Volumetric Determination of Nitrate Ion Simultaneous Determination of Nitrite OTTO RICHARD GOTTLIEB and MAURO TAVEIRA MAGALHAES lnstituto de Qdrnica Agrkola, Minisi6rio da Agricultura, Rio de Janeiro, Brazil ,The reaction between sulfamic acid and nitric acid is used in the determination of nitrate ion by titration of the excess sulfamic acid with standard sodium nitrite solution, using an external indicator such as starch-iodide to detect the end point. Quantities of potassium nitrate between 5 and 100 mg. were determined with a deviation from the correct value oscillating between &0.6% and an average The method error of less than 0.1%. is sufficiently selective. Nitrites and ammonium salts do not interfere. The procedure does not call for any special apparatus, and it is rapid and well suited for routine work.
A
Iyas needed for the determination of nitrate nitrogen which did not rely on special equipment, reagents, or techniques. Berglund (3) discovered that sulfamic acid is decomposed with effervescence, even a t room temperature, by nitric acid or by a mixture of sulfuric acid and a nitrate. He identified the gas as nitrous oxide. Baumgarten '(2)found that at room temperature the reaction occurs only with concentrated nitric acid, proceeding quantitatively according to the equation: METHOD
HNO'
+ HzN*SOsH
+
NzO
+ HzsOt
+ Hz0
Previous work on gasometric titration (8) showed that this reaction can be used for the determination of nitrate ion. The acids are freed from a mixture of nitrate and a known quantity of sulfamate, by adding concentrated sulfuric acid a t room temperature. After this rapid reaction is completed, the excess sulfamic acid is titrated with a standard solution of sodium nitrite. REAGENTS
SODIUM NITRITE. Crystallize reagent grade material three times from water. Dry the crystals to constant weight in a desiccator under vacuum. Dissolve a 5.5206-gram sample in distilled water and dilute t o 1 liter. This 0.08M solution can be used as a primary standard (9) for 90 days (6). SULFAMIC ACID. Dissolve about 7.77 grams of sulfamic acid (reagent grade,
G. Frederick Smith Chemical Co., Columbus, Ohio) in distilled water. Dilute to 1 liter, to prepare an approximately 0.08M solution. Starch, 3% aqueous solution. Potassium iodide, free from iodine, 20% aqueous solution. PROCEDURE
Prepare a neutral sample solution containing 0.4 to 0.6 gram of nitrate ion per 100 ml. Preliminary Test. If one drop of the sample solution after acidification with dilute sulfuric acid, 10% by volume, gives a feeble or negative starchiodide test (absence of nitrite), use 25.00-ml. portions of 0.08iM sulfamic acid solution in the procedure below. If a strong blue color appears in the test, treat 2 ml. of 0.086f sulfamic acid with 1 ml. of 10% sulfuric acid and titrate the solution rapidly with the sample solution to an approximate starch-iodide end point. If 4 ml. of sample solution have been used up and the end point has not been reached (presence of only small quantities of nitrite), follow the procedure as described with 25.00-ml. portions of 0.08M sulfamic acid solution. If, however, between 4 and 0.5 ml. of sample solution are required to reach the end point (presence of significant quantities of nitrite), increase the portions of 0.08M sulfamic acid solution in the procedure. It is best to use portions in the vicinity of (25.00 20.00:~)ml., z representing the volume of sample solution consumed in the titration above. If less than 0.5 ml. is sufficient to reach the end point (presence of very great quantities of nitrite), dilute the sample solution until a t least 0.5 ml. is needed to react with 2 ml. of the sulfamic acid solution. Determination of Nitrate. Transfer t o each of two 250-ml. beakers marked I and 11, equal volumes, 25.00 ml. or more (see Preliminary Test) of 0.08M sulfamic acid solution. Add t o each 0.5 ml. of 10% sulfuric acid. Run into each slowly 10.00 ml. of sample solution with stirring. Continue t o stir for 3 minutes. Add one drop of 0.1% phenolphthalein indicator and then dilute sodium hydroxide solution to slight alkalinity. Evaporate the contents of both beakers t o dryness on the steam bath. Heating should be discontinued a t the moment the residue completely solidifies; if heating is prolonged beyond this point, it will be less easily attacked by
+
Table 1. Effect of Procedure Conditions on Stability of Sulfamate
0.08M
Sulfamic Bcid solution Taken, Ml., 25,00 25.00 25.00
0.08M Sodium Nitrite
Solution, Used M1. 1.
