Ca2+] (11). The addition of a separate potential, arising from the solubility of Ca salts, to the Nernstian response could account for this observation. +
(11) R. Huston and J. N. Butler, ANAL.CHEM., 41,200 (1969). (12) J. T. Davies and E. K. Rideal, “Interfacial Phenomena,” Academic Press, New York and London, 1963, p 322. (13) A. S. Kertes, “Solvent Extraction Chemistry of Metals,” Proc. Int. Conference, Sponsored by U.K.A.E.A., 1965, p 377. (14) N. Pilpel, Chern. Reu., 63,223 (1963). (15) A. L. Myers, W. J. McDowell, and C. F. Coleman, J . Inorg. Nircl. Chem., 26, 2005 (1964). (16) M. J. Brand and G. A. Rechnitz, ANAL,CHEM., 41,1185 (1969). (17) Y.Marcus and A. S . Kertes, “Ion-Exchange and Solvent EXtraction of Metal Complexes,” Wiley-Interscience, London, 1969, p 538. (18) Ibid.,p 697.
ANALYTICAL CHEMISTRY, VOL. 44, NO. 11, SEPTEMBER 1972
The rate of initial decrease of emf is rapid and is dependent upon stirring in the aqueous phase. The first stage is, therefore, probably governed by the diffusion of H+ or Ca2+from the solution to the interface. The second slow process is likely to take place within the membrane and to consist of ionic migration restoring the system to equilibrium. We suggest that the slow process is due to a barrier layer formed a t the membrane-solution interface at low pH and high [Ca2+1 which hinders the passage of ions between solution and the interior. This layer might consist of a closepacked monomolecular layer of oriented chains or perhaps a cross-linked region of high viscosity. Monolayers of monooctyldecylphosphoric acid (MODP) and di-octyldecylphosphoric acid (DODP) spread on aqueous substrates do show marked changes in structure with substrate pH. On a substrate Z = O.lm'NaC1, a DODP film contracts from an expanded structure to a solid structure when the p H decreases from 5.7 to 3.0 and a MODP fdm changes from an expanded structure to a solid structure to a liquid or gel-like film (19). The presence of Mg2+ in the substrate leads to the formation of a condensed film of MODP, whereas without Mg2+ an expanded film is formed. The action of alkaline-earth ions, particularly Ca2+, forming condensed films with long-chain acids is well known (20). The absence of the slow process at low [Ca2+] or when interference by Li+ takes place is explained, in the first case, by the greater solubility of the film under these conditions (21), and, in the second case, by the greater solubility of the Li salt formed in the exchange. The bulk solubility of dialkylphosphoric acids is also reduced markedly in acid solutions, e.g., dibutylphosphoric acid has a solubility in O.lm H N 0 3 of the value in water (22), but the solubility is increased very markedly in contact with L i N 0 3solutions (23). (19) H. C. Parreira, J. Colloid Sci.. 20,742 (1965). (20) E. D. Goddard and J. A. Ackilli, ibid., 18,585 (1963). (21) J. Bagg, M. P. Haber, and H. P. Gregor, J. Colloid Interfac. Sci., 22,138 (1966). (22) A. S . Kertes, A. Beck, and Y . Habousha,J. Inorg. Nucl. Chem., 21,108 (1961). (23) V. B. Shevchenko and V. S . Smelov, At. Energy (USSR),5 , 542 (1958).
Table I. Selectivity Parameters, KE, for CaX2-HX Solutions This work PH a8/al1I2 KE 4.0 1.04 x 10-2 1.29 X 101 3.5 3.40 x 10-2 2.32 X 10' 3.0 1.13 x 10-l 1.71 X 10' 2.5 3.86 X lo-' 1.69 X 10' Average 1.83 X 101 Ross ( I ) , pH 3.0-5.0, [CaCln] = 10-1-10-4m Average (1.2 + 0.2) X IO2 Ross (9,pH 3.0, [CaC12] = 10-1-10-4m Average (0.8 + 0.3) X lo2 Orme (6), pH 3.0, [CaCh] = 10-3m Average (1.4 + 0.2) X 10'
Slow extractions have been reported of Be, Fe, and A1 by dialkylphosphoric acids from acid solutions. For example, in the extraction of BeZ+ by didodecylphosphoric acid from acid solutions, equilibration times of 300-500 min were found (24, 25). The equilibrium time increased with decreasing pH and increasing [Be2+]. Large differences in extraction rate were found with pure dialkylphosphoric acid compared to an impure acid with monoalkyl impurity. Equilibration times were reduced from 300 min, to 15-20 min for a mixture of acids containing 20% monoalkyl acid. The monoalkyl acid has a higher solubility than the dialkyl acid and this difference probably accounts for the increased rate. The presence of monoalkyl acids in the liquid-membrane might similarly be expected to change the electrode response and could account for differences between batches of ion-exchanger.
RECEIVED for review January 6, 1972. Accepted May 2, 1972. (24) C . J. Hardy, B. F. Greenfield, and D. Scargill, J . Chem. SOC., 1961,174. (25) R. A. Wells, D. A. Everset, and A. A. North, Nucl. Sci. Eng., 17,259 (1963).
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