the reaction mixture has a similar concentration with respect
to iodide and iodine. Interferences. First of all the influence of the titration vessel, made of Duran glass, and of the Teflon-coated stirrer bars was investigated. Noticeable positive errors arose only when a fresh stirrer bar was used for the first time. This phenomenon, which is probably due to the delivery of traces of fluoride from the Teflon material, is without significance after a few runs. The influence of several cations and anions on the described determination of fluoride was examined. The results listed in Table I indicate the maximum permissible amount (pmole) of the respective ion in our working range. Twe series of measurements were performed, according to whether the direct influence on the catalyst zirconium itself or the influence on the zirconium-fluoride system was to be investigated. In the first case, the solution of the interfering ion was added directly to the reaction mixture; in the second case, the solutions of fluoride and the ion under study were mixed before being transferred to the titration vessel. Care was taken that the pW value of the reaction system was the same as described for the fluoride determinations. The different ions in Table I are divided into several groups, according to their mode of interference. A. In this group ions are included which act as inhibitors similar to fluoride. B. These ions have a competitive effect with respect to the complex formation with fluoride.
Determination o elective Ion
C. Mercury forms complexes with iodide. D. Molybdenum is exhibiting a strong catalytic activity in the perborate/iodide system (I3,14). E. Different possibilities exist for the interfering effect of the ions in this group: catalytic activity, complex formation with fluoride and-in the case of iron (111)redox reaction with iodide. F. Practically no interference is shown by these ions, It is interesting to note that thorium has no catalytic activity under our experimental conditions, but acts sensitively as a “fluoride catcher.” This is shown by the fact, that without fluoride, thorium can be present up to 100 pmoles. The described method should be applicable to air pollution studies; too high concentrations of sulfate and phosphate in the air would be the main interferences. The serious interferences caused by some organic chelating agents (see Table I, group A) suggest their determination by the described method. ACKNOWLEDGMENT
We wish to express our thanks to C. Cuntz and H. Denzinger for helpful assistance. RECEIVED for review May 13, 1970. Accepted August 3, 1970. We are grateful to the “Deutsche Forschungsgemeinschaft” (D.K.) and to the “Deutscher Akademischer Austauschdienst” ( M A G . ) , which made this work possible by providing research grants.
ium Nitrate in Sodium Nitrite t rode Measurement
Douglas G. Gehring, William A. Dippel, and Robert S. Boucher Eastern Laboratory, Explosives Department, E. I . du Pont de Nemours & Company, Gibbstown, N . J . 08027 A nitrate selective ion electrode was used to measure small amounts of sodium nitrate in sodium nitrite after first destroying the nitrite with hydroxylamine sulfate (HAS). It was found that, under the conditions of the nitrite-hydroxylamine reaction, formation of 0.02% nitrate (as NaN03) occurs through disproportionation of nitrite. Nitrate concentrations were determined by relating the potentials of sample solutions to calibrating solutions of about the same ionic strength and composition. The precision (95% confidence) was calculated to be +0.02% at the 0.20% sodium nitrate concentration level. This method is rapid, uncomplicated, and has the advantage that a series of samples can be routinely analyzed at one time.
ACCURATE MEASUREMENT of nitrate in nitrite is difficult because nitrite is the most significant interference in nearly all the chemical and physical methods commonly employed for nitrate determination ( I ) . This problem is magnified in the case of sodium nitrite salt where the nitriteinitrate ratio is normally greater than 200. The failure to find nitrate listed as an impurity on the labels of commercial reagent grade nitrites is witness to the analysis difficulties, particularly since (1) A. P. Kreshkov, A. N. Yarovenko, and K. A. Komarova, J. Anal. Chem. (USSR), 22, 650 (1967).
