Rapid Determination of Nitrate Nitrogen in the Presence of Ammonia and Urea R. M. ENGELBRECHT
and
F. A. McCOY
Research Department, Lion O i l Division, Monsanto Chemical Co.,
A method for the rapid determination of nitrate nitrogen is described. It is especially well suited to mixed fertilizers, as nitrate may be determined in the presence of ammonia, sodium, potassium, chloride, sulfate, phosphate, and urea. This method is an improvement over conventional methods in both analysis time and accuracy.
T
HE determination of nitrate nitrogen in the presence of
ammonia and/or urea is important, especially in the fertilizer industry. Existing methods are very time-consuming and desirable accuracy is rather difficult to obtain. The conventional AOAC method ( 1 )is a typical example. With only ammonia and nitrate present the sample is digested n i t h Devarda’s metal (an alloy of copper, aluminum, and zinc), which reduces the nitrate to ammonia. The ammonia is then distilled into an acid escess and the excess determined by a caustic titration. The nitrogen value thus obtained includes both the nitrate nitrogen and the ammonia nitrogen; it is obvious that the nitrate nitrogen value is dependent on the accuracy of the ammonia determination. The following calculation shows that any error in the ammonia analysis is magnified 4.43 times in the nitrate determination:
The analysis time is from 2 to 4 hours, depending on the method used to determine the ammonia nitrogen. If urea is present together with the ammonium and nitrate ions, another scheme of analysis is followed. The test solution is digested for 2 hours with the enzyme urease, which quantitatively converts the urea into ammonium carbonate in a neutral solution. After the digestion period the ammonia is distilled into an acid excess and determined in the usual manner. This gives the value for ammonium and urea nitrogen. To the residue in the distillation flask Devarda’s metal is added and the nitrate nitrogen is determined as described above. The accuracy of this method is not very good, because mechanical difficulties such as excessive foaming and bumping are encountered; furthermore, the accuracy of the nitrate determination is dependent on the completeness of the urea distillation. Another deterring factor in this analysis scheme is the 6-hour analysis time. A literature search revealed an oxidimetric procedure for nitrate determinations. This entails the reduction of the nitrate with an excess of ferrous sulfate solution and determination of the excess by a permanganate titration. There is nothing novel about this reaction, as Pelouze and Fresenius ( 5 )first performed the analysis in 1857. Kolthoff, Sandell, and Moskovits ( 2 ) determined the nitrate by a back-titration of the ferrous excess with potassium dichromate. However, these investigators used a time-consuming procedure, carrying the reaction out in a carbon dioxide atmosphere to avoid air oxidation. Leithe ( 4 ) gave a rapid method of analysis by the ferrouspermanganate and the ferrous-dichromate titrations. This reaction was carried out in a 60% sulfuric acid solution, which speeds the reaction time and aids in the elimination of air oxidation. The only data given in this work were for pure salts such
El
Dorado,
Ark.
as sodium nitrate and potassium nitrate. The work reported hrre was done to determine the accuracy of the method in the presence of ammonia, potassium, sodium, chloride, sulfate, phosphate, and urea. PREPARATIOh OF REAGE-ITS
Potassium Permanganate, 0.5S. Dissolve ahout 32 grams of reagent grade potassium permanganate in 2 liters of distilled water. Let stand for 2 days in a dark place; shake well several times dqring this period. After 2 days filter the permanganate through a glass nool-asbestos filter and store, well sealed, in a dark place. Standardize against sodium oxalate in the usual manner. Ferrous Sulfate, 1.ON. Weigh 270 grams of ferrous sulfate heptahydrate into a 1-liter volumetric flask. Dissolve in distilled water and a little sulfuric acid. When the ferrous sulfate has dissolved, add 40 grams of sodium chloride and dissolve. Add 100 ml. of concentrated sulfuric acid. Let cool and dilute to the mark with distilled water. EXPERIMENTAL PROCEDURE
Transfer 25.00-ml. aliquot of sample that contains no more than 0.30 gram of nitrate to a 250-ml. Erlenmeyer flask, and add 25.00 ml. of the ferrous sulfate solution. If a series of samples is to be run, stopper the flask m l l to prevent any possibility of air oxidation. Immediately prior to the determination, add 20.00 ml. of concentrated sulfuric acid, slowly and with gentle swirling.
