2401
Anal. Chem. 1984, 58, 2401-2405
Robinson, W. 0.; Dudley, H. C.; Williams K. T.; Byers, H. Q , Ind. Eng. Chem. Anal. Ed. 1994, 6 , 274. Mulford, C. E. At. Abs. Newsl. 1988, 5 , 135. Beal, A.R. J . Fish Res. Board Can. 1975, 3 2 , 249. Holak, W. J . Assoc. Off. Anal. Chem. 1978, 5 9 , 650. Olson, 0. E.; Palmer, 1. S.; Whitehead, E. I. Methods Biochem. Anal. 1073, 2 1 , 39. Analytical Methods Commlttee Analyst (London) 1979, 104, 778. Levesque, M.; Vendette, E. D. Can. J . SoilScl. 1971, 5 1 , 85. Nygaard, D. D.; Lowry, J. H. Anal. Chem. 1982, 5 4 , 803. Masson, M. R. Mikrochlm. Acta 1978, I , 419. Cappon, C. J.; Smith, J. C. J . Anal. Toxicol. 1078, 2 , 114. Neve, J.; Hanocq, M.; Molle, L.; Lefebvre, G. Analyst (London) 1982, 107, 934. Valenta, P.; Rutzel, H.; Nurnberg, H. W.; Stoepler, M. 2.Anal. Chem. 1977, 285, 25. Nurnberg, H. W. Electrochim. Acta 1077, 22, 935. Princeton Applied Research Workshop Manual, Princeton, NJ, 1973.
Adeloju, S. B.; Bond, A. M.; Brlggs, M. H.; Hughes, H. C. Anal. Chem. 1083, 5 5 , 2076, and references cited thereln. Gorsuch, T. T. Analyst (London) 1959, 84, 135. Watklnson, J. H, Anal. Chem. 1008, 3 8 , 92. Olson, 0. E.; Palmer, I. S.: Cary, E. E. J . Assoc. Off. Anal. Chem. 1075, 5 8 , 117. Clinton, 0. E. Analyst (London) 1977, 102, 187. Bajo, S . Anal. Chem. 1978. 5 0 , 649. Pahlavanpour, B.; Pullen, J. H.; Thompson, M. Analyst (London) 1980, 105, 274. Subramanlan, K. S.; Meranger, J. C. Analyst (London) 1982, 107, 157. Andrews, R. W.; Johnson, D. C. Anal. Chem. 1975, 4 7 , 294. Andrews, R. W.; Johnson, D. C. Anal. Chem. 1978, 4 8 , 1058. Posey, R. S.; Andrews, R. W. Anal. Chlm. Acta 1981, 124, 107. Christian, G. D.; Knoblock, E. C.; Purdy, W. C. J . Assoc. Agric. Chem. 1985, 4 8 , 877. Lloyd, B.; Holt, P.; Delves, H. T. Analyst (London) 1082, 107, 927. Janghorbani, M.; Tlng, B. T. G.; Young, V. R. Am. J . Clln. Nu&. 1981, 3 4 , 2818. Ihnat. M. Anal. Chlm. Acta 1078, 8 2 , 293. Ihnat, M.; Miller, H. J. J . Assoc. Off. Anal. Chem. 1977, 60, 1414. Ihnat, M.; Thompson, B. K. J . Assoc. Off. Anal. Chem. 1980, 6 3 , 814.
for review February 24, lgg4* Accepted May
151
1984.
Applications of Cyclic Voltammetry in the Characterization of Complexes at Low Ligand Concentrations H. M. Killa,' Edward E. Mercer, and Robert H. Philp, Jr.* Department of Chemistry, University of South Carolina, Columbia, South Carolina 29208
A simple and general simulation for calculation of cycllc voltammetrlc curves for reversibly reduced complexes In the presence of low ligand Concentration Is described. Calculated curves are presented whlch demonstrate effects of ligand concentration and formation constants on the forms of curves for single complex and binary systems. Analysis of curve shape allows calculatlon of PI values for systems not accessible by conventional polarographic analysis. log P, values of 5.5 for the Cu(I1)-oxalate system and 5.6 for the Cd( I I)-propylenedlamlne system were obtained by this method. Both are In good agreement with values reported from POtentlometric analysis.
