following very small changes in even highly conducting solutions, it opens the door to conductometric analysis in other systems (such as redox reactions and ion exchange monitoring) where it has previously been impossible or very impractical to make conductance measurements. DISCUSSION We have just shown two of the many possible applications for this conductance instrument. This instrument can be used anywhere that a conductance bridge apparatus can be used, including the accurate determination of equivalent conductance or cell constants. The absolute accuracy is limited to the accuracy of a reference resistor used for standardization, and the accuracy of the voltage measuring device. The precision and noise level is less than 0.002 over most of the resistance range (200 Q-1 MQ). With the present instrument, a precise measurement can be made every 50 psec if necessary,
and that value held with a drift of only 0.1 of full scale/sec. With a repetition rate of at least 100 Hz, a continuous analog output can be obtained. This method causes very little solution heating, as the power output is very low (0.2 pwatt/Hz maximum); thus, constant temperature is easy to obtain. And, most important, there is no capacitance balancing involved as this method is completely independent of the parallel cell capacitance and almost entirely independent of the double layer capacitance. Thus, this method should prove very useful for fast, accurate conductance measurements under a wide variety of experimental conditions. RECEIVED for review September 16, 1969. Accepted December 15, 1969. One of us (D. E. Johnson) gratefully acknowledges a National Science Foundation Traineeship. This work was partially supported by NSF Basic Research Grant GP-3404.
Direct Spectrophotometric Simultaneous Determination of Nitrite and Nitrate in the Ultraviolet James. H. Wettersl and Kenneth L. Uglum Central Michigan Unisersity, Mount Pleasant, Mich. 48858 The ratio of absorbance of aqueous sodium nitrite at 355 m p to that at 302 mp is 2.50 & 0.02. As nitrate does not absorb at 355 mp but has a characteristic band at 302 mp, absorbance due to nitrate can be calculated by dividing nitrate absorbance at 355 m p by 2.50 and subtracting the quotient from total absorbance at 302 m p . I n 1-cm cells the lower detection limit for nitrite is 0.02 mg/ml and for nitrate, 0.09 mg/ml. Below pH 5, nitrite forms nitrous acid, with an absorbance maximum at 357 mp which i s twice as sensitive to concentration as the 355 mp nitrite peak. Effects of heat, time, and reagents at various concentration levels are tabulated. Molar absorptivity of nitrite at 355 mp is 23.3 f 0.8; at 302, 9.12. At 302 m p nitrate has a molar absorptivity of 7.24 i 0.04.
NITRITE AND NITRATE have been determined simultaneously in the ultraviolet by three methods. Hamaguchi, Kuroda, and Endo ( I ) resolved a simple two-component system of nitrite and nitrate at a pH greater than 5 by using two sets of standard curves of absorbance us. concentration at their characteristic wavelength maximums. They found that at a pH less than 5, the characteristic peak of nitrite disappears and unique peaks of nitrous acid begin to develop. They did an interference study with hydrogen ion concentration and various salts. Recovery of nitrite and nitrate was good, and they were able to resolve interference due to aluminum by making the solutions basic. Meerman (2) found that nitrite in a mixture could be determined directly at 353 mp; after reducing the nitrite with sulfamic acid, the nitrate could be determined at 302 mp. If the nitrite concentration was less than half the concentration of the nitrate, the step involving sulfamic acid was not necessary. Haddad and MacDonald (3) resolved a mixture of nitrite and nitrate in 4.9M potassium hydroxide, using a constant ratio of 1
Present Address, Dow Corning Corporation, Midland, Mich.
48640 ~
~
(1) H. Hamaguchi, R. Kuroda, and S. Endo, Bunseki Kaguka, 7, 409 (1958); C.A., 54, 7421e (1960).
(2) G. Meerman, Dissertation Abstr., 20,4507 (1960). (3) L. Haddad and J. C. MacDonald, “The Simultaneous Deter-
mination of Nitrite-Nitrate Using Ultraviolet Spectrophotometry,” Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Cleveland, Ohio, March 1969.