Table I. Determination of Yitrate in Presence of Potassium, Ammonium, Chloride, Sulfate, and Phosphate NO8 Teighed, G.a
KoaTitrated, G.
0,3307 0,2502 0.3217 0.1470 0.1549 0.1671 0,2353 0,2450 0.2762 0,3487 0.3560 0.2253 0.1522 0,2272 0.3008 0.1104 0.1779 0.2545 0.3378
0.3314 0.2512 0,3230 0,1470 0.1544 0.1673 0,2353 0,2450 0.2760 0.3486 0.3648 0.2246 0.1514 0.2261 0 2998 0.1099 0.1770 0.2651 0.3385
Av. 0
Recovery, % 100.2 100.4 100.4 100.0 99.7 100.1 100.0 100.0 99.6 100.0 99.9 99.7 99.5 99.5 99.7 99.5 99.5 100.2 100.2 __ 100 0
+ 0.3
NaN03 or NHaKOa used as source of nitrate.
With nitrate present a dark-brown solution is obtained. Boil gently until the solution color becomes orange; it usually takes from 3 to 5 minutes of gentle boiling to reach the permanent orange color. As soon as the solution color has become orange, pour the contents into a flask containing about 1 liter of cold water. Rinse the reaction flask twice with distilled water and titrate the solution immediately with standard permanganate to a permanent pink end point. Make a blank for the ferrous sulfate solution in the same manner, but use 25.00 ml. of distilled water in place of the sample. Run the blank concurrently with a sample series.
1619
ANALYTICAL CHEMISTRY
1620 The per cent nitrate is calculated in the following manner
bubbling and bumping were encountered during the boiling stage. The results of an experiment to determine the effect of the ( 4 - B , ( K I I I: I O norniality)(6.2) ~ C b nitrate = sodium chloride addition are shown in Table 11. With the excep3 (sample weight; tion of the added amounts of urea and sodium chloride, these experiments were performed in the same manner as a blank titrawhere A is the volume of potassium permanganate used in the blank determination and B is the volume of permanganate retion. Because it as necessary to use different stock solutions quired for the sample titration. of ferrous sulfate and potassium permanganate which varied in strength, a factor had to be calculated to illustrate the effect of sodium chloride and urea on the titrations. Table I1 shows that This procedure is used for nitrate in the presence of potassium, when either of these compounds is present alone, the blank is sodium, chloride, sulfate, phosphate, and the ammonium ion. unaffected and, consequently, the calculated factor is 1.000. Five grams of sodium chloride are added prior to the addition However, the presence of both causes an increase in the amount of the ferrous sulfate when urea is present, to avoid a harmful of potassium permanganate required for the back-titration. effect of the urea on the subsequent decomposition of the nitrate. proportional to the amount of sodium chloride present. Except for the addition of phosphoric acid, the same method The following reactions are postulated in an effort to explain is employed when potassium dichromate is used for the backthis. titration. The apparent superiority of dichromate over permanganate is solution stability. However, indicator change (ferrous phenanthroline) for the dichromate titration was not 0 sharp enough. The permanent pink end point of the.permanganate is very readily detectable and the stability of the permanganate was found to be most satisfactory over n 2-week period. DISCUSSION OF RESULTS
Table I gives the results obtained by this method for solutions that contained potassium, sodium, chloride, sulfate, phosphate, and ammonium ions. The test solutions were prepared by weighing potassium OT sodium nitrate on an analytical balance in a 250-ml. volumetric flask, adding salts containing the other ions, dissolving these in 25.00 ml. of distilled water, and following the experimental procedure described above. The added salts were weighed on a torsion balance in quantities to simulate various types of mixed fertilizers. The accuracy of results is the same as given by Leithe ( 4 ) for pure nitrate salts-within f 0.3%. In addition to the accuracy, which is as good as, if not better than, the Devarda distillation, the and5 sis time is 15 minutes compared to 2 hours.
