Benzyl and Allyl Halides. Benzyl and allyl halides give an interfering color which obscures any possible test. This is in agreement with Kornblum’s investigations in which he found t h a t only at low temperatures were the nitro compounds formed. Dihalides. No test was obtained with dichloromethane, ethylene dihalides, or 1,3-dichloropropane. The only other commercially available higher dihalide, 1 ,bdiiodopentane, did give a weak but definite test (0.15 mmole of nitro compound obtained, 5% yield). Difunctional Compounds. A positive test was obtained with only one halide derivative containing a second functional group, 3-chloro-1-propanol. Either no color o r interfering colors were obtained with the following: 2-Chloroethanol 2,3-Dichloro-3-propaol 1-Chloro-%propanone Chloroacetic anhydride Chloroacetic acid 2-Chloroethyl acetate Ethyl 3-chloropropionate 2-Chloroacetamide 4Chlorobutyronitrile Epichlorohydrin Bis(Zchloroethy1) ether Procedure Using Dimethyl Sulfoxide. Sodium nitrite is more soluble in dimethyl sulfoxide than in dimethyl-
formamide and Kornblum has recommended (8) this solvent for formation of the nitroalkanes. Essentially, the procedure previously described was used except that a 10% solution of sodium nitrite in dimethyl sulfoxide was substituted for the dimethylformamide reagent. Positive tests were obtained with all primary chloroalkanes through l-chlorodecane. However, most of the red color develops on adding alkali to the original reaction mixture. Dropwise addition of sulfuric acid gives only a slight intensification of the color. The red color disappears in acid solution and is regenerated in basic media. Secondary chloroalkanes generally give a deeper color using the dimethyl sulfoxide reagent in comparison to the dimethylformamide. An exception is chlorocyclopentane which gives only a faint test using the sulfoxide solution. 2-Chloro-Zmethylpropane gives a negative test using the sulfoxide procedure. %Bromo-Zmethylpropane and 2-iodo-%methylpropane give a yellowish tan color when the sulfuric acid addition is made. This conflicts with the test for a primary halide. CONCLUSIONS
This test should be applied only to
those compouhds which give a positive test with alcoholic silver nitrate. With such compounds proper classification will be obtained with simple alkyl and cycloalkyl halides within the limits specified. ACKNOWLEDGMENT
Frank V. Kitko is indebted to the National Science Foundation for support with a Summer Fellowship for Secondary School Teachers during part of this investigation LITERATURE CITED
(1) Kornblum, N., Larson, H. O., Black-
wood, R. K., Mooberry, D. D., Oliveto, E. P.. Graham, G. E., J . Am. Chem. SOC. 78,1497 (19%). (2) Kornblum. Powers.’ J. W.. J. Ora . .Chem. 22,455N.. (1957).
(3) Kornblum, N., Taub, B., Ungnade, H. E., J . A m . Chem. SOC.7 6 , 3209 (1954). ( 4 ) MeXer, V., Locher, J., Ber. 7, 1510 (1874). RECEIVEDfor review April 13, 1960. Accepted July 14, 1960. Division of Analytical Chemistry, 136th Meeting, ACS, Atlantic City, N. J., September 1959. Taken from a portion of a thesis to be submitted by Frank V. Kitko to the Graduate School o f Western Reserve University in partial fulfillment of the requirements for the degree of doctor of philosophy.