3.37 f 0.03 of molar absorptivities at 356 and 301 mp for nitrite, They also found that the molar absorptivity of nitrate at 301 mp and nitrite at 356 mp decreased with increasing potassium hydroxide concentration. The value of E for nitrite at 301 mp increased with increased KOH concentration. They found E for KNOz at 356 mp to be 21 .O i 0.3 and at 301 mp, 6.17 & 0.10; that for KNO, was 6.13 i 0.35. Vandenbelt et af. ( 4 ) determined the molar absorptivity of KNO, to be 7.064. The method of the present paper is based on the observation that the ratio of absorbance of aqueous sodium nitrite at 355 mp to that at 302 mp is constant at 2.50. Nitrate does not absorb at 355 but has a characteristic band at 302. This technique is more applicable to computer calculation than the methods of Hamaguchi and Meerman. A study was made of the effects of acids, bases, salts, reducing agents, and heat. EXPERIMENTAL
Apparatus. A Cary 14 spectrophotometer with 1-cm quartz cells was used. Reagents. Dilute nitric acid was prepared from concentrated, reagent grade nitric acid. Reagent grade sodium nitrite and sodium nitrate were used in preparing aqueous nitrite and nitrate solutions of different concentrations. Other concentrated acids, bases, and salts of various kinds were also of the highest quality available. Procedure. The spectrum was scanned from 400 mp downward at X 10 speed to 290 mp on solutions of nitrite and nitrate having absorbances between 0.1 and 1 . Mixtures of nitrite and nitrate were also scanned in the same absorbance range. An interference study was conducted by making dilutions in the presence of various acids, bases, salts, and reductants in a concentration range from to 1.OM. In each case the interference study was done with nitrite and nitrate by themselves and in mixtures, at the concentration of 2.14 X 10+M sodium nitrite and 6.90 X 10-2M sodium nitrate. These concentrations were chosen to give an absorbance of 0.5 in 1.0-cm cells. These concentrations were obtained by diluting a 15.00-ml aliquot of a 0.357M sodium ~~
~~
~~~
~
~
(4) J. M. Vandenbelt et al., IND. ENG. CHEM.,ANAL.ED., 17, 235 (1945). ANALYTICAL CHEMISTRY, VOL. 42, NO. 3, MARCH 1970
335
~~
_
_
_
_
~
~~
~
Table I. Absorbance of HN03at 302 mp Concn HN03 Absorbance E 1 X 10-2M 0.072 7.19 5.2 x 0.372 7.14 1.0 x 10-1 0.729 7.29 1.3 X 10-1 0.934 7.19 1.37 x 10-2 0.099 7.24 8.22 X 10-2 0.595 7.23 1.22 x 10-1 0.880 7.21 ~
~~
Table 11. Absorbance of NaN03 at 302 mp Concn NaN03 Absorbance E 7.18 X 10-2M 0.521 7.26 1.0 x 10-1 0.725 7.25 1.30 X 10-l 0.942 7.23 4.04 x 10-2 0.291 7.21 Table 111. Absorbance of NaNOz at 355 mp Concn NaNOz Absorbance E 2.23 X 10-2M 0.521 23.4 3.15 X 0.721 22.9 4.02 X 0.927 23.1 3.83 x 10-3 0.092 24.0 2.11 x 10-2 0.485 23.0 4.07 x 10-2 0.944 23.2 nitrite solution to 250 ml and a 15.00-ml aliquot of 1.15M sodium nitrate solution to 250 ml. In each interference set, a control sample of the known nitrite-nitrate combination was run and used in the calculation of percent recovery based on absorbance. Two aliquots of a 0.357M stock solution of sodium nitrite were placed in beakers and evaporated to dryness at 105 "C. One was placed in another oven at 200 "C for 1.5 hours. Both were checked for recovery. Nitrogen dioxide generated by reaction of concentrated nitric acid with copper was passed into water at 25 "C and approximately 100 "C, 0.25N NaOH at 25 "C and 100 "C, and 1N H2S04 at the same temperatures. The resulting solutions were diluted to the proper concentrations and spectra run on the Cary 14. Solutions of nitrite and nitrate having absorbances of 0.02, 1.4, and 1.9 were prepared to determine the detection limits. The molar absorptivity was determined on various nitrate salts at a concentration of approximately 6.9 X 10F2M. Dilute solutions at 10-4M of sodium nitrite and nitrate were used in scanning down to 200 mp. Nitrite was found to have a band with molar absorptivity of around 4400 at 210 mp, while nitrate showed strong absorption at 204 mp with a molar absorptivity of about 9700. Strong absorption by water in this region makes these values approximate.