Electroanalytical techniques have long found wide application in the characterization of complex ions in solution. Early treatments were limited by the assumptions of a reversible electrode reaction, a single complex in solution, and the presence of a large excess of ligand ( 1 , 2). DeFord and Hume ( 3 ) extended the treatment to include stepwise formation of complexes, and Schaap and McMasters ( 4 )treated the case of mixed ligand complexes. The limitations of the original DeFord-Hume treatment have been pointed out by a number of workers (5,6) particularly in regard to the accuracy required in Ellz measurements. The advantages of using least-squares curve fitting programs in processing the data as opposed to the originally proposed graphical procedure have been pointed out (5, 7), and this is now standard practice. A comprehensive error analysis has been given by Klatt and Rouseff (5). These workers also addressed the question of necessary and sufficient conditions for detecting the presence of mixtures from polarographic data Present address: Faculty of Science, Zagazig University,Zagazig, Egypt. 0003-2700/84/0356-2401$01.50/0
and pointed out that although polarographic and potentiometric determinations of stability constants depend on the same functional relationship the potentiometric method is generally more powerful because it allows studies over a wider concentration range. Descriptions of current-potential curves in the absence of the restriction that a large excess of ligand be present have been presented by a number of workers (8-12). In earlier work (13) we reported results of calculated linear scan (LSV) and cyclic (CV) curves in the absence of excess ligand and suggested possible applications of these results in characterizing single complexes and mixtures of two complexes in systems where the ligand is not present in excess. Subsequently it has been shown that these simulations reliably predict the form of LSV and CV curves for previously wellcharacterized systems with a single predominate complex (14). These simulations were somewhat cumbersome and time consuming and were limited to the case with P,[X]q >> 1where p, is the overall formation constant Pq
=
[MXqI
[MI [XIq
and [XI is the free-ligand concentration (charges omitted). The purpose of this report is to present a much simpler and more versatile simulation and to compare calculated and experimental curves. In addition, applications of this approach in the determination of p1values for two complex systems in which the DeFord-Hume treatment is not applicable will be given.
EXPERIMENTAL SECTION All current-potential curves were obtained by using an IBM EC/225 voltammetric analyzer with a 7424 M x-y recorder. A conventional three-electrode cell with a saturated sodium chloride calomel electrode (SSCE) was employed. The working electrode in linear scan and cyclic measurementswas a PAR 9323 hanging mercury drop electrode (HMDE),and the electrode area normally 0 1984 American Chemlcal Society
2402
ANALYTICAL CHEMISTRY, VOL. 56, NO. 13, NOVEMBER 1984
employed was 0.0408 cm2. The scan rate was 100 mV/s in all data reported. Tast polarography with a controlled drop time of 1 or 2 s was employed for measurement of diffusion coefficients and for measurement of Eo’ which was taken as the measured Ellz in the absence of ligand. Reagent grade chemicals, deionized water, and triple-distilled mercury were employed in all cases.
RESULTS AND DISCUSSION Description of t h e Simulations. The calculations followed the general procedures outlined by Bard and Faulkner (15) and were based upon several assumptions that are commonly made. (1)Chemical and electrochemicalprocesses are both fast on the experimental time scale, so that equilibrium is maintained at all times. (2) All species diffuse at the same rate. (3) Concentration profiles may be approximated by a series of thin layers of both the solution and the electrode material, with mass transfer occurring from one layer to another by diffusion. (4) The electrode is planar. For each simulation values are entered for total metal (CM) concentrations, total ligand (Cx)concentration, the number of electrons for the electrode process (n),a starting potential (E)relative to E O ’ , a switching potential, a potential step size, and values of pq’s for the complexes assumed present. Each cycle of the simulation consists of the following steps: (1) Establish electrochemical and chemical equilibrium in the surface layer. (2) Approximate the diffusion out to layer 4.2k’/2 where k is the iteration number and compute the current function from the surface flux of I&.(3) Reestablish chemical equilibrium in the surface layer. (4) Increment the potential. (5)Return to step (1). For a series of mononuclear complexes in chemical equilibrium M + qX * MX,
(1)
w
-
0
.
4 1 0.0
0.0
0.0
E
I
where C is 0.0389 from the Nernst equation and AE is the potential applied to the electrode, relative to Eo’ for the electrode process. When it is demanded that both chemical and electrochemical equilibrium are simultaneously established it can be shown that at the electrode surface
In the program used, the value of [XI is refined by using the Newton-Raphson method. The initial guess required for [XI varies according to the particular chemical system and the values for the 0s;. A single subroutine was used in steps 1and 3 of each cycle by setting R in eq 4 to zero when only chemical equilibrium is computed. Because of the assumption that all diffusion coefficients are the same, the total ligand concentration in each solution layer is unchanged and equal to Cx.This means that it is only necessary to consider the diffusion of one species, I&, to obtain a reliable simulation of the current function. As a result of
I
-0.1 -0.2
Flgure 1. Calculated CV curves for the MX, system at Cx/Cu = 2. Values of log @, from left to right: 4, 6, 7, 8, and 12.
