Spectrophotometric Determination of Nitrite And of Nitric Oxide in Furnace Atmospheres HERMAN A. LIEBHAFSKY AND EARL H. WINSLOW, General Electric Company, Schenectady, N. Y.
G
The maximum absorption of visible light by the azo dye evidently occurs a t 5200 A*, and this wave length was accordingly chosen for quantitative colorimetric work. Beer’s law in the form
RIESS (6) first discovered that nitrous acid could be detected by reacting it with sulfanilic acid (I),coupling the resulting diazonium compound with a-naphthylamine (11) to form a highly absorbing pink azo dye (111) :
f (JJ
A = log I B / I =
H
0
8
s
~
N
=
N
~
H
“2
I
I1
I11
Perhaps the most sensitive method of detecting nitrite or nitrous acid as such (6, 7, 8), the formation of this dye can also be used to measure amounts of any substance that will yield nitrite in known proportion-for example, traces of oxygen could be determined by the use of nitric oxide in large excess. Over a year ago, the problem of estimating the small amounts of nitric oxide in certain furnace atmospheres arose in this laboratory, and the authors were led to investigate the usefulness of Griess’s reagent for this purpose. Recently they discovered that Bennett (8) had already completed a similar investigation, employing visual methods where they used a recording spectrophotometer (16), and that their results in general confirm his earlier findings. Griess’s reagent, as modified by Ilovsay and Lunge, was prepared approximately according to the directions of Dennis (4). Sulfanilic acid (0.5 gram) was dissolved in 150 cc. of 5 N acetic acid; 0.1 gram of a-naphthylamine was boiled with 20 cc. of water, filtered while hot, and the filtrate added t o 150 cc. of 5 N acetic acid. The two solutions were mixed and kept in a glassstoppered bottle that stood in the dark. A known aqueous potassium nitrite solution containing 10 micrograms per cc. was prepared from the c. P. salt (84.5 per cent KN02). Standard nitrite solutions for colorimetric work were prepared by adding the proper volume of this stock solution to 20 cc. of 0.5 N sodium hydroxide, and shaking after the mixture had been acidified with 1 6 cc. of glacial acetic acid. After the further addition of 4 cc. of the Griess reagent, dilution t o 27 cc., and thorough mixing, the solutions were allowed t o stand 20 minutes before their transmissions were measured on the spectrophotometer in a cylindrical quartz cell (inside length, 5.15 em.; inside diameter, 2.40 em.).
(1) where I B is the intensity of transmitted light at 5200 A. for 2the reagent blank, I is the corresponding intensity for the sample, and rn is the micrograms of added potassium nitrite, will be used to show the concordance of the curves a t the absorption maximum. The average IC, which is used in computing m (calculated), is the arithmetical mean of all values except that for 0.5 microgram. The data show: (1) that above rn = 0.5, m and rn (calculated) are in excellent agreement: Beer’s law is obeyed within the experimental error (usually less than 2 per cent); (2) that, after the first 20 minutes, standing up to 30 minutes more has little or no effect on the results. This concordance, which might be increased even further, establishes this as one of the more accurate colorimetric methods; it shows that Bennett (8) was correct in attributing his larger (5 to 10 per cent) errors to the personal factors associated with visual comparisons. It suggests that the conversion of nitrite into azo dye, a relatively complex process, is virtually complete even a t these very low concentrations of nitrite. The colorimetric method might be a valuable tool for investigating the mechanism of the reactions involved in the conversion. The molar extinction coefficient of the azo dye can be calculated from k if complete conversion of the added nitrite is
80
z
0
5- 6 0 3,
TABLE I. CONCORDANCE OF CURVES m A k
m (calcd.)
m A k m (calod.)