Table 11.
Effect of Sodium Chloride and Urea on Blank
Urea Added, G.
NaCl Added, G.
Correction Factor'
0.5
... ...
1.000 1,000 1,000 1,000 1.010
1.0
... ...
5.0 10.0
0.5 0.5 0.5
10.0
1.0 1.0 1.0
1.5 1.5 1.5 (I
Correction actor =
5.0
1,020
15.0
1.029
5.0 10.0 15.0
1.011 1.022 1.032
5.0 10.0
1.013 1,025 1.033
15.0
volume KMnO4 (urea and NaCl present) volume KMnO4 (no urea and NaCl present
.
When urea is present, 5 grams of sodium chloride are added to the sample aliquot before addition of the ferrous sulfate solution. Results obtained for nitrate determinations with urea present and with no sodium chloride added were nonuniformly low. The addition of 5 grams of sodium chloride brought the nitrate value to within 2% of the theoretical value. Ten and 15 grams of sodium chloride were also added. However, with this concentration of salt, which approaches the saturation point, excessive
0
When only urea is present, a hydrolysis takes place similar to Reaction 1. With only sodium chloride present, a reaction similar to Equation 2 takes place and the hydrogen chloride is evolved during the boiling stage. With both urea and sodium chloride present, Reaction 3 is given as a possible reaction to explain the increase in the amount of permanganate needed for the back-titration. Urea nitrate and urea-hydrogen chloride are known to form under reaction conditions very similar to those of this method ( 6 ) . If the compound urea-hydrogen chloride is formed, it is reasonable to expect this increase in the permanganate titration, rn the ferrous iron-permanganate reaction is known to induce the hydrogen chloride-permanganate reaction (3). With no sodium chloride added the urea probably prevents the complete decomposition of the nitrate by forming urea nitrate; this would explain the erratic low results obtained in earlier work. The addition of a large excess of sodium chloride gives a copious quantity of hydrogen chloride, which unites with the urea to form urea-hydrogen chloride, thus permitting the complete decomposition of the nitrate. With the 10- and 15-gram addition of sodium chloride the solubility limit is reached; furthermore, excessive bumping and bubbling were encountered and with this a probable sample loss. Therefore, when urea is present in a test sample, 5 grams of sodium chloride are added to the sample aliquot prior to the addition of the ferrous sulfate. Table I11 gives the results of nitrate nitrogen determinations in the presence of urea. The urea content was varied to give a nitrate-urea ratio from 1:4 to 3 : l . Five grams of sodium chloride were added to each sample. However, it was found expedient to perform the blank determination with no sodium chloride or urea present. This blank was then used to determine nitrate in either the presence or absence of urea. The results in Table I11 show a recovery of 9i.9 Z!C 0.7%. The uniformly low results
V O L U M E 28, NO. 10, O C T O B E R 1 9 5 6
1621
f!:) ( nitrate result brings the nitrate value to 100.0 f 1.0%.
97.9%, a factor of 1.022 which is - multiplied by the
Table 111. Determination of Nitrate in Presence of Potassium, Ammonium, Chloride, Sulfate, Phosphate, and Urea Added, Urea G. u p to u p to u p to u p to u p to u p to u p to u p to u p to u p to u p to u p to
Added, NaCl G.
Weighed, Nos- G.
Titrated, NO8 - G ,
Recovery, %
5.0 5.0 5 0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5,O 5.0 5,O 5.0 5,o 5.0 5.0
0.1741 0.1756 0.2590 0.1714 0.2619 0.1623 0,1539 0.2345 0.2903 0.3286 0.3132 0.2903 0.2953 0,2587 0.2911 0.2250 0.1397 0.3049 0.1943
0,1686 0.1701 0.2518 0.1683 0,2530 0.1591 0.1524 0.2305 0.2834 0 3232 0,3086 0.2874 0.2895 0.2510 0.2866 0.2217 0.1339 0.3019 0.1895
96.8 96.9 97.2 98.2 96.6 98.0 99.0 98.3 97.6 98.4 98.5 99.0 98.0 97.0 98.0 98.5 96.5 99.0 97.5
1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 0.25 0.50 1 .o 0.75 0.20 0.10 0.40
CONCLUSION
.
