materials were very carefully analyzed for reference purposes by different laboratories. The new method is applicable to niaterials containing from 0.1 to 100% sulfate. The titration with E D T A and Eriochrome Black T as indicator is characterized by high precision. The error in determining the end point was usually less than 0.01 ml. of a n 0.05di solution. The main advantage of the complexometric method over the classical barium sulfate method lies, however, in its rapidity; a determination of sulfate b y the latter method usually takes
1 t o 2 days while the new procedure gives results in 30 to 45 minutes (time depending mainly on the pretreatment required for the kind of material tested). The precipitation of sulfate with lead is yirtually free of interference and the procedure is therefore directly applicable to ores, concentrates, extracts, solutions, etc., which contain large amounts of uranium, iron, and other cations and anions normally found in smaller concentrations.
Mining and Refining, Ltd., for perniission to publish this paper.
ACKNOWLEDGMENT
RECEIVEDfor review August 12, 1957. Accepted February 13, 1958.
The author n-ishes to thank Eldorado
LITERATURE CITED
(1) Anderegg, G., Flaschka, H., Sallmann, R., Schxarzenbach, G., Helv. Chim. Acta 37, 113 (1954). ( 2 ) Fritz, J. S., Yamamura, S. S., Richard, M. J., AXAL.CHEJI.29, 158 (195i). (3) Kenny, F., Kurtz, R. B., Beck, I., Lukosevicius, I., Ibid., 29, 543 (1957).
Redox Behavior of Cobalt Chelates of Nitrilotriacetic Acid K. L. CHENG’ Westinghouse Electric Corp., East Pittsburgh, Pa. blntensely colored complexes are formed by the oxidation of the (ethylenedinitri1o)tetraacetic acid (EDTA) or nitrilotriacetic acid (NTA) complexes of iron(ll1) and cobalt[l\) with hydrogen peroxide in alkaline medium. Iron(ll1) -NTA i s not oxidized to a bluish or violet complex by perborate, but the cobaltic-NTA complex is. No other metals give a red or violet color under similar conditions. This colored complex i s not reduced by common reducing agents but by ferrous sulfate, which exists as the iron(ll)-NTA complex with high reducing power. Factors affecting color formation and possible analytical applications to the determination of hydrogen peroxide, NTA, and other compounds are presented.
hydrogen peroxide is used to oxidize either cobalt or iron in the presence of nitrilotriacetic acid (NTA), colored complexes are formed. However, when sodium perborate is used as the oxidizing agent, it gives a colored complex only for cobalt. The cobalt(II1)-STA complex is stable in both acid and alkaline media. The color reaction is highly selective and may be used for the spectrophotometric determination of cobalt. It may also be used to determine peroxide, perborate, nitrilotriacetic acid, or (ethylenedinitri1o)tetraacetic acid (EDTA). The iron(I1)-NTA complex is a strong HEX
1 Present address, Utica Metalc Division, Kelsey-Hayes Co., Utica 4,X. Y.
reducing agent and can be used to reduce the cobalt(II1)-NTA complex in alkaline medium. Besides (ethylenedinitri1o)tetraacetic acid and nitrilotriacetic acid, additional color tests for iron and cobalt were made with several recently reported complexing agents (IO). The results shown in Table I indicate that all of the complexing agents tested except diaminocycloheuanetetraacetic acid, without the addition of a n oxidizing agent, gave a yellow or reddish brown color viitli iron(II1) in ammoniacal medium. Ethylene glycol-bis - (0- aminoethyl ether)-,Y,S’-tetraacetic acid complex of either iron or cobalt did not show any color change upon addition of perborate. Cyanide decomposes tlie (ethylenedinitri1o)tetraacetic acid complev of cobalt (3, 6, 7 ) : It was found to have the same effect on the nitrilotriacetic acid complex of cobalt. The nitrilotriacetic acid complex of iron(111) (not perosy compound) is decomposed by phosphoric acid, which also masks the color of ferric ions. The blue color of copper(I1) if present, can be masked by thiosulfate. Pribil and AIalicky ( 8 ) used cerium(IV) to oxidize cobalt in the presence of (ethylenedinitri1o)tetraacetic acid in an acetic acid medium. The conditions for the oxidation must be carefully controlled. The oxidation cannot be conducted in the alkaline medium because of the precipitation of cerium as hydroxide. Pribil and Malicky also reported interference with oxidation due to the presence of manganese and nickel.
