Table 111. Related Phthalimides in Literature Cpd No. 19 20 21 22 23 24 25 26 27 28 29
Compound : Y = C&&NOz YCH?CBr=CH, YCH2CHClCHzCl YCHpCH=CHCHzY YCHpOCHpY YCHzCHpSCHzCHzY YCHzCHYCHpY YCHpCH(CH3)CHpY YCH2-CHBrCH2Y YCHpCHClCHzY YCHzCHOHCHpY YCHzCOCHpY
Mol wt 266 293 226-227 3 36 380 479 345 413 369 350 348
MP, "C 125-6 93 219 209.5-21 1 125 226-227 169-1 7 1 196-198 215-216 205 268
11). The identity of the remainder, with derivatives already recorded in the literature (Table III), was shown by comparison of melting points, mixture melting points (when the substances prepared by both methods were available), and by the use of instrumental methods. Physical Properties. The IR spectra were obtained using a Beckman IR20A and a Perkin-Elmer Model 137 instrument; the samples were prepared as pressed plates in potassium bromide or in Nujol mulls. Characteristic carbonyl stretching absorptions for the 5-membered-ring imide were found near 5.65 microns (weak) and 5.85 microns (strong). Mass spectra were obtained using a Hitachi-Perkin-Elmer RMS-4. All samples were introduced through the direct probe and volatilized in the temperature range of 80-150 "C. The analytical sample of 2-chloroallylphthalimide 2 showed m/e 221 (M, parent) 5.7%; 186 (M - C1) 35%; 160 (M - CC1=CH2) 22.9 %. A once-recrystallized specimen showed m / e 258 (molecular weight of dichloropropylphthali(20) A. E. Cashmore and H. McCombie, J . Chem. SOC.,123, 2884 (1923). (21) S. Gabriel and W.Michels, Ber., 25, 3057 (1892). (22) F. G. Mann, J. Chem. Soc., 1927,2907 (esp. 2912). (23) A. F. Fairbourne and G. W.Cowdrey, ibid., 1929, 129. (24) S. Gabriel and T. Posner, Ber., 27, 1042 (1894).
mide) 2-3 %, and less than 1 of phthalimide, in addition to the foregoing. (Phthalimide, itself, has mje 147, 104, 103, and 76). Pentaerythritol tetrabromide showed m/e 384 (M, parent ; Br = 79) 39%; 305 (M - Br) 100%; 225 (M - Br, HBr) 9.6%; 211 (M - Br,CH,Br) 18%,andmanyminorpeaks. The N M R spectra were measured on a Varian A 60 spectrometer. The chemical shifts are reported in ppm downfield from tetramethylsilane. The solvents were CDC1, and DMSO-de. 2-Chloroallylphthalimide 2; 4.53 (m, 2H, -NCH?-), 5.44 (m, 2H, =CH2), and 7.9 (m, 8H, arom.). 1,4-Diphthalimidobutane 6; 1.7 (t, 4H, insulated methylene), 3.6 (t, 2H, -NCH2CH20-), and 7.8 (m, 8H, arom.). ACKNOWLEDGMENT
We are pleased to have had the invaluable assistance of Miss T. J. Davis for some of the IR, D. Maier for the MS, T. H. Regan for some of the NMR, and D. Ketchum for the microanalyses-all of the Eastman Kodak Company Research Laboratory. RECEIVED for review January 3, 1972. Accepted April 17, 1972.
Selectivity Studies on Anion-Selective Membrane Electrodes Stig Back Department of Physiology and Biophysics, Medical Centre, C niaersity of Uppsala, Sweden
SANDBLOM,EISENMAN, AND WALKER (1) have deduced an expression for the potential of a liquid ion exchange membrane for the special case of two counter ions and strong association between ion and site in the membrane phase. This seems to correspond most nearly to the commercially available liquid membrane ion-selective electrodes. The following expression can be written for this potential: E
=
constant - (1 -
(
al
7) -In
RT F
F
+ (PQ( p+l + ps)'kt'a? ps).kl
(
In al
r =
)
p2s'kl. Ki.a? +pis*ki. Kz
(1)
(1) J. Sandblom, G. Eisenman, and J. L. Walker, Jr., J . Phys. Chem., 71, 3862 (1967). 1696
where p1, p 2 , p s , pis, and p Z s are mobilities in the membrane phase of ion 1, ion 2, the dissociated resin ion, the undissociated ion-pair of the resin ion, and counterion 2, respectively. In the external solution and membrane phase, KI and K2 are dissociation constants of ion-pairs of the exchanger ion with counterion 1 and counterion 2, respectively, and
(111
+
p s ( K ~ . p z s- Kz*pis) ps).p2s.Ki
-
(pz
+
ps).pis.K2
(2)
The selectivity ratio in the second logarithmic term of Equation 1 has been equated to the ion-exchange equilibrium constant since it is reasonable to assume that p s = ,uZa(1). However, in the derivation of Equations 2 and 3, the dissociation constants K1 and K? have been defined Only in terms Of concentrations of the various species in the membrane phase and
ANALYTICAL CHEMISTRY, VOL. 44, NO. 9, AUGUST 1972
not in terms of their activities.
is only the ion-exchange equilibrium constant.
