values correspond to a residence time of 9.5 sec. The expected effect of the dark reaction is calculated for atmospheric data (8) using a 10-second residence time. The measured concentrations and calculated atmospheric concentrations are shown in Figure 4. In obtaining this figure, the ozone concentration was taken to be that for oxidant. It should be noted that since (NO,) [= ( N O z ) (NO)] is conserved in the dark reaction, no error is expected in this quantity due to these considerations. O n the other hand, measured values of ( N O ) and (0,) will be systematically low, (NO2) will be high. The quantity ( N O ) ( 0 3 ) / ( N 0 2 )which , has been used as a measure of the photochemical equilibrium and the amount of sunlight, will be systematically low. The effect considered here is based entirely on the gas phase dark reaction. Reactions with the wall of the inlet system have been neglected as have reactions which may occur within the analytical instrument being used. Additional experimental data are required t o determine the magnitudes of these errors. It may be seen from Figure 4
+
(8) P. A. Leighton, “The Photochemistry of Air Pollution,” Academic Press, New York, N. Y . , 1961, p 273.
that large relative errors for a given species will occur only when the concentrations are low-near the detection limit for many of the wet chemical instruments now in use. However, with the advent of more specific and sensitive methods for measuring nitrogen oxides and ozone (9-11), consideration should be given to reducing the errors due t o the residence time. Decisions concerning the design of sampling lines must consider losses due to residence time, in addition to those resulting from the impaction of aerosols and chemical losses on the surfaces.
RECEIVED for review May 20, 1971. Accepted July 21, 1971. Partial support for this work was obtained from these grants from the Air Quality Office of the Environmental Protection Agency : 2-TO1 AP00029-07, Air Pollution Training; and A P 00336, Influence of Aerosol Characteristics of Visibility. (9) J. A. Hodgeson, K. J. Krost, A. E. O’Keeffe, and R. K. Stevens, ANAL.CHEM., 42, 1795 (1970). (10) A. Fontijn, A. J. Sabadell, and R. J. Ronco, ibid.,p 575. (11) G. W. Nederbragt, A. van der Horst, and J. van Duijn, Nature, 206, 87 (1965).
~~
Liquid Anion Membranes Dependence of Selectivity Factor on Organic Salt Concentration Pier Roberto Danesi, Giancarlo Scibona, and Bernard0 Scuppa Industrial Chemistry Laboratory, Commitato Nazionale per L’Energia Nucleare, Centro di Studi Nucleari della Casaccia, Rome, Italy
PREVIOUS WORK ( I , 2 ) on the behavior of liquid anion membrane electrodes has shown that the potential of these electrodes can be described by means of the equation (at 25 “C) P
P
Ji
J2
with ai, Ki, and ui, respectively, activity, distribution constant, and mobility in the organic phase of the ith ion. The two integrals
and
are given in Reference (2) and
can be neglected when the mobility of the organic cation is much smaller than that of the anions. By considering a twoanion system with anion mobility larger than that of the organic cation, Equation 1 becomes V = 59 log [(al‘
+ Pz.la2’)/(al’’+ PZ.IUZ”)I
(2)
The selectivity factor of the liquid anion electrode, Pz.1 is given by P2.i = K Z U Z / ~ U ~ (3) By using the equilibrium constant of the biphasic ion exchange 2 e 25 1, K2.1,and the ion pair formation reaction, . (where square brackets represent constant Kij = activities or concentrations), Equation 3 becomes
+
+ [u]/[i] [a
(1) P. R. Danesi, F. Salvemini, G. Scibona, and B. Scuppa, J . Phys. Chew., 75, 554 (1971). (2) J. Sandblom, G. Eisenman, and J. L. Walker, ibid., 71, 3862
(1967). 1892
P2.i
=
( u z / u ~ .> ( K ~ J / K .~ Kz.1 J)
= UZ/UI
. K*z,i
(4)
