Anal. Chem. 1983, 55,963-965
electrode used for thermoionic diodes. CONCLUSION Detailed understanding of the response of the LEI pulse to electric field and laser beam position is an important step in ensuring the sensitivity and the accuracy of the method. This is especially true since easily ionized substances have been demonstrated ,to modify the electric field in the flame. ACKNOWLEDGMENT The authors ;are greatly indebted to J. C. Travis for helpful discussions, access to his team's theoretical work prior to publication, anid comments on our manuscript. Acknowledgment is made to G. Baudin and G. Delarue for help and suggestions received throughout the course of this study. We are grateful to B. Fleurot for advice in the high-speed electronic devices realization. LITERATURE C I T E D (1) Penning, F. hd. Physlca (The Hague) 1028, 8, 137-140. (2) Green, R. B.; Keller, R. A.; Luther, G. G.; Schenck, P. K.; Travis, J. C. Appl. Phys. Lett. 1 9 W 29, 727-729. (3) Turk, G. C.; Travis, J. C.; De Voe, J. I?.; Oblaver, T. C. Anal. Chem. 1078, 5 0 , 817-820. (4) Turk, G. C.; 'Travis, J. C.; De Voe, J. R.; O'Haver, T. C. Anal. Chem. 1979, 51, 1890-1896. (5) Turk, G. C.; Mallard, VI/. G.; Schenck, P. K.; Smyth, K. C. Anal. Chem. W7g, 51. 2408-1410. (6) Green, R. B.; Havrilla, G. J.; Trask, T. 0. Appl. Specttrosc. 1980, 34, 56 1-569.
963
(7) Gonchakov, A. S.;Zorov, N. B.; Kuzyakov, Yu. Ya.; Matveev, 0.I . Anal. Lett. 1079, 12, 1037-1048. (8) Schenck, P. K.;Travls, J. C.; Turk, G. C.; O'Haver, T. C. J . Flhys. Chem. 1981, 85. 2547-2557. (9) Turk, G. C. Anal. Chem. 1981, 5 3 , 1187-1190. ( 1 0 ) Mavrodineanu, R.; Boiteux, H. "Flame Spectroscopy"; Wiley: New York, 1965. (11) Smyth, K. C.; Mallard, W. G. Combust. Sci. Techno/. 1981, 26, 35-41. (12) Mallard, W. G.; Smyth, K. C. Combust. Flame 1982, 44, 61-70. (13) Lawton, J.; Welnberg, F. J. "Electrlcal Aspects of Combustion"; Clarendon Press: Oxford, 1969.
T h i e r r y Berthoud* J a c e k Lipinsky Centre d'Etudes NuclBaires Fontenay aux Roses, France P i e r r e Camus Laboratoire Aim6 Cotton
CNRS Orsay, France Jean-Louis Stehle SOPRA Bois Colombes, France
RECEIVED for review May 17, 1982. Resubmitted November 15, 1982. Accepted December 30, 1982. This work is supported by the DGRST-MinistBre de la Recherche et de 1'Industrie under No. 80.7.0569 and 81.Y.0805.
Cationic Indicator Bases for the H+ Scale in Concentrated Sulfuric Acid Sir: The Hammett acidity function, HO( I ) , was introduced to describe the prototropic behavior of neutral bases in concentrated mineral acid. The Ho function measured the tendency of the imediunn to transfer a proton to an uncharged base. It was felt (2), however, that this function would not adequately describe the prototropic equilibria of charged bases such as singly or doubly charged cations. Thus attempts were made to establish the H+function based on the protonation of singly charged cationic bases. Theoretically, the H+function was expected to be more negative than the Ho function because of the electrostatic contribution of the charges on the indicator ions to their activity coefficients (3). One study ( 4 ) showed that plots of log ([base]/[acid]) vs. percent sulfuric acid for several anilines as well as 4-nitro1,2-phenylenediamine (inflection region in 30-55% sulfuric acid) and 4-aminoacetophenone (inflection region in 75-95% sulfuric acid) were parallel. From these results, these investigators concluded that the H+and Hoscales were either identical or differed by a small constant amount. Other workers (5) found that the second protonation of the aminopyridines followed the Ho function. In investigation of the prototropic behawior of singly charged cations from 0.02 to 15 M sulfuric acid (6),it was found that the H+ acidity function was slightly more negative than the Ho function up to about 7 M acid. Thereafter, the former became less negative than the latter and did not increase as rapidly with increasing molar acid concentration. Below 7 M acid, primary anilines were used to establish the H+ function. However, above 7 M acid quinoxalines, except for 3-nitr0-1,2phenylenediamine, were used as indicators. It seemed likely to us that the leveling off of the H+ function was due to the use of quinoxalirie cations.