-
0
4
h
-‘-0:2
-0:3’
E -E o’(Volts)
-
+ .2
-0.3
Flgure 2. Calculated CV curves for the MX, system for log 0, = 12 as a functionof Cx/Ct,,,. Values of Cx/CMfrom left to rlght: 2,4, 10, and 50.
Table I. Calculated Curve Parameters for the MX2 System at CX/CI = 2 for Different Values of P2”
[MI [XI9
Electrochemical equilibrium establishes the ratio
io
-EoiVolts)
0,= the average number of ligands bound per metal at equilibrium is given by
I
82
current function at max
104 106 107 108 1012
0.446 0.431 0.385 0.357 0.347
peak potentials, mV vs. Eo‘
EPC
E,.
-14 -22 -42 -73 -193
15 10 3 -6 -134
oCalculated for two-electron reduction. these simplifying assumptions and the reduction in the calculations that must be performed, it is practical to carry out these simulations on an 8-bit microcomputer using an interpreted language. In typical simulations the starting potential was 100 mV and the potential step size was 1mV. The switching potential was set at least 75 mV beyond the peak potential. An n value of 2 was employed in all simulations presented. A copy of the program is available from the authors upon request. Calculated Curves for Single-Complex Systems. The MX2 System. Calculated CV curves for the MX2 system in which the ligand is present in stoichiometricamounts (CX/CM = 2) are shown in Figure 1. It can be seen that a t low values of Pz the curves show a small negative shift from Eo’ (or Ell2)and that the maximum value of the current function (d2X(at))deviates only slightly from the theoretical value (0.446)for uncomplicated reversible systems. Curve parameters for this system are given in Table I. The value of the current function of 0.347 at p2 = 10l2is identical with that previously reported (13)for the limiting case. Figure 2 shows changes in calculated curves with changes in total ligand concentration for a strong (p2 = MX2
ANALYTICAL CHEMISTRY, VOL. 56, NO. 13, NOVEMBER 1984
-- .
-.
II
I
0.0
0.0 E .Eo'(volrs)
0.:
1
I
I
I
I
1
1
2
3
4
5
6
I
I
8
Figure 3. Calculated CV curves for the MX system at C x / C M= 0.5 as a function of 6,. Values of log from left to right: 3, 4, 5, 6, and
@,
7.
I
I
....................
+
0.4. N -,
-0.4-t
-
0.0
"
E
0.0
"
0.0
I
-Eo'(volts)
Figure 4. Calculated CV curves for the MX, system at Cx/CM= 1 as a function of p,. Values of log & from left to right: 6, 7, 8, and 9.
system. As would be expected the maximum value of the current function attains the value of 0.446 and the CV curve exhibits the "normal" shape 59 EpO - Epc= - mV
n
at a c X / c M value of around 50. The MX System. Similar results are obtained assuming a single 1:l complex. In this system at stoichiometric ligand concentration the maximum value of the current function decreases to 0.385 at @I > lo6. Calculations at Substoichiometric Ligand Concentrations. In cases where the ligand is present in less than stoichiometric amounts calculated curves show splitting dependent on the value of @., Figures 3 and 4 show calculated curves for the MX case at C X / C M = 0.5 and for the MX2 case a t c x / c M = 1. The results are in general agreement with the calculations of Elenkova and Nedelcheva (12) for generalized currentpotential curves which predict a split cathodic wave for the MX2 case with c x / c M = 1when fl2 > loll. Application of simulations allows a much more accurate calculation of conditions under which splitting is expected to be observed. Mixtures of Complexes. Since stepwise formation of complexes represents the usual case in real systems it is essential that such systems be treated. Specific questions of interest would include the form of CV curves when mixtures are present, effect of Ppvalues on the forms and characteristic potentials, methods for detecting the presence of mixtures, and possible measurement of P, values a t low ligand concentration. Results of simulations indicate that, as would be expected, the forms of CV curves depend on the ratio of the two p, values. Figure 5 shows the value of maximum current function as a function of this ratio at c x / c M = 2. At low values of p2/& the maximum value of the cathodic current function is 0.420,
k
0.3-
3 I
I
I
I
I
2403
2404
ANALYTICAL CHEMISTRY, VOL. 56, NO. 13,NOVEMBER 1984
I
A
0.31
A
A
I
I
I
I
I
I
-3.5
-3.0
-2.5
CM
Flgure 7. Calculated and experimental values of current function at maximum C,IC, = 2,varying total concentration. Solid lines calculated, (A)Pb(diene),(0)Cu(en), (m) Cu(ox)+ Cu(ox),. Values of log 0, employed In calculat1on:for Pb(d1ene)log PI = 9.5;for Cu(en), log p2 = 20;for Cu(ox) Cu(ox), log p1= 5.4, log 0,= 9.2. Supporting electroyte 1 M NaN0, for Cu(ox) Cu(ox),and Pb(diene);1 M KCI for Cu(en),.