2b
25.05 25.03 25.05
25.05 25.04 25.03
The sulfamic acid solution, after addition of sulfuric acid, was titrated with sodium nitrite solution. * The sulfamic acid solution mas subjected to the conditions of the procedure and then titrated with sodium nitrite solution. 0
sulfuric acid. Cool the residue to room temperature. To beaker I, add 2 ml. of concentrated sulfuric acid and make the mixture homogeneous. After the gas ceases to evolve (a few minutes), dilute n-ith distilled water to approximately SO ml. To beaker 11, add 80 ml. of distilled water and then 2 ml. of concentrated sulfuric acid. Titrate the excess sulfamic acid in both beakers with standard 0.OSM sodium nitrite solution, using a starchiodide solution as an external indicator. Stir the contents of the beaker constantly. Add the nitrite solution slowly (about 2 to 3 ml. per minute), stopping for 1 minute after each 5-ml. addition. To verify the end point of the titration, perform two spot tests, the first 1minute and the second 3 minutes after adding each drop of nitrite solution. The equivalence point is considered reached, when the indicator shows a slight blue color within 15 seconds in both tests. Perform these spot tests by adding a t the moment of their execution, in the cavity of a plate, one drop of starch solution and one drop of potassium iodide solution and touching this misture with the tip of a glass rod which has been immersed in the titrated solution. Determination of Nitrate and Nitrite. To beaker 111, add the same volume of 0.08M sulfamic acid solution used in beakers I and I1 and proceed as outlined for beaker 11, omitting the sample solution. CALCULATION
The per cent nitrate (or nitrate nitroVOL. 30, NO. 5, MAY 1958
* 995
gen) and nitrite (or nitrite nitrogen) in the sample may be calculated, respectively, through the expressions:
P (a - a) M -
lo.p
and
Q (c - b ) M -
lo.p
where P is the molecular weight of the nitrate (or the atomic weight of nitrogen), Q is the molecular weight of the nitrite (or the atomic weight of nitrogen), p is the weight in grams of the sample in the aliquot taken for analysis, 11f is the molarity of the standard nitrite solution, and a, b, and c are the volumes in milliliters of standard nitrite solution consumed, respectively, in the titration of the content of beakers I, 11, and 111. DISCUSSION
Because large amounts of nitrite may use up significant quantities of sulfamic acid, it is necessary to make a preliminary check for nitrite and, if this is present, to add a sufficient excess of
Table
II.
Determination of Pure Potassium Nitrate
Mg. Found
Mg.
Present 80.92"**
a
80.60 80.75 80.87
Error,
7%
-0.40 -0.21 -0.06 -0.06 -0.02
80.87 80.90 80.92 0.00 80.95 +0.04 80.98 +0.07 80.98 $0.07 81.03 $0.14 81.11 $0.23 81.17 $0.31 81.17 +0.31 81.42 $0.62 Av. 80.98 i0.15 f 0 . 0 7 8.09"*c 8.02 -0.87 8.10 +o. 12 8.10 $0.12 8.15 +0.74 Av. 8.09 =t0.04 $0.03 Quantity contained in 10.00 ml. of
eolution.
* 25.00 ml. of 0.08M sulfamic acid solu-
tion added. Titration performed with
0.08M sodium nitrite solution. c 20.00 ml. of 0.016M sulfamic acid
solution added. Titration performed with 0.016M sodium nitrite solution.
Table 111.
Samde
sulfamic acid to compensate for the reagent used up by the nitrite ion. If the sample does not contain nitrite ion, the order of addition of sulfamic acid solution and sample is irrelevant, and treatment of the sulfamic acid solution by sulfuric acid and slow addition of sample solution are also unnecessary. Only very concentrated aqueous solutions of nitric acid react with sulfamic acid ( 2 ); hence reaction betxeen the two substances is prevented by dilution. By performing tIvo parallel operationsLe., reacting nitrate with sulfamic acid in beaker I and preventing the reaction in beaker 11-problems pertinent to the stability of the sulfamic acid solution are excluded. The data of Table I show that the sulfamic ion is not appreciably altered under the conditions of the analysis. Quantities of concentrated sulfuric acid larger than 2 ml. make the starch-iodide end point rather difficult to observe. Quantities smaller than 0.5 ml. make a thorough suspension of the evaporation residue difficult and decrease the velocity of the reaction between the nitrite and sulfamate ions. The optimum quantity of sulfamic acid to be titrated with the 0.08V nitrite solution was established from data given by Bowler and Arnold (4). External indicators, such as starch-iodide
Table IV. Determination of Nitrate Nitrogen in Commercial Products
Nitrate Nitrogen
Found, yo AOAC
method (1)
Fertilizer"
Proposed procedure
4.6
Calcium nitrateb
4.67 4.70 4.70 14.14 14.11 14.13
14.1
4 Organic nitrogen, 1.7y0 as N; chloride, 4,5% as C1; phosphorus soluble in 2% citnc acid, 11.7y0 as PzO5; total phosphorus, 12.2y0 as Pp06; potassium soluble in water, 5.1% aa KIO. * Product, which contained 1.6% ammonium nitrogen, was labeled 15.5% N.