EXPERLMENTAL
The presence of certain metals required addition of varying amounts of complexing agents for complete oxidation (see Table IV). Among the metals tested, only large amounts of manganese(11) gave difficulty in the coniplete oxidation, Too much free nitrilotriacetic acid should be avoided. More than 2.5 mrnoles of this acid, being only slightly soluble in an acid medium, will precipitate in 0.5M phosphoric acid. The perborate oxidation of cobalt can be carried out in the presence of a large number of ions with the exception of cyanide and sulfide; these form a very stable complex or precipitate with cobalt. Preliminary experiments showed that the nitrilotriacetie acid complex of cobalt(II1) was not reduced in the alkaline medium by common reducing agents such as sulfite, hydroxylamine hydrochloride, ascorbic acid, nitrite, or thiosulfate, but it was readily reduced by ferrous sulfate. I n the presence of excess conipleving agent, the nitrilotriacetic acid complex of iron(I1) was formed upon addition of the ferrous salt. The strong reducing action of the complex was due to its existence in the lower valence state in the comple\iag compound ( I , 2, 4 ) . Pribil and Svestka (9) used chromous chloride to reduce the (ethylenedinitril0)tetraacetic acid complex of cobalt(II1) in the sulfuric acid medium. Chromous chloride is a poIverfu1 reducing agent, but is very unstable and not conimercially available. An attempt was made to titrate the VOL. 30, NO. 6, JUNE 1958
1035
Table 1.
Color Reactions of Iron and Cobalt with Complexing Agents in Ammoniacal Solution
Iron
Cobalt
Complexing Agent Sitrilotriacetic acid (Ethylenedinitri1o)tetraacetic acid Diethylenetriaminepentaacetic acid
Unoxidized Yellow Reddish b r o m Lemon yellorv
Oxidizeda Yellow Deep purple Deep purple
Lnoxidized Light pink Light pink Light pink
-1'-Hydroxyethylethylenediaminetriaceticacid
Tellon
Light pink
Diaminocyclohexanetetraacetic acid Ethylene glycol-bis-(p-aminoetliyl ether)-N,N'tetraacetic acid a Perborate oxidntion.
Brownish ppt. I-elloiv
Light wine red (not stable) Brownish ppt. Tellobv
Oxidized" Purple Violet blue Very light violet blue Violet blue
Very light j-ellox Very light pink
Purplish red Very light pink
Table II.
Formal Potentials of Complexes of Iron and Cobalt
Reaction
E", Volt 0.ii
FeEDT.1-
0.36 0.329 (12) -0.11 (ca. j Present n-ork, carbonate medium 0.117 ( I f )
+ e = FeEDT.1--
++ + + +
0.1 ( 2 ) -0.16 (Fa.) (Z), sodium h\-tlroxide and tartrate
duction and precipitation by ferrocyanide were very slow in sulfuric acid medium, because the pH was too lo^ to permit the formation of cobalt ferrocyanide precipitate. Cobalt(I1) was also precipitated by ferrocyanide in the presence of nitrilotriacetic acid in acetic acid or phosphoric acid medium.