+
+
coefficients that give maximum current functions in these cases within 2% of 0.446. In order to compare predictions from simulations with experimental observations we have examined several previously studied systems. The Pb(II)-diethylenetriamine (dien) system has been reported to form a single complex with log p1 = 9.5 (16).The Cu(I1)-ethylenediamine (en) system has been studied by a number of workers (4,14,17-21).Values of pz reported from polarographic measurements range from log b2 = 19.7to 20.5. The value of log p1= 10 has been recently reported from potentiometric measurements (21).The Cu(11)-oxalate (ox) system is reported (4,22-25) to form two complexes with log pz values of 9.2-9.5 and log p1of 5.7-6.6. Figure 7 shows results of calculated and experimental values of maximum current fundion in the case where c X / c M is kept constant at 2 and the total concentration is varied. Experimental values for Pb(diene) and Cu(en)2follow closely the calculated value characteristic of single 1:l and 1:2 complexes, respectively, while the Cu(I1)-ox system exhibits changes indicative of a transition toward a mixture in which the CuOX)^ concentration is increasing at higher total concentration. Values of log p1and log p2 employed in the simulation for Cu(I1)-ox were 5.4 and 9.2. Similar results are obtained by comparing calculated and observed values of E, - Ep,2as an alternate measure of curve shape. It would thus appear that this technique does provide a rapid method for detecting mixtures if the ratio of p, values is such that a transition of the type shown in Figure 5 is predicted. A useful application of these results is in the measurement of p1values in systems with two complexes for cases in which the normal potential shift technique is not applicable. This in effect allows extention of polarographic methods to lower ligand concentrations, After establishing that a mixture is present the value of pZmay be determined by conventional potential shift analysis at Cx/Cw > 50. The value of is then obtained at low ligand concentration ( c X / c M = 2) by using a curve similar to that shown in Figure 5 as a working curve. Two systems were examined by using this approach. Potential shift analysis for the Cu(I1)-ox system gave a value of log pz = 9.2 in good agreement with previously reported values (4,22-25). Application of nonlinear least-squares analysis indicated that no significance could be placed on values of PIobtained from these data. This is in agreement with the findings of Leggett (7). At c x / c M = 2 this system gave a value of 0.370 for the maximum current function based on a measured D value of 6.68 X lo4 cm2/s. This corresponds to a value of log p1 of 5.5. When a working curve based on Ep- E P pis used a value of log & = 5.3 is calculated. Both are in better agreement with
-0.41
I
I I
I
0.0
-0.1
I
I
I
-0.05 -0.1
I
I
-0.2
E -EoiVolts)
Figure 8. Calculated and experimental CV curves for the Cu(I1)-ox system. In 1 M NaNO,. Left: C,, = 5 X lo-' M, Cox= 1 X M. Right: C,, = 3 X lo-, M, Cox = 6 X M. Values used in calculation: Eo' = -0.017 V vs. SSCE, log PI = 5.4, log p2 = 9.2. Polnts are experimental.
I
0.AL
-0.41
1
0.0
I
I 1
'
1
-0.1 0.0 E -Eo'(Volts)
I
1
-0.1
-0.2
Figure 9. Calculated and experimental CV curves for the Cu(I1)-ox system. Left: CoxIC,, = 0.5. Right: C,,IC,, = 1. Values used in calculations: Eo' = 0.017 V vs. SSCE, log PI = 5.4,log p2 = 9.2. Broken line Calculated for log PI = 6.6, log 0,= 9.2. Polnts are experimental.