In this work the second protonation of several cations in sulfuric acid was studied. The aim was to use these indicators to check the accuracy of and to extend the published H+ function to 18 M sulfuric acid using the protonation of primary arylamines as the exclusive class of indicator reaction. A reliable H+function is necessary in order to extend the hydration parameter treatment which was applied previously to carboxamides (7,8) and lactams (9) to charged bases other than primary amineEi. To date there has been no study of the protonation behavior of doubly positively charged cations. Hence a limited study of two such indicators was also conducted. EXPERIMENTAL SECTION Chemicals. Proflavine (Allied Chemical and Dye Corp., New York, NY), Thionine (K&K Laboratories, Plainview, NY), 1,2diaminoanthraquinone, 2,6-diaminopyridine, 3-nitro-1,2phenylenediamine,4-aminopyridine,2-amino-6-methylpyridirte, and 2-aminopyridine were used without further purification. 4-Bromo-6-nitro-1,2-phenylenediamine was synthesized from 2,6-dinitroaniline (Pfaltz and Bauer, Stamford, CT) by the following procedure. 2,6-Dinitroaniline was dissolved in acetic acid and excess bromine added to the solution. This mixture was then refluxed on a water bath for 2 h (10). The precipitate was filtered and shown to be 4-bromo-2,6-dinitroaniline by melting point (150-162 OC) (11) and NMR. This compound was dissolved in ethanol and treated with ammonium sulfide (12). This solution was heated on a water bath for an hour during which time it became deep red. The solution was concentrated under vacuu.m and the precipitate filtered. This was checked for purity by TLC using methano1:chloroform (10:90) as developing solvent. This precipitate was found to be 4-bromo-6-nitro-1,2-phenylenediamine by melting point (360 "C) and NMR. Method. A stock solution of each indicator was prepared in
0003-2700/83/0355-0963$01.50/00 1983 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 55, NO. 6, MAY 1983
Q64
Table I. Regression Coefficients for the Protonation of Cationic Bases in Sulfuric Acid at 2 5 "C Ho vs. log I slope i intercept std dev std dev
compound
1,2-diaminoanthraquinone 1.05 i 0.03 2,6-diaminopyridine 0.94 i 0.01 3-nitro-l,2-phenylenediamine 1.05 i 0.01 4-bromo-6-nitro-l,2-phenylenediamine1.02 i 0.01 4-aminopyridine 0.85 i 0.01 2-amino-6-methylpyridine A,,, 0.98 + 0.02 1.05 i 0.02 0.99 i 0.02 2-aminopyridine A,,, 0.98 t 0.01 0.97 i 0.02 A255 0.95 i 0.01 4 9 8 proflavine A,,, 1.31 i 0.04 1.10 i 0.02 1.12 i: 0.02 1.12 ?r 0.02 A460 Thionine 0.96 ?r 0.01 -1o.m
-7.m
, 4
-3.62 -4.12 -4.38 -5.22 -6.36 -7.45 -7.41 -7.46 -8.03 -8.09 -8.19 -3.07 -2.99
-3.12 -2.97 -4.46
t
i i i i i
i
0.02
0.01 0.01 0.01 0.01 i 0.01 i: 0.01 i 0.01 i 0.01 i 0.01 i 0.01
i
0.02 0.01 0.01 0.01
i
0.01
i i i
corr 0.993 0.999 0.997 0.999 0.999 0.995 0.994 0.996 0.999 0.996 0.998 0.988 0.996 0.998 0.996 0.996
0.85i 0.76 i 0.82 i 0.61 i 0.24 t
0.02 0.01 0.01 0.02 0.01
-3.58 -3.98 -4.19 -4.76 -5.29
0.01 0.01 i 0.01 t 0.01 i: 0.01 i i
lit. pK
0.992 0.998 0.996 0.988 0.994
-4.10 -6.50
-8.1
1.01 i 0.98 i 0.98 i 1.06 i
0.01 -3.11 0.02 -3.08 0.01 -3.19 0.01 -3.07
f. i i i