- 0.1
-1ogexp 1.0-‘E E
Consequently, the quantity
RYF
I t is
ki.Kz interesting to note that the ion exchange coefficients of certain ions in the case of solid ion-exchangers show a decrease at low concentrations of the ion of interest in the external solution (2). Thus, the decrease in selectivity ratios a t low concentrations, observed in this study, might be related t o a similar decrease in the ion-exchange itself. While the specific origin of this special effect can only be summarized a t this point, it is nevertheless clear that the selectivities of the liquid membrane electrodes studied are largely consistent with current theories of “mobile” type membrane electrodes (3, 4). Side-reactions of the ion-pair itself, such as dissociation or association, should give a contribution to the potential. EXPERIMENTAL
The liquid membrane electrode utilized in this study was the ion-selective electrode (Model 92-20), obtained from Orion Research, Inc. E M F measurements were made usinga Corning No. 12 pH-meter with saturated calomel electrode as a reference electrode. Stock solutions of sodium salts of the anions were made. All chemicals were of analytical grade and potential measurements were made in the concentration range of l-10-5M. The body of the electrode was filled with 0. 1Mtetrabutylammonium perchlorate or 0.1M tetrapropylammonium perchlorate in methylene chloride. The reference solution consisted of 0.01M sodium chloride and 0.1M sodium perchlorate. Before measurement, the reference solution and the membrane solution were equilibrated with each other for five minutes. The membrane electrode was immersed to a constant depth in the solution. RESULTS AND DISCUSSION
30 -
AIo Pi-
20-
-
10
-
1
2
3
4
5 -loga2
Figure 1. Plots of Equation 9 for the interfering anions iodide and picrate with tetra-propylammonium perchlorate as ion exchanger -log exp E‘EV0.I
I
RVF
- 0.1
40-
I
Methods of Evaluating the Selectivity Ratios. Equation 1 differs from the empirical Equation 3 :
E
=
RT F
Eo - - In (al -t K . a 2 )
(3)
where a1 = activity of the primary univalent ion, uz = activity of any univalent ion 2, to which the electrode responds, and K = selectivity ratio for ion 2 for the given electrode. The potential of an ion-selective electrode in a solution containing only the primary univalent ion is given by:
E1
=
RT F
Eo - 2.303 - logal
(4)
If the solution does not contain the primary ion but any other univalent ion with a selectivity ratio K l , the potential of the electrode in such a solution can be expressed by Equation 5, obtained by substituting al = 0 in Equation 3.
RT log KI - 2.303 RT log uz E2 = Eo - 2.303 -
F
If a1
=
F
(5)
a2
log KIP=
E1 - Ez
20 -
10 -
I‘ I
1
2
3
4
5
’
-toga2
Figure 2. Plots of Equation 9 for the interfering anions iodide and picrate with tetra-butylammonium perchlorate as ion exchanger
~
(2) B. Chu, D. C. Whitney, and R. M. Diamond., J. Inorg. Nucl. Chem., 24, 1405 (1962). (3) . . H. Levin. W. J. Diamond, and B. J. Brown., Zird. Enn. - Chem., 51, 313 (1959). (4) K. Srinivasan and G. A. Rechnitz.. ANAL. CHEM.. 41. 1203 (1969). I
AIOR‘
2.303 RT
F
.
30-
I t has been usual to calculate KIZfrom values of E1 and Ez measured at a concentration of 0.1M (5, 6). However, K12 can be obtained from measurements a t any other concentra(5) T. S. Light and James L. Swartz., Anal. Lett., 1, 1113, 825 (1968). (6) G. A. Rechnitz, M. R. Kretsz, and S. B. Zamochnick., ANAL. CHEM., 38, 973 (1966).