Simultaneous Determination of Nitrate and Nitrite pvlg. Found hlg. Present
KNO, 80.92
NaKOS 11.07
KNOI
NaNOl
10.81 80.95 10.98 80.78 11.18 80.71 55.53 80.71 55.36 80.92 2 55.36 80.87 55.24 81.03 3 32.37 166.08 32.20 166.63 32.45 166.30 33.01 165.91 In samples 1, 2, and 3, respectively, 25, 35, and 50 ml. of 0.08M sulfamic acid solution were used, Titration was performed with 0.08M sodium nitrite solution. 1
996
ANALYTICAL CHEMISTRY
(4) or even the more sensitive sulfanilamide-1-naphthylamine (Y), are not satisfactory \\-hen solutions of nitrite and sulfamate ions more dilute than 0.0lM are titrated. This establishes a lower limit of about 5 mg. for the quantity of nitrate ion which can be determined by the procedure described. Preliminary experiments showed the possibility of performing the titration by electrometry. The amperometric titration of sulfamates with nitrite, as described by Hirozawa and Brasted (@, in combination with the method here outlined, might reduce the lower limit of nitrate ion determination. Slow reaction near the equivalence point is responsible for the wait before proceeding with the verification and the confirmation of the end point. RESULTS
Some results of the application of this method to an analytical sample of potassium nitrate are presented in Table 11. Table I11 illustrates results which can be obtained in the determination of mixtures of nitrite and nitrate. INTERFERENCES
Chloride ion causes loss of nitrate when the residue is treated with sulfuric acid. This interference can be eliminated by transforming the chloride ion to silver chloride with a slight excess of silver sulfate. The determination of nitrate in a fertilizer (Table IV) was performed in this manner. Although nitrite ion reacts with sulfamic acid, the double titration procedure eliminates its influence on the nitrate determination. By an additional titration, using the same solutions and technique as in the nitrate determination, nitrite can be determined simultaneously. The interference of other substances capable of reacting with sulfamic acid in aqueous solution is also eliminated by the double titration. However, substances which, like nitrate ion, react with sulfamic acid only when treated with concentrated sulfuric acid cannot be present. Phosphates, sulfates, carbonates, and many cations, including ammonium, do not interfere. Cobalt(I1) ion is transformed by nitrite to the cobalt(II1) complex, [Co(NO&]---, which is stable toward sulfamic acid (6). Hence cobalt(I1) interferes with the nitrite determination. The introduction of salts into the aliquot, as through some pretreatment of the sample, must be practiced with caution. Relatively large quantities of solids in the mivture with the nitrate make a thorough suspension of the residue in concentrated sulfuric acid difficult. This condition may cause low results. Organic substances which are easily oxidized under the conditions of the
method can be destroyed with peroxide (IO), which, in turn, must be eliminated before nitrate can be determined. Organic substances which are reasonably stable under the conditions of the method do not interfere considerably, even if they are nitrogenated. Preliminary experiments showed that satisfactory, although somewhat low, results can be obtained for nitrate in the presence of equal quantities of gelatin or urea. ACKNOWLEDGMENT
The authors wish to express their ap-
preciation to Kalter B. Mors for his interest in this work and t o Mary Nebel, Wayne State University, Detroit, hIich., for revision of the English manuscript. LITERATURE CITED
(1) Assoc. Offic. Agr. Chemists, “Official hlethods of Analysis,” 8th ed., p. 13, 1955. (2) Baumgarten, P., Bey. deut. chem. Ges. 71, 80 (1938). (3) Berglund, E., Bull. S O C . chim. [2], 29, 422 (1878). (4) Bowler, W. W., Arnold, E. A., ANAL.CHEN.19, 336 (1943).