Differential Colorimetric Determination of Nitrite and Nitrate Ions JACK L. LAMBERT and FRED ZITOMER Department o f Chemisfry, Kansas State University, Manhattan, Kan.
b Nitrite and nitrate ion in the ranges of 0 to 1.5 and 0 to 2.0 p.p.m., respectively, are determined colorimetrically by the formation of a red dye through the diazotization-coupling reaction. (4 Aminopheny1)trimethylammonium ion i s used as the diazotizable amine and N,N-dimethyl-1 -naphthylamine is used as the coupling agent. Nitrite ion is determined directly in acid solution. Nitrate ion is first reduced by zinc in ammoniacal solution of pH 10.8 in the presence of manganous hydroxide, and then determined by the same reaction as the nitrite ion. Nitrite ion i s eliminated in the nitrate determination by exchanging the diazonium cation formed onto a nuclear sulfonic ion exchange resin (potassium form) prior to the reduction of nitrate ion. Removal of residual chlorine, which interferes in the nitrate determination, is accomplished through reaction with the same ion exchange resin, presumably through
-
1684
ANALYTlCAL CHEMISTRY
chlorination of the resin. Residual chlorine and nitrite ion are not found together. Chloride, sulfate, and other commonly occurring substances do not interfere.
N
ION can be determined colorimetrically by its well known selective diazotization-coupling reaction to form an azo dye (1, 3). Bray and coworkers (6) used this reaction for the determination of nitrate ion in soils after reduction t o nitrite ion by zinc in acid solution in the presence of manganous ion. Middleton (a has shown that the reduction is quantitative only in ammoniacal solution of pH 10.2 to 11.2 and a t lower than room temperatures. Methods used most often for the determination of nitrate ion in water are the phenoldisulfonic acid (1) and thc brucine procdures (2). The phrnoldisulfonic acid method is time consuming and subject to intrrferencr by chloride ion. The brucine method an ITRITE
be used to determine nitrate ion if nitrit ion is absent; otherwise, nitrate is determined by difference. Both procedures produce yellow solutions which are d a c u l t to compare visually. The method described here can be used to determine nitrate and nitrite ions separately in the presence of each other. Nitrite ion is determined by the diazotization of (4aminophenyl)trimethylammonium ion and subsequent coupling with N,N-dimethyl-1-naphthylamine. A bright red dye is produced which is excellent for visual or spectrophotometric purposes. Nitrate ion is determined by the same reaction following quantitative reduction to nit.rite ion. When nitrate ion is determined separately in the presence of nitrite ion, the latter is first removed by diazotization of the (4-aminoplienyl)trimt.tlryl:tmmonium ion and ewhange of tlit, rcsulting tlipositive cation onto a nuclear SUIEonic ion cschnngr resin. WIP rr:.in 3150 servrs t c j rexnovc rc,>itlual
to the nitrate ion determination. Chloride, sulfate, and other commonly occurring substances do not interfere. SPECIAL REAGENTS
All reagents are of the best available grade. Standard nitrite ion solution, 10oO p.p.m., 1.500 grams of sodium nitrite per hter of solution. Standard nitrate ion solution, 10oO p.p.m., 1.631 grams of potassium nitrate per liter of solution. Amine solution, 1% (4aminophenyl)trimethylammonium chloride monohydrochloride. This solution should be stored in a dark bottle. Coupler solution, 0.5% N,N-dimethyl-1-naphthylamine in 1% hydrochloric acid. This solution should be stored in a dark bottle. Ammonia solution, 120.0 ml. of concentrated ammonium hydroxide per liter of solution. Manganese solution, 1% manganous chloride tetrahydrate. Preparation of (4-Aminopheny1)trimethylammonium Ion. T o 43.0 grams of N,N-dimethyl-p-phenylenediamine monohydrochloride dissolved in 650.0 ml. of water a t 50' C., add 30.0 ml. of acetic anhydride and stir. Add a solution of 45.0 grams of anhydrous sodium acetate in 150.0 ml. of water. Cool to 20' C. in an icewater bath. Add a several hundred per cent excess of concentrated ammonium hydroxide to complete precipitation of the product. Collect the precipitate formed l5y filtration. Wash the precipitate three times on the filter by working it into a slurry with equal portions of a solution containing 75.0 ml. of concentrated ammonium hydroxide and 150.0 ml. of water, and air dry (about 1 to 2 days). Dissolve the precipitate in lo00 ml. of warm benzene. Filter and add 1000 ml. of benzene a t room temperature. Pour the solution into a 2000-ml. standard taper flask and add 32.0 ml. of methyl iodide. so that minimum air space remains. Mix thoroughly and stopper tightly without stopcock grease. Store in a dark cool place for 4 days. Filter, and recrystallize the product from methanol. &dissolve the product in a minimum amount of hot methanol. For each gram of product in the preceding step, dissolve 0.7 gram of lead acetate triliydrate in hot methanol. Mix the two alcoholic solutions, ,cool, and filter. Bubble hydrogen sulfide gas into the filtrate until precipitation of lead sulfide is complete, and filter, Add 0.6 gram of concentrated hydrochloric acid to the filtrate for each gram of the product collected in the recrystallization from methanol above. Concentrate to about 100 to 200 ml. (or to where crystals form) with forced evaporation on a steam bath. Cool and add just enough methanol to dissolve all the crystals Triple the total volume by dilutior. with ether. After 1/2 hour 07 more, filter the product. Dissolve in a minimum quantity of methanol and repeat the ether dilution. Collect thr.