1686
nitrate is a major contaminent. The Raman spectroscopic procedure reported by Brooker and Irish ( 2 ) is perhaps the best example of a method for nitrate whereby nitrite interference is minimized. However, it is doubtful whether this technique could yield the necessary accuracy at low nitrate levels or that costly Raman equipment and expertise could be adapted to routine laboratory environment. Methods in which nitrates are determined after decomposing the nitrites are generally inaccurate at low nitrate levels since additional amounts of nitrate are invariably produced ( I ) . Using a nitrate selective ion electrode, quantitative measurements in nitrite solutions are possible, provided the concentration of nitrite ion is not too large relative to nitrate [electrode response to nitrite is about 10% of nitrate response (3)]. Unfortunately, this condition does not exist in sodium nitrite salt where the nitrite,hitrate ratio is in the order of 200 to 500. Large amounts of nitrite effectively flatten the Nernstian response so that the electrode becomes relatively insensitive to nitrate. Therefore, it is necessary to eliminate or minimize the nitrite contribution if the electrode potential (2) M. H. Brooker and D. E. Irish, Can. J. Chem., 46, 229 (1968). (3) S. S. Potterton and W. D. Shults, Anal. Left., 1(2), 11 (1967).
ANALYTICAL CHEMISTRY, VOL. 42, NO. 14, DECEMBER 1970
measurements are to yield analytical results of maximum sensitivity and accuracy for nitrate. The reduction of nitrite to N20 with hydroxylamine, first described by V. Meyer (4), seemed ideal because of the speed and simplicity of this reaction. Seaman and coworkers (5) reported that some nitrate was formed during this reaction but they gave no quantitative data. We have found that, under the conditions of the nitrite-HAS reaction described, formation of only 0.02% nitrate (as NaN03) takes place through some disproportionation of nitrous acid. This nitrite-HAS reaction followed by selective ion electrode measurement forms the basis for the analytical procedure presented here.
for the standard. Finally, the sample solution potential was measured in the “REL MV” position after one minute while stirring at the same rate as the standard. The weight percentage of sodium nitrate was calculated from the following expression: NaNO, = milligrams NaN03/250 ml from curve grams of sample
Reagents. All reagents were commercially available reagent grade chemicals and were used without further purification. The reference electrode filling solution was prepared by dissolving 70 grams of ammonium sulfate and 0.50 gram of potassium chloride in 100 ml of water-then adding 0.1-0.2 gram of silver sulfate, stirring for 10 minutes, and filtering through Whatman No. 42 paper. This solution was used in preference to the Orion junction solution (No. 90-00-01) in order to eliminate contamination of the test solutions with nitrate which is a major component of the Orion solution. Apparatus. All potential measurements were read from a digital pH/mV meter, Orion model 801, Orion Research Inc., 11 Blackstone St., Cambridge, Mass. 02139. The nitrate indicator electrode was an Orion Model 92-07 and the reference electrode was a single junction silver-silver chloride electrode, Orion Model 90-01, filled with the reference solution described above. Calibration. A semilog calibration curve, mV us. milligrams NaN03/250 ml was prepared by pipetting 1.0, 1.5, 2.0, 3.0, 4.0, and 5.0 ml of standard N a N 0 3 solution (2.00 grams NaN03/liter) into 6 separate 250-ml volumetric flasks, each containing 40 ml of 1M HAS and 4 drops of 2 5 z HaPo?. The flasks, containing 2.0, 3.0, 4.0, 6.0, 8.0, and 10.0 milligrams of NaN03, respectively, were diluted to volume with water and mixed thoroughly. These solutions were found to be stable for at least 4 months. The potential of each standard was measured by transferring about 75 ml of solution, respectively, into 100-ml beakers and stirring (magnetic stirring bar) uniformly for one minute before recording the mV value. Sample Analysis. Sodium nitrite salt, 2.00 I 0.02 grams, was weighed into a 250-ml volumetric flask (pellet samples were pulverized with mortar and pestle before weighing). Forty milliliters of 1 M HAS was added rapidly from a graduate while, at the same time, washing down any sample crystals adhering to the neck or sides of the flask. When the reaction subsided (ca. 1 minute) 4 drops of 2 5 x H 3 P 0 4 were added while simultaneously swirling the flask. After standing for at least 7 or 8 minutes at which time evolution of N20 bubbles had virtually ceased, the contents of the flask were diluted to volume and mixed thoroughly. Before measuring the sample solution or a group of sample solutions, the potential of a calibration standard, containing nitrate at the expected concentration in the sample, was measured as before. The meter dial was then switched to the “REL MV” position and, by turning the Calibration control, the potential of the standard solution was adjusted to correspond exactly to the calibration curve millivolt value previously obtained (4) J. W. Mellor, “A Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. VIII, Longmans, London, 1931, p 289. (5) W. Seaman, A. R. Norton, W. J. Mader, and J. J. Hugonet, IND.ENG. CHEM., AWAL.ED., 14, 420 (1942).