RESULTS AND DISCUSSION Tables I, 11, and I11 show absorbances of pure substances at various concentrations. Molar absorptivity of nitric acid is 7.21 + 0.09 (for 2 sigma); that of sodium nitrate is 7.24 + 0.04; and that of sodium nitrite, 23.3 f 0.8. For nitrate, the values found for nitric acid and sodium nitrate are in disagreement with those obtained by Vandenbelt (4) and by Haddad and MacDonald (3). Data for separate solutions of nitrite and nitrate prepared to give absorbances of approximately 0.72 at their respective maximums are given in Table IV. Absorbance data for solutions of sodium nitrite and sodium nitrate after mixing are also shown in Table IV. Note that absorbance data for pure sodium nitrite are given at 355 and 302 mp for the same concentration. The absorbance of the mixture at 302 mp corrected 336
ANALYTICAL CHEMISTRY, VOL. 42, NO. 3, MARCH 1970
Table 1V. Determination of NO3- and NO2- in a Mixture Concn of anion X Max. Absorbance E 1 X 10-'MNOa302 mp 0.725 7.25 3.15 X 10-2MNOz355 0.721 22.9 3.15 X 10-2MNOz302 0.287 9.12 Mixture of NO3- and NOZ1 X 10-lMNOs302 1.010 ... 3.15 X 10-2MNOz355 0.730 23.1 1.010 ,4302 of mixture - 0.287 A 3 0 2 of nitrite = 0.723 A302of nitrate. Hence E of NO3- in mixture is 7.23. Table V. Comparison of Sensitivity of 355-mp Maximum with 302-mp Maximum of NO2Absorbance Concn of NO2302 mp 355 mp A3dA3~2 2.23 X 1W2M 0.215 0.532 2.48 3.15 x 0.287 0.721 2.51 0.936 2.50 4.02 x 0.375 3.83 x 10-3 0.037 0.092 2.49 4.07 X 1W2 0.377 0.944 2.50 Table VI. Effects of Various Acids in Presence of 2.1 X 10-+M NaNOz Acids giving nitrous acid spectra at stated concentrations Concentration, Acid M Acid Concentration, M Tartaric 0.5 1.04 HClO4 KHSO4 0.5 0.972 HCl Na2SiFs 0.03 (pH3.8) 1.04 HzSOa 0.05 SnC14 0.125 (pH 0.4) SnC14 0.25 (pH0.2) HOAc 1.03 0.05 (pH 1.4) 0.893 HCOOH KzSz08 0.5 Alz(S04)a 0.1 (PH 2.9) CHzClCOOH MnSO4 0.25 (pH2.3) Citric 0.25 0.5 KHCeHaOa 0.25 (PH 3.8) Oxalic 0.982 &PO4 Acids not converting nitrite to nitrous acid at the stated concentrations Percent recovery, Acid Concentration, M 355 mp 107 NaH2P04 0.5 (pH 4.4) 103 Na2HP04 0 . 5 (pH 8.8) 101 Na3P04 0.25 (pH 10.9) 0.1 (pH 5.7) 99 101 &Boa 0 . 5 (pH 4.8) 103 0.54(pH 4.6)
for the absorbance for nitrite alone gives the absorbance of nitrate at the same wavelength. This corrected absorbance is in agreement with the absorbance of the nitrate by itself. In Table V the absorbance of nitrite at 302 and at 355 mp is recorded for various concentrations. The ratio of absorbance at 355 mp to absorbance at 302 mp was calculated to be 2.50 f 0.02. Concentration of nitrite in the solution is determined from the absorbance at 355 mp. This absorbance is divided by 2.50, with the quotient subtracted from the total absorbance at 302 mp to give the absorbance at 302 mp due to the nitrate. Based on an absorbance of 0.01, the lower detection limit of nitrate is 0.09 m g / d in I-cm cells and for nitrite 0.02 m g / d in 1-cm cells. When 10-cm cells are used, the lower detection limit would be 9 pg/ml for nitrate and 2 pg/ml for nitrite. In the working range of 0.02 to 1 absorbance, the concentration range is 0.2 to 9 mg of nitrate per ml; that of nitrite is 0.4 to 2 mg per ml. This is in good agreement with Hamaguchi ( I ) .