20 Minutes’ Standing. IB= 2.00 1.00 4.00 0.0785 0,1488 0.3115 0.0785 0.0744 0.0779 1.02 1.93 4.03 Av. k = 0.0772 After 50 Minutes’ Standinz. I B = 0.50 1.00 2.00 -4.00 0.050 0.0777 0.1483 0.3074 0.101 0.0777 0.0742 0.0769 0.65 1.01 1.94 4.02 Av. k = 0.0766 After 0.50 0.049 0.098 0.63
76.2% 6.00 0.4620 0.0770 5.99
4
a
I-
10.00 0,7816 0.0782 10.1
IY (L
75.7% 6.00 0.4591 0,0765 5.99
k (m)
40
P
10.00 0,7753 0,0775 10.1
The standard solutions were prepared in this way mainly because aqueous sodium hydroxide was used as absorbent in the nitric oxide determinations; thus the troublesome purification of the reagents could be circumvented. Also, the sodium acetate, by decreasing the acidity, may speed up the reactions involved in the formation of the azo dye. The transmission curves for the standard nitrite solutions are shown in Figure 1.
NUMBERS SHOW MICROGRAMS OF K N O p A D D E D
I 4000
I
I
1
1
I
I
I
l
I
6000 e000 WAVE LENGTH IN ~ N G S T R O M S
I
l
l
I 7 IO
FIGURE 1. TRANSMISSION CURVESOBTAINED WITH THE POTASSIUM NITRITESTANDARD SOLUTIONS 189
INDUSTRIAL AND ENGINEERING CHEMISTRY
190
assumed. The result, E = 34,400 for 5200 A., is in close agreement Tith the maximum coefficients for dithizone (300,400 at 6200 A.) and copper dithizonate (35,600 a t 5080 A.) (13); since the magnitude of E is an index of the sensitivity of a colorimetric method, it may be that 35,000 or thereabouts is an upper limit that maximum extinction coefficients in colorimetric methods may be expected to approach; E of the zirconium-quinalizarin lake, for example, is only 10,000 (14). An extensive literature testifies that the stoichiometry of the reaction between nitric oxide and oxygen is a complex and somewhat controversial subject; no complete discussion can be given here. Baudisch and Klinger ( 1 ) first proved that the reaction
4NO
+ Oa + 4KOH = 4KNOz + 2HzO
4N0
+ 302 + 4KOH = 4KNOs + 2Hz0
(2) could be used for the determination of nitric oxide; stoichiometric conversion into nitrite occurs, provided the hydroxide is sufficiently dry (18) and the oxygen is admitted to nitric oxide that is already in contact with the alkali. If these conditions are not carefully observed, the reaction (3)
also occurs. Reaction 3 is known to involve the intermediate formation of nitrogen dioxide, which can react with water (perhaps also with the alkali) to give nitrate and nitrite, for example: (4) This earlier work would indicate that nitric oxide can never be completely converted to nitrite if it is mixed with an excess of oxygen and permitted to stand in contact with dilute alkali. It was accordingly necessary to test the stoichiometry of the reaction between nitric oxide and oxygen. Pure nitric oxide was prepared by reacting potassium nitrite, potassium iodide, and sulfuric acid (9); and was passed through 90 per cent sulfuric acid, then through 50 per cent potassium hydroxide. The resulting gas was diluted with nitrogen and added in known amounts to two 4-liter flasks containing 100 cc. of 0.5 N sodium hydroxide and about 10 cm. of air. Nitrogen sufficient t o raise the pressure to about 1 atmosphere was finally added. The flasks were shaken vigorously for a minimum of 10 minutes, whereupon duplicate nitrite determinations were made on aliquot parts of the two alkaline solutions. TABLE11. NITRITEDETERMINATIONS Sample 1 1 2 2
Nitrite Found Micrograms of KNOi 171 168 367 351
Nitric Oxide Found
Nitric Oxide Added
P. p . m. 12.7 12.5 27.2 26.0
P. p . m.. 8.0 8.0 22.8 22.8