0.1
- 0.02
(1)
RESULTS AND DISCUSSION
Nitrite-HAS Reaction. The reduction of sodium nitrite with HAS (4) is given as: 2 NaNOz
EXPERIMENTAL
x
+ (NH20H)z-HzS04+ 2 NzO + Na2S04
+ 4 HzO
(2)
When moderate strengths of reactants were combined, this reaction was found to be spontaneous at room temperature, rapid, and highly exothermic. However, to be analytically useful for subsequent nitrate measurements, it was important to establish reaction conditions whereby nitrite was nearquantitatively removed in the shortest possible time and the formation of nitrate ion through disproportionation of nitrite was negligible. A series of reactions with recrystallized sodium nitrite and varying concentrations of HAS were carried out using the well known Griess reagent (6) as a qualitative test for nitrite disappearance as a function of time, temperature, and concentration of reactants. These test conditions and results are given in Table I. At the end of the indicated time periods, the reactions were quenched with cold water, diluted to 250ml volume, and 20-ml portions withdrawn for the Griess Test. Two-gram portions of sodium nitrite were used in all tests as previous work (7) indicated that if smaller nitrite samples were taken for analysis, possible errors in the nitrate answers might be introduced because of poor representative sampling. HAS, rather than the more common hydrochloride salt, was used because sulfate interference to the nitrateselective electrode is only one tenth that of chloride interference (8). These test experiments were not intended as a quantitative kinetic study of this reaction but as a guide to establishing reaction conditions for the contemplated analytical procedure. The reaction was essentially completed within 30 seconds in each case, but beyond this time further reduction of the remaining nitrite proceeded very slowly as evidenced by the positive Griess tests. However, we observed that addition of acid immediately after the bulk of the reaction was completed caused the remaining nitrite to react quickly and completely, thus totally eliminating nitrite interference during the subsequent nitrate measurement (Table I, tests 7 through 10). By adding the acid after the bulk of the nitrite had already reacted, formation of nitrate through disproportionation of the nitrite was minimized. Phosphoric acid was selected primarily because the electrode response to phosphate anion is about 3000-fold less than nitrate response (8). Based upon the Griess test alone, the reaction conditions of test 5 appeared to satisfy the analytical requirements of near-complete (6) “Standard Methods for the Examination of Water and Waste Water,” 12th ed., Amer. Pub. Health Assoc., Boyd Albany, N. Y . , 1965, p 205. (7) Unpublished results, Eastern Laboratory, E. I. du Pont de Nemours & Co., Gibbstown, N. J. (8) “Instruction Manual,” Nitrate Ion Electrode, Orion Research Inc., 11 Blackstone St., Cambridge, Mass. 02139.
ANALYTICAL CHEMISTRY, VOL. 42, NO. 14, DECEMBER 1970
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Table 1. Test Reactions and Nitrate Measurements Following Reaction between 2.0 Grams of Recrystallized Sodium Nitrite and Hydroxylamine Sulfate (HAS) Drops Mole ratio Total Max react. Reaction of 25% % NaNOa Test HAS/NaN02 volume, ml temp, "C time, min H3PQ4 Griess test, color found 1 1.03 30 77 5 0 Red . . .a 10 s. pink 2 1.55 50 61 5 0 Red ... 8 0 s. Red ... 10 0 s. Pink 3 0.69 30 55 5 0 Red *.. 10 0 Pink ~ . . 4 0.69 20 76 10 0 Pink ... 5 0.86 25 66 5 0 s. Pink ... e
.
.
e . .