Although K I J , K ~ Jand , K2.1 are thermodynamic quantities, a concentration variation of the liquid ion exchanger can affect the macroscopic dielectric properties of the organic phase and, consequently, their values. Of course, the (u&1) ratio can also be concentration dependent. A dependence of P2.*on the liquid exchanger concentration is therefore with the expected. In the present work the variation of P2.1 concentration of the liquid exchanger for membranes consisting of benzene solutions of tetraheptylammonium nitrate interposed between aqueous solutions containing the nitrate and chloride anions has been experimentally studied. The implications of the P2.1variations on the analytical determination of the ion concentrations by means of liquid ion exchanger membrane electrodes are also evaluated. EXPERIMENTAL
Reagents. HC1, H N 0 3 , LiC1, and LiN03 of analytical grade purity, supplied by Carlo Erba, have been used in the experiments. Benzene of the same type of purity, Carlo Erba, has been used. Tetraheptylammonium iodide (THAI), Eastman Kodak, has been used t o prepare the other alkylammonium salts. The preparation of THACl and THAN03 has been already reported (3). In order to take into account (3) P. R. Danesi, M. Magini, and G . Scibona, “Solvent Extraction Research,” A. S. Kertes and Y. Marcus, Ed., John Wilev & Sons, New York, N.Y., 1969,p 185.
ANALYTICAL CHEMISTRY, VOL. 43, NO. 13, NOVEMBER 1 9 7 1
80 70
60
Figure 1. Membrane potential, V(mV), us. logarithm of nitrate ion concentration ratio ([N03-]”/[NOa-]’) in the two half cells, Equation 5
56 40
Experimental points refer to the following THAN03 concentrations: X 0.05M, 0 0.1M, 0.2M, 0 0.3M. Solid lines calculated by Equation 2 with the Pz.1 values of Table I
30
20
10
0 10-
possible concentration variations due to solubilization of the alkylammonium salt or exchange reactions (in the case of biionic potentials), all aqueous and organic solutions have been analyzed by standard analytical procedures before and after the membrane potential measurements. Silver-silver chloride electrodes have been prepared according to Reference (4). Membrane Potential Measurements. The biionic membrane potential measurements have been performed with stirred (organic and aqueous phase) cells of the type
Log
([NO;]’j/
10-
5
[NO:)’)
Table I. Variation of Selectivity Factor and Mobility Ratio with the Membrane Composition THAN03 concn M E (Exp) P2.1 (exp) K*z.I UNOa/UC1 0.3 0.2 0.1 0.05
3.68 f 0 . 0 2 2.98 i 0 . 0 2 2.58 f 0 . 0 2 2.44 i 0 . 0 2
20k 1 24 f 2 33 i 2 40+ 4
19 54 125 189
1.1 f 0 . 5 0.44 =t0 . 0 3 0.26 rt 0 . 0 2 0.21 i0.02
+
Ag, AgCljLiCIMcl, LiN03, M I C O ~ ( M CMh-03 I = 0.1 ; M c 1 1 M ~constant)ITHANO.?M 03 in
+
benzenelLiC1M~1,LiN03Mh-03,( M C I M K O=~ 0.1; M c l l M ~variable)IAgCl, ~3 Ag ( 5 ) The emf of these cells is given by E = V - RT In (a’cI/a’’cl) where V is given by Equation 2. ~ C represents I the chloride ion activity which in our case has been replaced by its concentration since a constant ionic strength was used. The following THANO.] concentrations have been used in the experiments 0.05,0.1,0.2, and 0.3M. Four series of measurements have been so obtained (one for each THANOa concentration). In each series of measurements the organic and aqueous solutions used in the cells had been previously equilibrated, the organic phase with the aqueous solution containing the lowest LiCl concentrations and the aqueous phases with the benzene solution of THANOP. The cell used was the same as that described in Reference (5). The emf values have been measured by using a Cary 31 vibrating reed electrode electrometer at 25 + 1 “C. As a consequence of the high impedance of the system, the electrochemical cells were placed inside a copper Faraday cage to perform noiseless electrical potential measurements. Dielectric Constants. The dielectric constant measurements have been performed by following a pulse echo technique. The instrument used was a Hewlett Packard time domain reflectometer Model 1415A. The rise time of the (4) J. Janz, “Reference Electrodes,” Academic Press, New York, N.Y., 1961, Chap. 4. (5) G. Scibona, L. Mantella, and P. R. Danesi, ANAL.CHEM., 42, 844 (1970).