final product by filtration and dry in air. If the product is not colorless, dissolve it in methanol and decolorize with carbon. Precipitate with ether as previously mentioned and dry. Yield, approximately 17 grams. Preparation of Cation Exchange Resin. Place the desired amount of Amberlite IR-lSO(H) nuclear sulfonic cation exchange resin (hydrogen form) on a filter. Wash repeatedly with 20% potassium hydroxide until the wash solution remains blue t o litmus. Rinse with distilled water until the wash solution is neutral to litmus. Dry in air until the resin beads no longer stick together. EXPERIMENTAL
Nitrite Calibration Curve. Prepare nitrite ion samples in the range of 0.1 t o 1.5 p.p.m. To 30.0 ml. of the samples add 1.00 ml. of the amine solution and stir well. Add 1.00 ml. of the coupler solution and acidify with 5 or 6 drops of concentrated hydrochloric acid. Dilute to exactly 50.0 ml. After 20 minutes, record absorbance readings with a spectrophotometer a t 522 mp. Nitrate Calibration Curve. Prepare nitrate ion samples in the range of 0.1 t o 2.0 p.pm. To 30.0 ml. of the samples in a 125-ml. Erlenmeyer flask add 5.00 ml. of the ammonia reagent and 1.00 ml. of the manganese reagent. Place the flask in a large crystallizing dish containing water at 14' to 15' C., which is mounted on a magnetic stirrer having two thin asbestos sheets on its surface to serve as insulation against heat. Add 0.2 gram of zinc dust and stopper tightly with an unlubricated glass stopper. Stir for 6 minutes a t a constant rate, using a 15/rinch Teflon-covered stirring bar. Filter the solution with suction through a sintered-glass filter. After the filter has been sucked dry, rinse the flask with two small portions of distilled water and run them through the filter. Wash the filter with one small portion of distilled water. Add 1.00 ml. of the amine solution to the clear filtrate and stir well. Acidify dropwise with 1.8 ml. (about 15 drops) of concentrated hydrochloric acid and stir well. Add 1.00 ml. of the coupler solution and stir well. After 20 minutes, dilute to exactly 50.0 ml. and record absorbance readings on a spectrophotometer a t 522 mw. If a constant temperature bath is available, it may be convenient to circulate water between the bath and the crystallizing dish. The water bath described above will maintain constant temperature for sufficient time to make one determination. Quantitative formation of nitrite ion is critically dependent upon the time and rate of stirriilg. It is essential that these variables be determined on standard samples prior to attempting an unknown analysis. The stirring rate may be conveniently regulated by ac",Tsting the stirrer just below that point a t which splashing occurs. The time can then be varied by fixed increrncnts below and above the 6 minutes
recommended until the correct period is determined. The amount of zinc dust used is not critical. Once a general idea of the volume associated with 0.2 gram has been established, subsequent portions may be estimated with a spatula. Test for Nitrate Ion in Presence of Nitrite Ion. Test the sample for nitrite ion by placing about 1 ml. on a spot plate and adding several drops of the amine and coupler reagents along with a drop of concentrated hydrochloric acid. A red color indicates the presence of nitrite ion. If present, remove a 