ANALYTICAL CHEMISTRY, VOL. 44,
NO. 9, AUGUST 1972
1697
~~~
~
In Equation 9, the potential of the electrode in a solution containing both the ions is utilized in the calculation. Combining Equations 3 and 4, one obtains:
Table I. Selectivity Ratios for Iodide (I-) and Picrate (Pi -) on the Perchlorate-Selective Electrode with Tetra-Propylammonium Perchlorate (TPrA + Clod-) or Tetra-Butylammonium Perchlorate (TBA + C104-) as Liquid Ion Exchanger Activities (a*-) Potentials, +mV
(exp E1 - Ed RT
TBA+
TPrA+*
c10-,
0.0755 0.00899 0.000964 0.000099 0.0000099
which can be rearranged t o give a n expression for K I Z
c10-,
I-
Pi-
I-
Pi-
110.0 137.0 157.0 170.0 176.2
93.0 108.7 139.0 173.5 184.0
107.0 145.0 171.2 197.0 215.7
115.0 124.6 170.8 196.2 210.0
tion, as long as the two ions are a t the same activity in their solutions. If Equation 2 and 3 are combined with E1 = E z , we get: (7) where a1 and a2 are the activities of the two ions which produce the same potential when present separately. Equations 6 and 7 involve measurements in solutions containing only one of the ions in any test solution.
UI
- a1
Plots of Equation 9 are given in Figures 1 and 2. SELECTIVITIES AND INTERFERENCES
From the studies in Table I, it is observed that the quaternary ammonium compound shows the highest selectivity when the interfering ion is highly extractable, i.e. the extraction constant ( E Q c L o -is ~ )of lower order than that of picrate. The extraction constants are found to be in the sequence Pi- > C10-4 > I- for symmetric quaternary ammonium compounds (7,8). RECEIVED for review August 9, 1971. Accepted January 24, 1972. (7) K. Gustavii and G. Schill., Acta. Pharm. Suecica, 3, 249 (1966). (8) G. Schill., So. Kern. Tidskr., 80, 10 (1968).
Polarographic Reduction of Tin(lV) in the Presence of a-Mercaptopropionic Acid Atsuyoshi Saito and Sadayuki Himeno Department of Chemistry, College of General Education, Kobe University, Nada-ku, Kobe, Japan
THE POLAROGRAPHIC REDUCTION of tin(1V) is generally irreversible. To show two distinct waves, it is ordinarily carried out in such complexing agents as chloride ( I ) , bromide (2, 3), pyrogallol ( 4 ) , chloranilic acid (3, and P-mercaptopropionic acid (6). In the presence of a-mercaptopropionic acid (a-MPA), tin(1V) shows a reversible two-step reduction wave at p H less than 2.5 and an apparent single step reduction wave above p H 2.5. The reduction wave is well formed and appears suitable for analytical purposes. EXPERIMENTAL Apparatus. The polarograms were recorded on a Yanagimot0 Model PA-102 polarograph, with a digital voltmeter ( i l mV) (Model 156-A, Kikusui Electronics Co., Tokyo, Japan) employed for precise potential measurement. p H D. E. Sellers and D. J. Roth, ANAL. CHEM., 38, 516 (1966). T. Kitagawa and K. Nakano, Rec. Polurogr., 9, 121 (1961). T. Kitagawa, Jup. Analyst, 6 , 603 (1961). A. J. Bard, ANAL. CHEM., 34, 266 (1962). 5 ) C. C. Boyd and J. G . Wardeska, ibid., 42, 529 (1970). 6 ) S. L. Phillips and R. A. Toomey, ibid., 37, 607 (1965).
1) 2) 3) 4)
1698
values were measured with a Hitachi-Horiba type M-4 p H meter. The capillary had a value of m 2 ’ 3 f 1 of / 6 1.990 a t -0.35 V os. SCE and of 1.844 a t -0.80 V cs. SCE in a solution of 0.2mM tin(IV), 0.10M HCIOd, 0.10M NaC104, lOmM a-MPA, and 0.0015 Triton X-100. An external saturated calomel electrode (SCE) was used as the reference electrode. The cell was immersed in a water bath maintained a t 25 i 0.1 “C. Reagents. Standard tin(1V) solutions were prepared by dissolving granulated 99.999 tin (Mitsuwa Chemicals Co. Osaka, Japan). a-MPA was reagent grade (Tokyo Kasei Kogyo Co., Tokyo, Japan) and used without further purification. All other inorganic chemicals were reagent grade. Procedure. To a known volume of supporting electrolyte, sufficient a-MPA was added so that the final solution was 10mM, unless indicated otherwise. An aliquot of 0.01M tin(1V) stock solution was then added. This sequence of addition was followed t o avoid hydrolysis of tin(1V). After adjusting the p H to the desired value with 1M perchloric acid or 1 M sodium hydroxide, the solution was deaerated with nitrogen gas for a t least 10 minutes. The polarograms were obtained in the usual manner.
ANALYTICAL CHEMISTRY, VOL. 44, NO. 9, AUGUST 19172