I
0.4c
E
- EoiVolts)
Flgure 10. Calculated and experimental CV curves for the Cd(I1)-prop system. Left C,,IC,, = 0.5. Right: Cpmp/C,, = 1. Values used in calculations: Eo' = -0.585 V vs. SSCE, log 8, = 5.4, log p2 = 9.7. Broken line calculated for log p1 = 7.9,log 6, = 9.7. Points are experlmental. the value of 5.4 from potentiometric measurements than the polarographic values of 5.7-6.6. Similar treatment of the Cd(I1)-propylenediamine (prop) system gave log PI= 5.6 in much closer agreement with the potentiometric value (4)of 5.4 than the polarographic values of 7.9-9.0which have been reported (27-29). Using the simulations one can readily make direct comparisons between calculated and observed CV curves. Figure 8 shows such a comparison for the Cu(I1)-ox system at C x / C M = 2. Curves at substoichiometric ligand concentration are as noted quite sensitive to @, values. Figure 9 shows curves for the CU(II)-OX system at &/CM = 0.5 and at c X / c M = 1. For comparison a calculated curve using ow experimental 02 value and a reported log p1value of 6.6 (23)is included in the figure. Figure 10 shows similar results for the Cd(I1)-prop system. In both cases the p1value which gives best agreement with
Anal, Chem. 1904, 56, 2405-2407
experiment closely conforms to & values reported from potentiometric measurements. In summary we have shown that by use of simulations it is possible to extend CV analysis for reversibly reduced complex systems to systems with low ligand concentration. It should be noted that this approach must be applied with some caution. At very low ligand concentration the possibility is greater for complications due to complexes between the metal ion and hydroxide and/or anions of the supporting electrolyte. In the cases studied the good agreement between predicted and experimental results would indicate that this is not a serious problem. Extension to more complex systems and to other electroanalytical techniques is in progress.
LITERATURE CITED (1) Von Stackelberg, M.; Von Freyhold, H. 2.Elektrochem. 1040, 46. 120. (2) Llngane, J. J. Chem. Rev. 1941, 29, 1. (3) DeFord. D. D.; Hume, D. N. J . Am. Chem. SOC. 1051, 7 3 , 5321. (4) Schaap, W. B.; McMasters, D. L. J . Am. Chem. SOC. 1061, 83, 4699. (5) Klatt. L. N.; Rouseff, R. L. Anal. Chem. 1070, 42, 1234. (6) Irvlng, H. Adv. Polarogr. 1960, 1 , 42. (7) Leggett, D. J. Talanta 1080, 27, 767. (8) Koryta, J. Pfog. Polarogr. 1062, I , 295. (9) Buck, R. P. J. Elecfroanal. Chem. 1063, 5 , 295. (IO) Butler, C. 0.; Kaye. R. C. J. Nectroanal. Chem. 1964, 8 , 463.
2405
(11) Macovschl, M. E. J . Nectfoanfll. Chem. 1968, 16, 457. (12) Elenkova, N. 0.;Nedelcheva, T. K. J . Elecfroanal. Chem. 1976, 69, 305. (13) Spell, J. E., 11; Phllp, R. H., Jr. J. .€/ecffmnal. Chem. 1080, 712, 281. (14) Kllla, H. M.; Phllp, R. H., Jr. J . Elecfroanal. Chem., in press. (15) Bard, A. J.; Faulkner, L. R. “Electrochemlcal Methods”; Wlley: New York, 1980; Appendix B. (16) Nikolov, T.; Isolfova-Tsaneva, M.; Lyakov, N. Mefallurgiya (Lenlngrad) 1977. 23(a). ,-,. 1314. (17) Laltlnen. H. A.; Onstott, E. I.; Bailar, J. C., Jr.; Swann, S. J . Am. Chem. SOC.1940, 7 1 , 1550. (18) Spike. C. G.; Parry, R. W. J. Am. Chem. SOC. 1053, 75, 3770. (19) Basolo, F.; Murmann, R. K. J . Am. Chem. SOC. 1952, 74, 5243. (20) Cotton, F. A.; Harris, F. E. J . Phys. Chem. 1055, 59, 1203. (21) Avdee, F. A.; Zabronsky, J.; Stutlng, H. H. Anal. Chem. 1983, 55, 298. (22) Shah, S. K.; Suyan, K. M.; Gupta, C. M. Talanta 1080, 2 7 , 455. (23) Placeres, C. R.; Leon, J. J.; TruJlllo, J. P.; Monterlango, F. G. J. Inorg. Nucl. Chem. 1081, 43, 1681. (24) McMaster, D. L.; DiRalmondo, J. C.; Jones, L. H.; Lindley, R. P.; ZeRmann, E. W. J. Phys. Chem. 1062, 66, 249. (25) Clavatta, L.; Vlllafiorlta, M. Gazzerta 1055, 95, 1247. (26) Carlson, 0. A.; McReynolds, J. P.; Verhoek, F. A. J. Am. Chem. SOC. 1945, 67, 1334. (27) Gupta, K. D.; Gaghel, S. C.; Gaur, J. N. Monatsh. Chem. 1070, l l a , 657. (28) Sharma, R. S.; Gaur, J. N. J. Elecfrochem. SOC.India 1078, 2 7 , 261. (29) Maheshwarl, A. K.; Jaln, D. S.; Gaur, J. N. Monafsh. Chem. 1975, 106, 1033.