0.05 0.999 0.01 0.997 0.01 0.998 0.01 0.998
Table 11. Revised Values of H + as a Function of Molar
/ +Concentration of Sulfuric Acid at 2 5 "C
/ t
I
[HZSO, M
I
/+-
7.00 7.40 7.80 8.00 8.40 8.80 9.00 9.40 9.80
A
1
10.00 om
H , (published) vs. log I slope i intercept i std dev std dev corr
361
72. 'lala,
.mco
13 a1 t
31 0
-4 4
.E
10.40 10.80 11.00 11.40 11.80 12.00
io
kr'C4
Figure 1. Plot of H+ vs. molar concentration of sulfurlc acid: (0) publlshed data of Vestesnlk et al. (6); (0)revised H+ scale this work.
sulfuric acid or water depending on the solubility of the indicator. Working solutions were prepared in 10-mL volumetric flasks as previously described (7,8). The procedure for the recording of spectra was also previously described. No spectral shifts due to medium effects were observed.
RESULTS AND DISCUSSION The dissociation constants for each indicator were determined by making plots of Ho (13) and H+ (6) vs. log I , the logarithm of the molar concentration ratio [BH+]/ [BHZ2+]. log I was calculated from the absorbances as previously described (8). The slopes and intercepts and their standard errors were obtained by the method of least squares. These are shown in table I. New values of H+as a function of molar concentration of sulfuric acid are given in Table 11. These values of H+ are plotted in Figure 1. From the definition of a Hammett acidity function, H, [base] H, = pK,,id log [acid]
+
a plot of H, vs. log I should be linear with unit slope if H, adequately describes the protonation behavior of the base. In Table I1 the slopes of plots of Ho vs. log I and of H+ (from the present work) vs. log I for all the singly charged cationic indicators in this study are unity, except for 4-aminopyridine. However, plots of H+ (taken from the previously published scale (6)) vs. log I have slopes less than unity. This would confirm that the leveling off of the published values of H+as a function of H2S04concentration is due to the use of qui.noxaline type indicators. In the case of the doubly charged
I,
H+ revised
-3.34 -3.58 -3.78 -3.88 -4.08 -4.28 -4.38 -4.58 -4.82 -4.93 -5.16 -5.38 -5.50 -5.75 -6.00 -6.12
[H,SO, M
I,
12.40 12.80 13.00 13.40 13.80 14.00 14.40 14.80 15.00 15.40 15.80
16.00 16.40 16.80 17.00
H+ revised
-6.36 -6.61 -6.71 -6.95 -7.20 -7.32 -7.56 -7.80 -7.91 -8.1 5 -8.39 -8.50 -8.74 -9.03 -9.18
cationic bases, plots of H+ vs. log I for proflavine have unit slopes, whereas plots of Ho vs. log I have slopes greater than unity. For Thionine, a plot of Hovs. log I has unit slope. These results suggest that the electrostatic contributions to the activity coefficients of solutes which are important in dilute sulfuric acid solutions exert little or no influence in concentrated sulfuric acid, that is greater than 7 M. From the limited study of the doubly charged cationic bases, it would appear that the H+and the H++acidity functions are collinear. More extensive studies would be required to confirm the above observation. The anomalous behavior of 4-aminopyridine may be explained by the fact that this compound has been postulated to exist as an amidine rather than a primary amine (13).