Table VII. Decomposition of Nitrite by Acids Acid HCl Hap04
Concn, M 0.97
5 7 0 7 0 7 0 6
0.93
HOAc
1.03
HCOOH
Day 0
0.89
A302
A336
A346
A351
A311
0.073 0.121 0.176 0.063 0.108
0.426 0.046
0.675 0.065 0.083 0.722 0.326 0.691 0.051 0.175 0
0.932 0.094 0.078 1.004 0.446 0.922 0.066 0.238 0
0.977 0.097 0.063 1.053 0.467 0.910 0.068 0.248 0
...
0.451 0.210 0.461 0.038 0.109 0
0.093 0.136 0.052 0.050
A385
... 0.056 0.036
... 0.273
... 0.038
... 0
Table VIII. Effect of Hydrochloric Acid Concentration on Recovery of 2.14 X 10-zM Sodium Nitrite Concn of HC1 PH (by (by dilution) measurement) A365 Recovery, A302 Recovery,
z
6.6 6.3 5.7 4.4 2.5 1.3
10-6 10-5
10-4 10-3
10-2 10-1 a
0.517 0.521 0.517 0.543a 0. 775n 1.044“
0.201 0.201 0.193 0.187” 0. 127a 0.045”
100.0 100.7 100.0 105.0 149.9 202.2
z
103.6 103.6 99.5 96.4 65.5 23.2
Acid formation caused interference. Nitrous acid has a maximum at 357 mp instead of at 355 mp.
In Tables VI through XVIII are shown effects of various acids, bases, salts, and reductants on recovery of nitrate and nitrite. In Table VI, in cases in which hydrogen ion concentration is greater than 5 X loe2, the spectrum obtained was that of nitrous acid only. At pH less than 5, the nitrous acid begins to appear. This is in agreement with the observations of Hamaguchi ( I ) . Acid salts such as SnC14,KzS2Os,MnS04, and Ai2(S04)3gave pH values less than 5 and therefore produced the nitrous acid spectrum. Basic solutions of these salts do not cause interference with nitrite. Table VI1 shows the influence of four acids in decomposing nitrite over a period of a week. Nitrous acid has absorbance maximums at 372, 357, and 346 mp. There are also absorbance shoulders at 385 and 336 mp. The formation of nitrous acid is explained by Equation 1:
+
H+ NOz“02 (1 ) The dissociation constant, K, for “ 0 2 is 4.5 X (5). Absorbance at the maximum of HNOz at 357 mp is approximately twice as sensitive as that of nitrite at 355 mp. The decrease in absorbance of HNOz with time is explained by the following equations.
+ NOz + HzO H+ + NOS- +
2HN02 -+ NO
3HN02 + 2N0 HzO
+
AGO = 2.06 kcal(6) (2) AGO = -3.22 kcal (6) K = 30 (7)
(3)
Reaction 3 explains the increase of absorbance at 302 mp which is due to the formation of nitrate. Table VI11 shows the effect of various concentrations of acid on nitrite. In acid solution there is no longer a maximum at 355 mp, but there is one at 357 which is due to the nitrous acid.
Table IX. Effect of Various pH Buffers on Recovery of 1.79 X 10-*M Sodium Nitrite pH of Recovery, Recovery, A355 A302 buffer
z
z
10 6 5 4.63 4 3
0.451 0.420 0.421 0.455 0.456” 0.613
103.4 101.7 101.8 104.4 108.7 148.4
0.188 >2 >2 0.177 >2 >2
115.2
...
109:3
... ...
Maximum shifted to 357 mp.
Because nitrous acid has a greater molar absorptivity than nitrite, recovery increases with increasing acid concentration. The recovery at 302 mp decreases because of conversion of nitrite to nitrous acid. A recovery study was performed in the presence of various commercial buffer solutions as shown in Table IX. Again, nitrous acid spectra formed at pH less than five. In the cases in which the absorbance was greater than two at 302 mp, this was due to unidentified species. In very dilute solutions, these species had absorption maximums at approximately 270 mp. Nitrogen dioxide absorbed in water forms nitrous acid and nitric acid, according to the following equation: 2N02
+ H20 (cold)
-+
“02
+
“03
K
=
IO5 (7) (4)
The spectra obtained by absorbing NO2 in hot water indicate the presence of both nitrous and nitric acids, in disagreement with the equation (8) 3N02
+ H20 (hot)
-+
2HN03
+ NO
( 5 ) N. V. Sidgwick, “The Chemical Elements and Their Com-
AGO = -12.62 kcal (6) (5)
pounds,” Clarendon Press, Amen House, London E.C. 4, 1952,
Nitrogen dioxide is much more soluble in cold water than in hot. Nitrogen dioxide is also soluble in 1N sulfuric acid, producing nitrous and nitric acids in the same proportions as in
p 694. (6) W. M. Latimer, “Oxidation Potentials,” Second Ed., PrenticeHall, Inc., Englewood Cliffs, N. J., 1952 (fourth printing, 1859), p 92. (7) W. M. Latimer and J. H. Hildebrand, “Reference Book of Inorganic Chemistry,” Third Ed., Macmillan, New York, 1951, p 207.