incident pulse was about 150 psec, which is equivalent to a band width of 2.4 GHz. The dielectric constant value ( E ) was determined through the measurement of the amplitude of the reflection step by the equation, E = [(l - @/(I S)12 with 6 reflection coefficient.
+
RESULTS AND DISCUSSION Figure 1 reports the experimental results as V (membrane potential) us. logarithm of nitrate ion concentration ratio in the two half cells, Equation 5. The P2.1values reported in Table I have been calculated through Equation 2 by using the method reported in Reference ( I ) . In Equation 2, a1 and a2 stand for the C1- and NO3- activities, respectively. P2.1 represents the selectivity factor of the tetraheptylammonium nitrate in the benzene membrane to the nitrate ion in the presence of chloride ions. Activities have been replaced by concentration since a constant ionic strength has been used in the experiments. Some information on the factors ~ the alkylammothat can promote the dependence of P Z .on nium salt concentrations can be obtained by considering the influence of the macroscopic dielectric constant of the oil phase on Kg.1, K ~ JK, g J ,and u2/u1(see Equation 4). The exchange constant K2.1and the ion pair formation constants, K l and ~ K Z J ,are thermodynamic quantities and, therefore, independent of the salt concentration. However, this independence is true when the physical properties of the phase are not altered by the increased presence of the salt molecules. In the case of benzene solutions of THANOa, how-
ANALYTICAL CHEMISTRY, VOL. 43, NO. 13, NOVEMBER 1971
1893
AC
A V(mV)
lxld2- 2xlG2
13
(a)THANO,
1x18- 2x1d2
15
(b)THAN03 a05 M
1x164- &tb4
0.5
(a)THANO,
1
(b)THAN03 4 0 5
d4
1 x 164-
0.3
0.3
M
M
M
\
+a
t.
'
'
"
'
I " " [
8
1
, " , I
-2
-3
-4
I
'
" ' I
-1
Figure 2. Membrane potential of cells (a) and (6)(corrected for fixed concentration term VO)as function of logarithm of nitrate concentration Vo =
-
15.4 mV for THAN03 0.3M (a) and VO = - 32.8 for THAN03 0.03M (b)
ever, the dielectric constant, E , of the oil phase increases with the salt concentration as shown in Table I (column 2 ) and a dependence of K2.1,K ~ Jand , KZJ on the salt concentration is expected, Such a dependence has to be clearly understood in terms of the change of physical properties of the oil phase as a consequence of the increased presence of salt molecules. The dependence of K 2 . 1 , Kij, and KZJ on E can be in first approximation estimated by using the equations A G N O ~ C=I -RT In K N O ~=C ~
-(L 1 - L)}(6) E,
rc1-
moa-
and AGij
=
-RT In Kif = -(Ne2/d,eKT)
(i
=
1, 2) (7)
The models and approximations implied in Equations 6 and 7 are described in references (6) and (7). The numerical calculations have been performed by using the following numerical values: E , (dielectric constant of water) = 78.54; NO^- (ionic radius of N03-oion) = 2.07 A ( I ) ; rcl- (ionic radius of C1- ion) = 1.84 A ( I ) ; r T H A (ionic radius of the tetraheptylammonium ion) = 7.16 A ( I ) ; dl = r c l ITHA; dz = r N o a f r T H A ; Ne2/?= 163.8 A Kcal/mole; E (dielectric constant of the oil phase) = see column 2 of Table I.