Registry No. 1,2-Diaminoanthraquinone,1758-68-5;2,6-di3694-52-8; aminopyridine, 141-86-6;3-nitro-l,2-phenylenediamine, 4-bromo-6-nitro-1,2-phenylenediamine, 84752-20-5; 4-aminopyridine, 504-24-5; 2-amino-6-methylpyridine,1824-81-3; 2aminopyridine,504-29-0;proflavine A, 92-62-6;thionine, 581-64-6; sulfuric acid, 7664-93-9; 2,6-dinitroaniline, 606-22-4; 4-bromo2,6-dinitroaniline, 62554-90-9.
LITERATURE CITED (1) Hammett, L. P.; Deyrup, A. M. J . Am. Chem. SOC. 1932, 5 4 , 2721-2739. (2) Hammett, L. P.; "Physical Organic Chemlstry"; McGraw-Hill: New York, 1940; pp 267. (3) ~, Brand. J. C. D.: Hornlna, W. C.; Thornley; M. B. J . Chem. SOC. 1952, 1374-1303. (4) Bonner, T. G.;Lockhart, J. C. J . Chem. SOC. 1957, 364-367.
Anal. Chem. 1983, 55,965-967 (5) Brlgnell, P. J.; Johnson, C. D.; Katritzky, A. R.; Shaklr, N.; Tarham, H. D.; Walker, G. J. Chem. SOC. 1067, 1233-1237. (6) Vestesnlk, P. J.; Bielavsky, J.; Vecera, M. Collect. Czech. Chem. Commun. 1068, 33, 1687-1692. (7) Lovell, M. W.; Schulman, S. G. Anal. C h h . Act8 1081, 127, 203-207. (8) Lovell, M. W.; Schulman, S. G. Int. J . Pharm. 1082, 7 1 , 345-354. (9) Schulman, S. G.; Vogt, B. S. J. W y s . Chem. 1081, 85,2074-2079. (IO) Deorha, D s.; Joshi, s. s.; Masesh, v. K. J. Indian Chem. soc. 1062, 39,534-536. (11) Robertson, G.R. "Organic Syntheses"; Wiley: New York, 1941;Collectlve Vol 1, pp 52-53. (12) Jorgenson, M. J.; Hartter, D. R. J. Am. Chem. Sot. 1063, 85, 878-883.
965
(13) Albert, A.; Goldacre, R. Nature (London) 1044, 753,467-4139,
Michael W. Lovell Stephen G . Schulnnan* College of Pharmacy University of Florida Gainesville, Florida 32610 for review November 8, 1982. Accepted January 21, 1983.
Nitric Oxide Interference in the Determination of Dissolved Oxygen by the Azide-Modified Winkler Method Sir: The reliable determination of dissolved oxygen in natural waters is essential to both water treatment engineering efforts and studies of the redox stability of chemical species in aqueous solution. The development of electrometric methods for dissolved oxygen determination has made such measurements3 rapid, precise, and convenient ( I , 2). Difficulties with the use of these sensors include membrane or electrode poisoning, the necessity to calibrate electrode response vs. parallel Winkler iodometric titrations, and the potential bias introduced by sample collection or handling efforts ( 3 , 4 ) . The latter problems are particularly acute for groundwater simples which are frequently not in equilibrium with the atmosphere, may contain high levels of dissolved sulfides or H2S, and may undergo considerable depressurization during pumping and collection (4, 5). The determination of low levels of dissolved oxygen in these systems requires careful attention to sampling technique and the analyst should also be aware of potential analytical interferences from reduced chemical species. Nitrite ion, an intermediate in microbial transformations of ammonia and nitrate, is known to significantly affect the results of the Winkler oxygen method. The Alsterberg modification was devised to prevent nitrite interferences with oxygen determinations under the reducing conditions encountered in wastewater and activated sludge systems (6). Subsequent improvements in methodology by Carpenter (7) extend the usefulness of the azide-modified Winkler method to natural waters with oxygen levels as low as 0.05 mg.L-' in the presence of up to 5 mg.L-l nitrite. In our studiles of the chemistry of contaminated groundwater systems, we have encountered another source of chemical interference with the azide-modified Winkler method (8). In this instance, the massive contamination of an alluvial aquifer had occurred, resulting in ammonia and nitrate concentrations averaging 2000 and 1300 mg.L-l, respectively. Microbial transformations