.
~
RECEIVED for review March 26,1984. Accepted June 11,1984.
Determination of Nitrite Ion and Sulfanilic and Orthanilic Acids by Differential Pulse Polarography S. T. Sulaiman Department of Chemistry, College of Science, University of Mosul, Mosul, Iraq
Nltrlte Ion can be determlned wlth a hlgh degree of accuracy and sensltlvlty by dlfferentiai pulse polarography utlilzlng the rapld and quantitative reactlon between nltrlte Ion and sulfanlllc acld or orthanlllc acid at pH 1.5. The experimental detectlon limit Is shown to be 8.6 X lo-’ M (as NO,-) In slmpie aqueous soiutlon. The method Is further used to determlne concentratlonsof sulfanlllc acld down to 4 X lo-’ M and orthanlllc scld down to 1.6 X lo-’ M under optlmum condltlons.
the quantitative reaction of nitrite with diphenylamine to yield diphenylnitrosamine (DPN), with a practical working limit of 4.6 ppb. DPN is toxic, may be a carcinogen, and may act as nitrosating agent via transnitrosating (8). The purpose of this paper is to demonstrate the differential pulse polarographic behavior of diazonium salts obtained from diazotation of sulfanilic and orthanilic acids for the trace determination of nitrite. The method is simple, sensitive, and rapid.
The nitrosation of most aliphatic and aromatic amines with nitrite leads to the formation of N-nitrosamines, many of which have been shown to be potent carcinogens (1-3), and, when correlated with other nitrogen forms in water, can provide an index of organic pollution (4). Thus a sensitive and rapid method for the determination of nitrite is desirable. Numerous methods have been proposed for the determination of nitrite, among them are the spectrophotometric (4, 5), ion selective (3), and polargraphic methods (6-8). The spectrophotometric methods have limited sensitivity and dynamic range; most of them depend on unstable colors and are time-consuming. Differential pulse polarographic determinations of nitrite have been reported either by direct measurement of nitrous acid (6) (detection limit 0.5 ppm) or indirectly by the enhancement of the ytterbium peak (7). A detection limit of 46 ppb of nitrite ion was achieved under ideal conditions. Another report appeared by Harrington et al. (B), in which nitrite was determined by differentialpulse polarography using
EXPERIMENTAL SECTION Apparatus. Polarographic curves are recorded with a Metrohm Polarecord E 506 in conjunction with an E 505 polarography stand equipped with mechanical drop timer. A three-electrode system was used the working electrode was a dropping mercury electrode; the reference electrode was an Ag/AgCl, KCl with ceramic liquid junction; and the counterelectrode was a platinum wire. The differential pulse mode was used with a 100 mV pulse, a 2-5 drop time, and scan rate of 3 mV s-l, except where otherwise indicated. All polarographic measurements were performed at room temperature (20 O C ) . The solution was deareated by passing through it a slow stream of helium for 15 min. Reagents. All chemicals used were of analytical grade. Standard nitrite solution 2.17 X M was prepared by dissolving 0.1499 g of sodium nitrite in twice-distilled deionized water, followed by the addition of a pellet of sodium hydroxide and 1 mL of spectroscopic grade chloroform, and diluted to 1L. Dilute solutions were prepared from this stock solution by appropriate dilution or by direct pipetting into samples. Sulfanilic acid and orthanilic acid were obtained from BDH Stock solutions M) of each compound were prepared in deionized water. A series of the modified Britton-Robinson universal buffer (BRB) solutions (pH 1.8-2.0) were prepared as given by Britton
0003-2700/64/0356-2405$0 1.50/0 0 1984 American Chemlcal Soclety