97.9 -I- 0 . 7
AY.
were expected, as neither sodium chloride nor urea was present in the blank. To compensate for this the following calculation is used for nitrate in the presence of urea
yo nitrate
=
Thie has been applied routinely and found most satisfactory. The use of a correction factor may be avoided if a blank is determined to which about 1 gram of urea and 5 grams of sodium chloride have been added. In either case the accuracy of this method is better than that of the urease-Devarda method. The analysis time is again 15 minutes compared to 6 hours.
( A - B)(KMnOn normality)(6.2)( 1.022) 3 (sample weight)
where A is again the volume of permanganate used for the blank determination and B is the permanganate volume required for the sample titration. As the recovery of nitrate with 5 grams of sodium chloride added to the test sample was found to be
The oxidimetric determination of nitrate is an improvement over the conventional methods from the standpoint of time and is equally good, if not better, with regard to accuracy. The method is especially well suited to nitrate nitrogen determinations in mixed fertilizers. However, it may be used to determine nitrate in any mixture free of substances that will react with the ferrous sulfate and potassium permanganate. In this study very good results were obtained when the method was applied to pure lithium, lanthanum, and uranium salts. LITERATURE CITED
(1) Assoc. Offic. Agr. Chemists, “Methods of Analysis,” 7th ed., p. 14, 1950. (2) Kolthoff, I. II.,Sandell, E. B., Xloskovits, B., J . Am. Chem. SOC. 55, 1454 (1933). (3) Kolthoff, I. M., Stenger, V. A., “Volumetric Analysis,” vol. I, pp. 168-77, Interscience, New York, 1942. (4) Leithe, Wolfgang, ANAL.CHEM.20, 1082 (1948). ( 5 ) Pelouze and Fresenius, Ann. 106, 217 (1857). (6) Soci6t6 d’Etudes Chimiques pour l’Industrie, Brit. Patent 179,544 (July 12, 1923). RECEIVLD for review Aprll20, 1956. Accepted May 14, 1956. 11th Southwest Regional Meeting, ACS, Houston, Tex.. December 1 to 3, 1955
Determination of Free and Combined Formaldehyde Using Modified Chromotropic Acid Procedure Application to Determination of Piperine Content of Pepper LEONARD A. LEE Eastern Utilization Research Branch,
USDA, f h i / a d e / p h i a 78,
A modified colorimetric chromotropic acid procedure has been developed for the quantitative determination of formaldehyde liberated from the methylenedioxyl group in piperine, the principal substance associated with the sharp taste of pepper. The method thus provides a simple, rapid procedure for the analysis of the piperine content of pepper. The conditions selected are also advantageous for the determination of free or combined formaldehyde in other compounds.
ERY little work has been published on the specific determination of piperine, which probably is the major pungent component of pepper (Piper nigrum). Use has been made of such general methods as Kjeldahl nitrogen estimation (15)or the weighing of a partially purified residue following extraction of the ground spice (6). Fagen, Kolen, and Hussong ( 8 ) have recently reported an ultraviolet spectrophotometric method for
Pa.
the determination of piperine in oleoresin of pepper. Their method is both specific and sensitive, but cannot be used by laboratories that have only visual range photometers. The desire by the spice trade and other segments of the food industry for a simple, rapid method for the determination of piperine led to an investigation of the possibility of developing such a method. Piperine forms colored complexes with a number of alkaloidal reagents ( 8 ) and colored salts with some concentrated acids (3,I S ) . None of these, however, was found useful for a quantitative colorimetric procedure; reactions were qualitative, or colors were unstable, or absorption maxima were located in the ultraviolet region, just beyond the range of several commonly used photoelectric colorimeters. Attempts to take advantage of the weakly basic properties of piperine (which is the piperidide of the unsaturated piperic acid) were not successful. The amido linkage could not be titrated, even in nonaqueous systems. However, piperic acid contains a methylenedioxyl group, which should yield formaldehyde on acid hydrolysis. Bricker and Vail