of both ammonia and nitrate were indicated by the field results. The oxygen-measurement problem was discovered when apparently high oxygen levels were measured in samples which exhibited nitrite concentrations in excess of 20 mg.L-l and ferrous iron concentrations averaged -0.05 mg-L-l, five times greater than the practical detection limit. The coexistence of these reduced species in oxygen-oversaturated groundwater samples is inconsistent with their known redox stability (9). We, therefore, initiated a study of the potential bias on Winkler oxygen results by other recognized intermediates of denitrification OF nitrate reduction processes in soil and sediment/water systems. The effects of hydroxylamine, nitrite, nitric oxide, and nitrous oxide were investigated. Nitrous oxide has been observed a t concentrations up to 6-8 mg.L-l in soil waters receiving similar 0003-2~00/83/0355-0965$0 1.50/0
imputs of inorganic nitrogen at pH and redox levels similar to those a t our field site (IO).
EXPERIMENTAL SECTION The background electrolyte chosen for the experiments was 5X M NaHC03 made up in a double-distilled water. Solutions of hydiroxylamine (NH20H)and nitrite were prepared gravimetrically. Stock solutions saturated with the gases (WzO, NO) were prepared at 24 "C and at atmospheric pressure. Tabulated values for solubility at this temperature were used as stock concentrations. Dilution of these stock solutions was performed by rapidly delivering the stock aliquots below the surface of the dillution water. Reagent chemicals were ACS reagent grade and the gases were delivered from lecture bottles of > B Y 0 purity (Matheson Scientific). Experimental concentrations ranged from 0.5 to 50 1mg.L-l. Parallel experiments were run using airsaturated and nitrogen-purged background solutions. This was done to determine if any observed interferences were related to the oxygen content of the sample. No significant differences were noted in the reriults after corrections were made for dilution. All determinations, were made in duplicate at both 16 and 24 "C following the azide-modifiedWinkler method (7). Once the solutions were prepared in 300-mL glass-stoppered bottles, the manganous sulfate and alkaline azide reagents were added and the samples were stoppered and agitated. After initial settling of the resultant precipitate, the samples were agitated again. After storage periods ranging from 0.5 to 8 h, the samples were acidified and excess iodine was back-titrated with 0.0109 N NazS203. Thiosulfate solutions were restandardized daily vs. standard potassium biiodate [KH(I03),] solution. Both starch and amperometric end point detection methods were used in the titrations. Experiments were also run to determine the usefulness of increased sodium azide (NaN,) concentrations on observed interferences.
R,ESULTS AND DISCUSSION The accuracy of the azide-modified Winkler analysis of oxygenated samples relative to biiodate standards was approximately 5% at the 5 mg 02.L-l level. Reproducibility on untreated (not spiked) samples from the experimental runs averaged *2% relative standard deviation. The precision of treated samples averaged *8% relative standard deviation. No significant decreases in pH were observed in the final dilutions of nitric oxide stock solutions. This ensured that alkaline conditions persisted during storage and that, under the conditions of the test, nitric oxide was not appreciably oxidized to nitrous acid. No significant positive bias on oxygen determinations was noted for hydroxylamine or nitrous oxide levels between 0.5 and 50 mg.L-l. Nitrite, at levels in excess of 5 mg.L-l (0.11 mM), resulted in a slight positive error equivalent to 0.01 mg 02.mg-l NO2-. The azide masking agent was in 10-fold excess over nitrite at 0 1983 American Ctiemlcal Society