50.0-ml. sample, and add 1.00 ml. of the amine and a drop of concentrated hydrochloric acid. Stir with a magnetic stirrer for 4 minutes. Add 5 grams of the cation exchange resin and continue stirring for another 4 minutes. (This amount of resin may be more conveniently measured as one level standard teaspoonful.) After the resin has completely settled, test 1 ml. of the clear liquid for nitrite ion. If the test is positive, continue stirring until no red color is obtained. Carefully decant the solution from the resin. Run a nitrate ion determination on a 30.0-ml. sample as described above. This procedure was successfully used on nitrate ion samples containing 5.0 p.p.m. nitrite ion or less. Test for Nitrate Ion in Chlorinated Water. Test a portion of the sample for residual chlorine. If present, remove a 100-ml. sample and add 20 grams of the cation exchange resin. Stir for 20 minutes on a magnetic stirrer. Test 1 ml. again for residual chlorine. If the test is positive, continue stirring until none remains. Decant the solution from the resin. Run a nitrate ion determination on a 30.0-ml. sample as described above. If residual chlorine is Enown to be absent, dechlorination is unnecessary. The presence of residual chlorine can be determined by the o-tolidine method or by the starch-iodide reaction in neutral solution. This procedure was successfully used on samples prepared by making standard additions of nitrate ion to chlorinated drinking water. DISCUSSION
A mechanism for the formation of nitrite ion by reduction with zinc and manganous hydroxide in ammoniacal solution is not immediately obvious. Experiments conducted in this laboratory indicate that somewhat less than 70% nitrite ion recovery is realized when the reduction is performed in the absence of manganous hydroxide. Presumably, the portion not recovered is further reduced to lower oxidation states of nitrogen and possibly ammonia. The function of the manganous hydroxide may be to form a nitrite complex that is stable to further reduction a t the temperature used. A t higher temperatures the reduction evidently goes beyond the nitrite stage. Equations depicting the over-all nitrate ion reduction proceed accordingly. VOL 32. NO. 12, NOVEMBER 1960
168s
NOI-
+ Zn + 4 NHj + Ha-. NOz- + Zn(NHJ).* + 2 OH-
+H
-N( CH3)I.
+
Figure 1 illustrates the relationship between the nitrite ion and nitrate ion calibration curves. It is apparent from the diagram that 100% recovery of nitrite ion is obtainea. This is based on the premise that 1.00 gram of nitrate ion is stoichiometrically capable of producing 0.74 gram of nitrite ion. Each point on the curves is an average of five trials. Table I lists the results of standard additions of nitrate ion to natural and treated water samples. The values shown represent an average of five trials for each concentration indicated. Samples contained from about 0.8 to 3.0 p.p.m. nitrate ion. Aliquot portions were taken when the nitrate ion concentration exceeded the limit of the calibration curve. Reduction takes place in a solution having a pH of about 10.8. This falls into the range of 10.2 to 11.2 described by Middleton (6). Selection of the amine was based on its speed of diazotization, coupling ability, and stability in aqueous solution. Pinnow and Koch (7) described a simpler method of preparation for this
Table 1.