6 7 8
1.38 1.73 1.38 1.38 2.76 0.86 0.86 0.86 1.38 1.38 1.73 1.03
9 10 11 12 13
14 15 16 17
40
SO 45 40 40 25 50
75 65 90 50 30
15 10 10 15 8
53 52 52 53 55 66
5
8 15 40 25
0 0 2 2
4 4 4
... ... ... ...
40
4
52 77
15 8
4 4
4
4 4
...
Pink V.S. Pink V.S. Pink V.V.S. Pink Clear ,..
*.. ... ...
... *.. ,..
0.082
0.046
... ...
0.035 0.038 0.074 0.037 0.033 0.032 0.033 0.035 0.060
Indicates no test run.
reduction of nitrite in the shortest possible time using a minimum amount of HAS. Nitrate Formation. It still remained to establish if, during the hydroxylamine reaction, any nitrate ion was formed through some disproportionation of nitrite according to the following equations : NaNOz
3 "0%
+ H+
4
"03
+ Na" + 2 NO + H20
+-
HNOz
(3)
(4)
The possibility that side reaction 4 might be occurring was supported by a previous report by Seaman and coworkers (5) and from the fact that the pH of 1M HAS was about 3.0. Therefore, the nitrate concentrations in several of the test solutions (Table I) were measured by the nitrate-selective electrode procedure, employing cslibrating solutions containing sulfate concentrations corresponding to the final sulfate concentrations in the test solutions. Significant differences in percentage of nitrate found indicated to what extent side reaction 4 was taking place. Examination of the % N a N 0 3 found column in Table I clearly indicates the reaction conditions of tests 5, 11, and 17 gave more nitrate than the other tests. The relative violence, speed, and higher temperatures of these three test reactions (less water present) was felt to directly enhance reaction 4. Test 10 shows that, by using 2M HAS, the nitrite reduction is completed in about 5 minutes with essentially the same nitrate found as using 1M HAS (test 9) which required about 8 minutes reaction time. Unfortunately, the use of more concentrated HAS introduces more sulfate into the system which flattens the Nernstian slope during final electrode measurement and thereby reduces the precision of the potential readings. Therefore, test 9 represents the best combination of the desired analytical reaction conditions, namely, rapid nitrite reduction time, minimum concentration of HAS, and minimum formation of nitrate during the reaction between HAS and nitrite. 1688
In order to ascertain if the small percentage (0.035x) of sodium nitrate found in the recrystallized nitrite was either present in the sample or was in fact nitrate resulting from unavoidable disproportionation of nitrite, an attempt was made to prepare pure sodium nitrite completely free of nitrate as determined by the nitrate-selective electrode procedure. One hundred thirty grams of sodium nitrite, which initially contained 0.42 sodium nitrate, was recrystallized from water a total of four times. A two-gram portion of the recrystallized nitrite was analyzed for nitrate after each recrystallization. The nitrate content after the last three recrystallizations was about 0.02%, indicating that either 0.02% nitrate (as NaN03) was formed during the HAS reaction or, less likely, that 0.02 % nitrate could not be removed by further recrystallizations. Identical results were obtained by analysis of sodium nitrite samples which were repetitively recrystallized from water-methanol-dimethylformamide mixtures followed by washing the nitrite crystals with acetone. Additional positive evidence that reaction 4 was still occurring was obtained by testing the evolved gases at the height of the HASnitrite reaction with a Drager tube (Dragerwerk Lubeck, New Jersey Safety Equipment Co., 1700 Stuyvesant Ave., Union, N. J.), specific for NO-N02 mixtures. Finally, the unmistakable odor of NO2 was faintly detectable during the reaction although the distinctive brown NOLfumes could not be seen presumably because of the expulsion of oxygen by, and the overwhelming presence of, N 2 0(reaction 2). We concluded from these tests that 0 . 0 2 z nitrate (as NaN03) was consistently formed under the conditions specified in this procedure and therefore may be subtracted as a constant correction from the per cent nitrate found (Equation
z
1).
Additional experiments were performed in an effort tQ eliminate the 0.02 % nitrate formation. These included lowering the reaction temperature to