(8) H. T. Briscoe, “College Chemistry,” Fourth Ed., Houghton Mifflin Company, Boston, Mass., 1951, p 349. ANALYTICAL CHEMISTRY, VOL. 42, NO. 3, MARCH 1970
337
creases, both at 355 and at 302 mp. This is in disagreement with the results of Haddad and MacDonald (3). Interference at 355 mp was outside the limits of 100 4% for bivalent ions of barium, zinc, copper, and lead, as shown in Table XI. In the cases of copper and lead, color effects were observed. The copper went from blue to green and the lead from colorless to yellow. The absorbance was greater than 2 instead of the expected value of 0.5. Effects of five reducing agents on nitrite and nitrate, alone or in mixtures, are shown in Table XII. In general, at concentrations of lo-* or less they did not react with nitrite. At concentrations greater than there was significant interference. No interference with nitrate was observed with hydrazine sulfate or hydroxylamine hydrochloride. There was no interference with nitrate by sodium hydrogen sulfite and sodium metabisulfite when these reductants were of 10-2M concentration. Recovery of nitrate cannot be obtained in the presence of 10-ZM L-ascorbic acid because of the ultraviolet absorbance of this reductant. With reductant concentration of 10-l and greater, recovery cannot be evaluated because hydrogen sulfite, metabisulfite, and L-ascorbic acid have strong absorption bands overlapping the band at 302 mp.
Table X. Effects of Various Bases on Recovery of 2.14 X 10+M Sodium Nitrite ReReConcn, covery, covery, Base M Aas5 Aaoz Ammonia 0.91 0.508 100.4 0.198 99.0 Sodium hydroxide 1.01 0.526 100.3 0.206 100.5 Lithium hydroxide 1.00 0.526 100.3 0.205 100.0 Potassium hydroxide 1.00 0.525 100.2 0.198 96.6 Potassium hydroxide 2.98 0.560 108.1 0.210 104.4 Potassium hydroxide 4.97 0.583 112.5 0.215 106.9
*
z
x
water. Nitrogen dioxide absorbs in 0.25M NaOH to form nitrite and nitrate in equal concentrations (8): 2N02
+ 2 OH-
NO,
+ NOS- + HzO
(6) Table X shows a recovery study of nitrite in the presence of bases. Recovery increases as the base concentration in-t
~
Salt Ammonium chloride Barium acetate Barium acetate Barium chloride Calcium chloride Cupric chloride Magnesium perchlorate Potassium chloride Sodium chloride Sodium oxalate Sodium potassium tartrate
~
~
~
~~~
~~~
0.526 0.520 0.561 0.530 0.516 >2
101.6 102.7 108.3 104.6 101.7
...
0.917 0.200 0.197 0.204 0.217 >2
98.0 100.0 98.0 101.9 108.5
0.25 1 1 0.215
0.499 0.527 0.515 0.497
98.7 101.7 98.3 97.8
0.197 0.197 0.192 0.240
98.5 98.0 93.7 123.0
0.25
0.536
103.1
0.217
111.3
0.533 0.531 0.423
102.7 102.4 83.7
0.206 0.197 0.245
102.4 98.0 122.5
0.5 1 0.25
Table XII. Effects of Various Reducing Agents on Recovery of 2.14 x 10-2M Sodium Nitrite and/or 6.90 X 10-2M Sodium Nitrate Concn, M A355 Recovery, Z A302 10-4
10-8 10-2 0.25
Hydroxylamine hydrochloride
10-4 10-3
10-2 0.5
Sodium hydrogen sulfite
10-3
10-2 10-1 0.5 10-8 10-2 10-1 0.25 10-3 10-2
Sodium metabisulfite L-Ascorbic acid
~
~
...
Recovery,
0.498 0.487 0.305 2 0.729 0.668 >2 >2 0.772 >2 >2 >2
102.9 96.8
...