Source of Sample A*
B"
Analysis of Typical Water Samples
NOS-, P.P.M.. Deviation, Added Recovered P.P.M. 0.06 0.04
0.10
0.11
0.30 0.50
0.33 0.55
1.00 0.10
1.00
0.03
0.12
0.04
0.30 0.48 0.90
0.02 0.10
0.30 0.50 1.00
0.05
0.04
A. Blue River water collected near Manhattan, Kan., with suspended .solids removed by filtrstion and/or centnfugation. B. Manhattan, Kan., city water supply. Esch value represents average of five determinations. * Nitrite ion removed by procedure described here. c Dechlorination accomplished by procedure described here. 0
1686
ANALYTICAL CHEMISTRY
RP.M.
Figure 1.
Comparison of nitrate and nitrite absorbance curves
amine, but the product described here is more suitable for analytical work. Solutions stored for more than 6 months in a dark bottle were still usable. If a slight yellow hue forms on standing, it can be removed with carbon at room temperature. Carbon thus used may introduce a trace of nitrate ion into the amine solution, but this does not interfere with the separate nitrite or nitrate ion determinations. The nitrate ion will not interfere with nitrite ion determination, and would not be added until after the zinc reduction step in the nitrate ion determination. A blank correction would be necessary when running a mixed determination of the two ions however, since then the amine is present during the zinc reduction step. The coupler was chosen for its ability to couple in acid solution and for the formation of a favorable color. It has been previously used in alcoholic solutions for colorimetric determinations (4, 8). Low resulta were obtained in this procedure when ethyl alcohol waa used as a solvent. This presumably was due to formation of an ether from the reaction of ethyl alcohol with the diazonium cation. The aqueous hydrochloride salt solution recommended here has a tendency to undergo decomposition and is prepared best often, even daily. Decomposition, probably due to air oxidation, may deet the pure N,Ndimethyl-1-naphthylamine once theoriginal bottle has been opened. Storage over granular, unamdgamated zinc apparently minimizes this decomposition. Clarity of solution is not an indication that no decomposition has occurred. The red color of the azo dye appears to be stable. Absorbance readings taken up to 1week after development do not vary significantly. N-(1-naphthy1)ethylenediamine does not yield 35 intense a color at a given nitrite or nitrate ion concentration when used as the coupler. When removing nitrite ion, it may be
necessary to permit a longer time for completion of diazotization than the 4 minutes recommended before introducing the cation exchange resin. This period should be determined on standard solutions. Dechlorination of the treated water samples is accomplished apparently by direct chlorination and possibly some addition across any unmturation in the RSh.
The following general equations rclating concentration to absorbance were derived from the calibration curves. Equation 1 applies to nitrite ion and Equation 2 to nitrate ion.
CNO*-= 1.798 CNO,-= 2.444
(1) (2 1
where C is concentration in parts per million and A is absorbance. The molar absorptivity for the dye formed is 26,200. LITERATURE CITED
(1) American Public Health Assoc., and
American Water Works Assoc., "Standard Methods for tfi" Examination of Water and Sewage, 9th ed.,pp. 69-73,
1946. ( 2 ) Fisher, F. L., Ibert, E. R., Beckman, H. F., ANAL. CHEM.30,1972 (1958). (3) Griess, P., Ber. 12,427 (1879). (4) Marshall, E. K., Jr., J. Biol. Chem. 122,263 (1938). (5) Middleton, K. R., Chem. & I d . (Londa) 1957, 1147.
(6) Nelson, J. L., Kurta, L. T., Bray, R. H., ANAL.CHEW26,1081 (1954). (7) Pinnow, J., Koch, E., Ber. 30, 2860
( 1897). (8) Spell, F. D., Snell, C . T., "Colorimetnc Methods of Analysis," Vol. 2, 3rd ed., p. 805, Van Nostrand, New York, 1949. RECEIVED for review December 10, 1958. Resubmitted J u n e 30, 1960. Accepted August 4, 1960. Supported in par: by Research Grant D-207 from the Sational Institutes of Health. Taken in part from a thesis presented by Fred Zitomer to ihQ?Graduate Faculty of Kansas State Uiriversity in partial fulfillment of the reqviirements fer t h e degree of m s t e r of scienre.