+
(6) G. Scibona, J. F. Byrum, K. Kimura, and J. W. Irvine, Jr., "Solvent Extraction Chemistry," North Holland, Amsterdam, 1967, p 398. (7) J. T. Denison and J. B. Ramsey, J. Amer. Chem. SOC.,77, 2615 (1955).
1894
Numerical values of KNOscl= (K~J/K~J)Kxo,c~ are reported in Table I, column 4. A decrease of K * N O ~with C I E is observed. The decrease of K * K with ~ ~E , ~and~ therefore with the increase of the T H A N 0 3 concentration, agrees with the experimental trend of PNOaCI. The main uncertainties in Equations 6 and 7, besides the use of the macroscopic dielectric constant, are due to the values of the ionic radii, ri, (i = NO3, Cl) and the closest approach distances, d,. The A G N O ~ Cvalue I is strongly affected by changes in rE values, while it is less sensitive to the dimension of the cation radius r T H A . From Equation 4 the mobility ratio, u ~ / u ~ can be calculated. Numerical values of u2/u1are reported in Table I, column 5 . However, as a consequence of the uncertainties involved in Equations 6 and 7, these values have to be cautiously considered. Experimental measurements of the ratio u2/ul are needed to test more quantitatively the approximations of the models and the validity of the equations used. There is then clear experimental evidence that the P 2 . 1 values depend on the concentration of the alkylammonium salt whose benzene solution forms the liquid membrane. Some guidelines on the factors affecting the P2.1values can be obtaining by using a crude model (Equation 6) and the current theories of the liquid membrane potential. The analytical implications of the dependence of the P z . 1 values on the alkylammonium salt concentration can be discussed at this point. To this purpose let us examine the two curves of Figure 2 . These two curves have been calculated for the systems : Ag, AgCI/[Cl-] = O.OlM, [NO,-] = 0.09MITHAN03 0.3M in benzenel[N03-] [Cl-] = 0.1M, [Nos-] = variableIAgC1, Ag with P2.1= 20 and V(mV> = - 15.4 59 log ([Cl-] 4 0 W 3-11
ANALYTICAL CHEMISTRY, VOL. 43, NO. 13, NOVEMBER 1971
+
+
+
Ag, AgCl/[Cl-] = O.OlM, [NO,-] = 0.09MITHAN03 0.05Min benzene [NO3-] [Cl-] = O.lM, [NO3-] = variablejAgC1, Ag with P2.l = 40 and V(mV) = -32.8 59 log ([Cl-] 40 WO3-1)
+
+
+
The data of Figure 2 show that in the case of the determination [NOB-] = O.lM, the of nitrate ions in a medium [Cl-I concentration of the membrane to be used is dictated by the order of magnitude of the nitrate ion concentration in the solution. In fact for a nitrate ion concentration of the order of 10-2M, system (a) gives a AV of 13 mV for AC = 1.10-22.10-2M, while system (b) gives a AV of 15 mV. In this case, since both values 13 mV and 15 mV can be read with the same precision, it can be worthwhile to use system (a). This system, which makes use of a T H A N 0 3 0.3M benzene solution represents a low impedance electrode with respect to system (6). When the nitrate concentration is of the order of 1 X 10-4M, the two systems give AV = 0.5 mV
+
(system a) and AV = 1 mV (system b) for a AC of 1.10-42.10-4M. It is clear in this case that system (b) is more convenient than system (a) despite its greater impedance due to the use of a THAN03 0.05M benzene solution. Therefore the choice of the liquid electrode concentration depends on the required value of the selectivity factor P2.1.The examples discussed show that a reasonable compromise between impedance and precision requested can be worked out when the characteristics of the electrode are well known. ACKNOWLEDGMENT
The authors thank Mr. F. Salvemini for his contribution in the membrane potential determinations and Mr. M. De Carolis for performing the dielectric constant measurements. RECEIVED for review April 9, 1971. Accepted July 23, 1971.
