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.
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(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 dillutionwater. 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-modified Winkler 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
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ANALYTICAL CHEMISTRY, VOL. 55, NO. 6, MAY 1983 DISSOLVED NITRIC DXlDE VERSUS WlNKLER DISSOLVED OXYGEN ERROR
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Figure 2. Effect of azide addition on the interference of nitric oxide in the Winkier dissolved oxygen determination. Results are shown for experlmental runs at 5.1 ( 0 )and 18.7 mg N O C ' (0)after 4 h of storage at 24 O C . DISSOLVED NITRIC OXIDE (mg ' L - l I
Figure 1. Effect of dissolved nitric oxide concentration on the error in Winkler oxygen determination. Results are shown for 24 O C experimental runs for three storage periods between fixing and acidification prior to titratlon. this concentration. Doubling the amount of azide in the alkaline-KI reagent can minimize nitrite effects proportionately up to 10 mg.L-'. Even under strongly reducing conditions it is unlikely that this level would be encountered in all but the most contaminated environments or in specialized laboratory applications. Nitric oxide exerted a much greater positive bias on the oxygen determinations. The results of a series of 24 "C runs are shown in Figure 1for NO levels between 0.5 and 30 mg.L-l. The magnitude of the error ranged from 0.1 to 0.6 mg O2-mg-l NO and was a strong function of the storage period between fixing and acidification prior to titration. The magnitude of the error increased in the initial 4 h of storage at which point the effect of nitric oxide leveled off at least up to an 8-h period. Due to the relative magnitude of the effects of low levels of NO and the decreased precision of the overall determination (*8% relative standard deviation) when NO was present, we were unable to detect any significant systematic error below 1 mg N0.L-'. All treated samples exhibited vigorous effervescence and a deep red color on acidification. When the residual aliquots of acidified samples were allowed to sit stoppered in the laboratory, the characteristic violet vapor of elemental iodine accumulated in the headspace. We presume that the loss of iodine during the titration resulted in the poorer precision of the determination in the presence of nitric oxide. The 16 "C experimental runs showed somewhat better precision; however the magnitude of the error was about half that observed at 24 OC for corresponding concentrations and storage periods. We also investigated the use of azide addition to mask the nitric oxide error. The averaged results for runs at 5.0 and 18.7 mg N0-L-l over a range of azide additions (up to the solubility limit of -10 mM) are plotted in Figure 2. The percent reduction in the interference with added azide was determined relative to that observed with the reagent at the method's nominal azide concentration ( 1mM) in the fixed samples. Increased levels of azide effectively eliminated the NO interference (at NO levels between 2 and 20 mg.L-l) when at least a 20-fold molar excess was added with the alkalineiodide reagent. A t the higher excess additions of azide most of the samples showed increased effervescence and uncharacteristically abrupt titration end points most similar to blanks or biiodate standards. This may be the result of the reduction N
of Iz by the N3- (azide) ion in the acidified samples. The acid catalyzed reaction
2N3-
+ 2H+ + I2 F? 2HI + 3N2
(1)
has been reported in the literature (11). Nitric oxide, itself, is remarkably inert in slightly alkaline aqueous solution (12). However, the Winkler reagents provide several potentially reactive species which may account for the observed nitric oxide error in the determination. If NO were to oxidize the manganous ion under alkaline conditions to higher oxides and oxyhydroxides, this would explain the positive error observed in the iodometric titrations. This possible mechanism was investigated by omitting the manganous sulfate reagent from the modified Winkler procedure. Ten samples containing nitric oxide at levels between 2 and 20 mg.L-' were processed without the addition of manganous sulfate parallel to a set including this reagent. The resultant "oxygen" values (blank corrected) in these paired samples differed less than i 8 % from each other. Thus, it was apparent that, under the conditions of the determination, nitric oxide does not act analogously to dissolved oxygen. Reaction with manganous ion either does not occur or takes place too slowly to contribute to the yield of iodine. The possibilities remain either that nitric oxide is involved in the oxidation of iodide to iodine or that NO reacts with the thiosulfate titrant. It is difficult to judge between these possibilities, since the sample medium is quite complex. The most plausible mechanism is that acidification of the sample prior to titration permits the oxidation of nitric oxide by molecular oxygen to nitrogen tetraoxide
2Hz0
+ 2NO s Nz04+ 4H+ + 4e-
(2)
which readily occurs in acid solution (12). Nz04is an oxidizing agent roughly equivalent to bromine under these conditions and may in turn cause the oxidation of iodide to iodine. This possibility is supported by the observed agreement between the blue starch-I2 complex and amperometric end point detection methods. If iodine was not generated in solution, these two methods should have shown a substantial lack of agreement. We conclude that the use of the azide-modified Winkler O2method is inappropriate for contaminated water samples with low oxygen levels (