The agreement of duplicate determinations for each sample is good, but about 4 parts per million too much of nitric oxide was found in both cases. This unexplained discrepancy may have been due to nitric oxide introduced as an impurity. At any rate, since the results are high, there is no indication of nitrate formation (Reaction 3). Bennett carried out similar tests-apparently, however, with much less shaking (2, p. 1150, line 9; and p. 1152, Figure 9)-but never obtained complete conversion to nitrite when the concentration of nitric oxide was near 10 parts per million. The authors’ results indicate that complete conversion can be obtained in a relatively short time with vigorous shaking, even in contact with aqueous sodium hydroxide and excess oxygen. A likely explanation for this apparent conflict with earlier observations (1, 11, 12) is not difficult to find. There can be little doubt that nitrate formation is usually preceded by the formation of nitrogen dioxide (Reaction 4), since nitrate formation is enhanced by the presence of water or of excess oxygen. The
VOL. 11, NO. 4
gaseous reaction by which nitrogen dioxide is formed from nitric oxide and oxygen follows the law (10) --d(Oz)/dt = k (NO)’ (02)
(5)
Equation 5 predicts that nitrogen dioxide formation will be retarded by reducing the nitric oxide pressure; consequently no nitrate may have formed in the authors’ experiments because the nitric oxide pressure was too low. But the whole subject deserves further attention-a kinetic investigation, using the colorimetric method, of the oxidation of nitric oxide a t very low pressures by oxygen in the presence and absence of aqueous sodium hydroxide should yield interesting and important information. Samples of furnace atmospheres were successfully analyzed by the method used in the experiments with.nitric oxide. When no oxygen detectable in an ordinary gas analysis was present, results below 1 part per million of nitric oxide were consistently obtained, showing that the discrepancy of 4 parts per million encountered above is not inherent in the analytical method. When a gas containing 0.7 per cent oxygen, 9.7 per cent carbon dioxide, no carbon monoxide, and 89.6 per cent nitrogen issued from a furnace in which the maximum temperature was 1390” C., about 80 parts per million of nitric oxide were found. This concentration is about .one seventh the equilibrium concentration for the reaction a t 1390” C. and far above the concentration for room temperature (S), indicating that the mixture of gases cooled rapidly enough to slow up the dissociation of nitric oxide (of course, equilibrium a t 1390” C. may never have been established). The colorimetric method might be used to study the rate of the reverse, and perhaps of the forward, reaction in the important Equilibrium 6. Attempts to collect the nitric oxide by passing the furnace gas mixed with air through liquid air and then through a bubbler containing dilute sodium hydroxide were unsuccessful, probably because there was insufficient time for the conversion of nitric oxide to nitrite. Summary The determination of very small amounts of nitrite with Griess’s reagent (sulfanilic acid and a-naphthylamine in acetic acid solution) has been studied with a recording spectrophotometer and found to be one of the more accurate colorimetric methods. The same reagent has been used successfully for the determination of nitric oxide in concentrations well below 10 parts per million. The possible use of the colorimetric method in several other connections has been pointed out. Literature Cited (1) (2) (3) (4) (5) (6) (7) (8) (9)
(IO) (11) (12) (13) (14) (15)
Baudisch and Klinger, Ber., 45, 3231 (1912). Bennett, J.Phys. Chem., 34, 1137 (1930). Clayton and Giauque, J . Am. Chem. Soc., 55, 4875 (1933). Dennis, “Gas Analysis,” p. 218, New York, Maomillan Co., 1913. Feigl, “Qualitative Analyse mit Hilfe von Tiipfelreaktionen,” p. 332, Leipeig, Akademische Verlagsgesellschaft, 1938. Griess, Ber., 12, 427 (1879). Hahn, Mikrochemie, 9, 31 (1931). Hahn and Jaeger, Ber., 58,2335 (1925). Johnston and Giauque, J . Am. Chem. SOC.,51, 3194 (1929). Kassel, “Kinetics of Homogeneous Gas Reactions,” p. 167, New York, Chemical Catalog Co., 1932. Klinger, Bet-., 46, 1744 (1913). Klinger, 2. angew. Chem., 27,7 (1914). Liebhafsky and Winslow, J . Am. Chem. Soc., 59, 1966 (1937). Ibid.,60, 1776 (1938). Michaelson and Liebhafsky, Gen. Elec. Rev., 39, 445 (1936).