Polarographic Determination of Chloride, Cyanide, Fluoride, Sulfate, and Sulfite with Metal Chloranilates Ray E. Humphrey and Clyde E. Laird Department of Chemistry, Sam Houston State University, Huntsville, Texas 77340 RELATIVELY INSOLUBLE METAL chloranilates have been used for the spectrophotometric determination of a number of anions. The less soluble or undissociated compound of the metal with the respective anion is formed releasing the absorbing chloranilate species into solution. For example, the determination of chloride involves the use of mercuric chloranilate as shown in Equation 1. HgCh
+ 2C1-
+ HgC12
+ Ch2-
(1)
The chloranilate ion released, represented by Ch2-, absorbs at 525 nm and 330 nm. Chloride, (1, 2) cyanide (3, 4) and sulfite (4) have been determined spectrophotometrically using mercuric chloranilate, fluoride has been measured using thorium chloranilate (5), and for sulfate ion barium chloranilate has been employed (6). Chloranilic acid and chloranilate ion undergo a reversible two-electron reduction at the dropping mercury electrode (7). In this work, the current from the reduction of the released chloranilate ion was measured and related to the concentration of the anion sought. No report was found in the literature for the determination of chloride by means of a polarographic reduction current. Cyanide has been determined polarographically by measuring the reduction current for mercuric cyanide (8). Sulfite is converted, by (1) J. E. Barney I1 and R. J. Bertolacini, ANAL.CHEM., 29, 1187 ( 1957). ( 2 ) R. J. Bertolacini and J. E. Barney 11, ibid., 30, 202 (1958). (3) E. Hoffmann, Z . Ami. Chem., 185, 372 (1962). 43, 1100 (1971). (4) R. E. Humphrey and W. Hinze, ANAL.CHEM., (5) A. L. Hensley and J. E. Barney 11, ibid., 32, 828 (1960). (6) R. J. Bertolacini and J. E. Barney 11, ibid., 29, 281 (1957). (7) J. Weissbart and D. Van Rysselberghe, J . Phys. Chem., 61, 765 (1957). (8) J. S . Hetman, Lab. Practice, 10, 155 (1961).
reaction with hydrogen ion, to sulfur dioxide which is polarographically reducible (9, 10). Fluoride has been determined by the decrease in the reduction current for Fe3+(11) or the A13+-Solochrome Violet RS complex (12) on adding this ion to a solution containing one of these species. Apparently, very little has been done in the way of a metathesis reaction to exchange chloride, cyanide, fluoride, or sulfite for a polarographically reducible ion. Sulfate has been determined polarographically by reaction with barium chromate to yield the reducible chromate ion (13). EXPERIMENTAL
Polarography. Polarograms were recorded with a SargentWelch Model XVI Polarograph in conjunction with a Sargent Constant Head Dropping Mercury Electrode assembly. A conventional H-cell, with a saturated calomel electrode (SCE) on one side, was employed. The drop time of the capillary electrode was in the range of 4-5 seconds. Solutions were purged with nitrogen prior to measurements. Polarograms were recorded for all of the reaction solutions. The maximum current value was used in establishing calibration plots and obtaining recovery data. Maximum pen travel was determined at the appropriate potential. For the lowest concentration, a current sensitivity setting of 0.01 pA/mm was used. Reagents and Solvents. Barium chloranilate and methyl cellosolve were Matheson, Coleman and Bell products. (9) D. B. Aulenbach and J. C. Balmat, ANAL. CHEM.,27, 562 (1955). (10) L. L. Ciaccio and T. Cotsis, ibid., 39, 260 (1967). ( 1 1 ) C. E. Shoemaker, ibid., 27, 552 (1955). (12) B. J. MacNulty, G. F. Reynolds, and E. A. Terry, Airalyst, 79, 190 (1954). (13) J. Mayer, E. Hluchan, and E. Abel, ANAL.CHEM., 39, 1460 (1967).
ANALYTICAL CHEMISTRY, VOL. 43, NO. 13, NOVEMBER 1971
1895