(5) Brasted, R. C., Zbid., 23, 980 (1951). ( 6 ) Zbid., 24, 1111 (1952). (7) Cumming, W.M., Alexander, W. A., Analust 68, 273 (1943). (8) Gottlieb, 0. R., “Titrimetria Gasombtrica,” Bol. 42, p. 47, Tnstituto
de Quimica Agricola, Ministbrio da Agricultura, Rio de Janeiro (1955);
Anal. Chim. Acta 13, 531 (1955). (9) Hirozawa, S. T., Brasted, R. C., ANAL.CHEM.25. 221 (19531. (10) Johnson, C . M., blrich, A,; Zbid., 22, 1526 (1950).
RECEIVEDfor review May 9, 1957. Accepted December 5, 1957. Work supported by Conselho Nacional de Pesquisas, Brazil.
Basic Behavior of Molecules and Ions in Acetjc Anhydride C. A. STREULI Analytical Research laborafory, American Cyanamid ,Titration of neutral and anionic bases in acetic anhydride has shown that a linear relationship exists between pK,(H20) values and half neutralization potentials in acetic anhydride. The relationship is, however, different for the neutral bases than for ions. Anions are stronger bases in acetic anhydride than in water relative to the neutral compounds. It is also possible to titrate halide salts directly in the solvent. Chlorides and iodides may be resolved, but this is not possible for chlorides and bromides or bromides and iodides. Anions which form weak acids in water solution are leveled in basic strength in acetic anhydride.
T
utility of acetic anhydride as a titration medium, either alone or in conjunction with another solvent, has been noted by Gremillion (4), Fritz (S), and Usanovich ( I S ) . A number of compounds which do not exhibit basic properties in water, acetic acid, or acetonitrile may be quantitatively titrated in the anhydride. The principal drawback of this solvent is its reactivity with some solutes to produce new products, usually weaker bases. The work described was initiated to determine if weak anionic bases might be titrated in this solvent, to give a more detailed picture of acid-base behavior in the anhydride, and to establish, if possible, a relationship between pK,(H20) and half neutralization potential (HNP) of weak bases in acetic anhydride. The substances used as standards were chosen to give as large a range of pK, values as possible, but were limited to those substances which showed a minimum chance of being acetylated by the solvent. In the case of the amines, HE
Co., Stamford, Conn.
this restricted the choice to tertiary amines and heterocyclics. Two assumptions must be made in viewing qualitatively acid-base behavior in acetic anhydride: Hydrogen acids behave (and bases react with hydrogen acids) in much the same manner as they do in a protonated solvent, such as acetic acid; and liquid junction potentials, while unknown, are constant and do not distort the titration curves.
PROCEDURE
Weigh out 1 mmole of the sample to be tested, dissolve it in acetic anhydride, and dilute to 100.0 ml. in a volumetric flask. Pipet 25.0 nil. of this solution into a EO-ml. beaker, add 75 ml. of anhydride, and titrRte with
REAGENTS AND APPARATUS
Reagents. The acetic anhydride used as solvent was Baker and Adamson reagent made. which was not further purifiedu. ‘ PERCHLORIC ACIDSOLUTION, 0.05N. This was prepared by diluting 4.2 ml. of 72% acid with about 500 ml. of acetic acid and then diluting this to 1 liter with acetic anhydride. The solution became highly cofored on standing, but color formation did not affect the normality. The acid was standardized against potassium hydrogen phthalate (11). Xost of the organic compounds mere Eastman Kodak White Label. Quaternary ammonium salts were obtained from the southwestern Chemical Co. Other salts were reagent grade. Apparatus. All titrations were performed using a Precision-Dom Recordomatic titrator. Solutions were agitated with a magnetic stirrer during the titration. The electrodes mere a conventional glass electrode and a silver-silver chloride reference cell. The latter was made from the shell of a Leeds & Northrup calomel electrode, and contained a coiled silver wire, electrically coated with silver chloride, immersed in acetic anhydride which had been saturated with silver and lithium chlorides.
‘ 0
ll-ullJ P 40 80 80 100 120 % Neutralization
Figure 1 . Titration of neutral nitrogen bases in acetic anhydride a. Acetanilide b. Acetamide c. Caffeine d. Urea e. Methylurea f. N,N-Dimethylaniline g. Quinoline h. Pyridine 1. N,N-Diethylanlline k. N,N-Dimethylbenzylamine I. Tri-n-butylamine
VOL. 30, NO. 5, MAY 1958
997