medium
Co+-- (as ) e = Co+' (xj CoEDT.1e = CoEDTh-COST.% e = COSTACo(OH)3 e = C O ( O H ) ~ OHCO(NHJ),+++ e = C O ( S H ~ ) ~ + e = co(Cs)&-' Co(cx)6---
1.8 0 6 (ca.) (8), acetic acid medium
SPECTROPHOTOMETRIC PROCEDURES
0.35 (ca.) Present work, carbonate medium
+
0.2 0.1 -0.83
+
nitrilotriacetic acid complex of cobalt(111) potentiometrically with ferrous sulfate in the alkaline medium. The nitrilotriacetic acid complex of cobalt(11) m s oxidized by hydrogen peroxide in the ammoniacal or sodium carbonate medium; sodium sulfite was added to reduce the excess peroxide. The results obtained from the titration curve were constantly 10 to 20% higher than the expected values (one electron change). It is riot easy to determine the forrnal potentials of iron(I1)-NTd-iron(II1)ST.4 and cobalt(I1)-KTA-cobalt (111) -STA accurately at high pH. For practical purposes they were e-timated
potentiometrically in the sodium carbonate medium. Those values were not measured under standard conditions, and only approximate values are given for comparison. Table I1 shows that the oxidation potential is greatly influenced by complexation and that iron(11)-nitrilotriacetic acid and iron(I1)EDTA are powerful reducing agents. An attempt was made to use ferrocyanide as a titrant. The nitrilotriacetic acid complex of cobalt(II1) was reduced by ferrocyanide x i t h the formation of precipitates in acetic acid or phosphoric acid medium. The re-
The formation of the violet cobalt(111)-XTA conip1e.r may also be applied to the determination of hydrogen peroxide, sodium peroxide, and sodium perborate. Reagents. Cobalt nitrate solution, 0.025 and O . l l l l . Kitrilotriacetic acid, 0 . 5 X . Sodium perborate tetrahydrate. Hydrogen peroxide solution, 0.25 mg. per ml. Cobalt. Prepare a calibration curye by pipetting 0, 2.0, 4.0. 8.0, 12.0, and 16.0 ml. of 0.025M cobalt nitratp solution into SO-nil. volumetric flask.. Dilute t o about 25 ml. and add 2 nil. of 0.5M S T A solution; adjust t h e pH to 7 . 5 t o 9.0 with sodium carbonate or sodium bicarbonate in a Becknian Model G (or equivalent) pH meter. Add approximately 0.1 gram of sodium perborate. Let t h e color clevelop for 10 minutes or longer, then
.r530 d
385
\
fNH,OHANDH,Pc,) ,547
w 0.6 V
-
(",OH)
4
m LO 0
m
a
0.4
-
0.2
-
0.ot 300
I
400
I
0.0
I
500 600 WAVE LENGTH,M#
700
Figure 1. Absorption spectra of cobalt (11)NTA complex in phosphoric acid and ammonium hydroxide
1036
ANALYTICAL CHEMISTRY
I
300
I
I
I
1
400
500
600
700
WAVE LENGTH, M A
Figure 2. Absorption spectra of cobalt (11)NTA complex in phosphoric acid and sodium carbonate
dilute to volume. Measure the absorbance a t 580 mp on a Beckman Model B (or equivalent) spectrophotometer. If it is desired to measure the absorbance at 560 mp in an acid medium, add 4 nil. of 8-11 phosphoric acid or 2 ml. of 9-11 d f u r i c acid before dilution. The absorption and calibration curves are shon n in Figures 1through 4. Beer's l a x i b obeyed in the regions specified. Hydrogen Peroxide and Sodium Perborate. T o determine hydrogen peroxide, prepare a calibiation curve from standard solutions. Add 0, 2.0, 4.0, 6.0, and 8.0 nil. of hydrogen peroxide solution to 50-ml. volumetric flasks containing 2 nil. of 0.1.V cobalt nitrate, 2 ml. of 0.5M N T A solution, and 2 nil. of 1Oyo sodium bicaibonate. Dilute to volume and let stand for 5 minutes or longer. XIeasure the absorbance a t 580 nip. Beer's law is followed from 0.32 to 1.28 nig. of hydrogen peroxide per nil. Prepare the calibration cur\-e for sodium perborate in the same way, using 0, 2 , 4, 6, 8, and 10 mg. of sodium perborate crystals of known hydrogen peroxide content ( 5 ) . The solutions conform to Beer's law from 1 to 14 mg. of perborate per 25 ml. The method can also be used to determine the peroxide content of sodium peroyide. Titration of Cobalt with Hydrogen Peroxide. To a solution containing 0.1 t o 1 mmole of cobalt, add approximately fire times as much nitrilotriacetic acid. ,idjust to pH 7.5 t o 9 n i t h sodium carbonate and dilute
to approximately 200 ml. Titrate with 0.04M hydlogen peioxide a t 580 mp.
,4 curve for a typical titration for cobalt is shoim in Figure 5 . For 0.2 nimole of cobalt, a titer of 5.16 ml. of 0.0486.M hydrogrn peroxide \I as obtained, compared n ith a theoretical valuc of 5.18 nil. DISCUSSION
pH. The oxidation of the complexes of cobalt(I1) and iron(II1) depends upon a t least three factors: (1) a strong coniplexing agent (see Table I ) ; ( 2 ) a suitable oxidizing agent; and (3) pH of the medium. When peroxide or perborate and iiitrilotriacetic acid are used, the oxidation must he done in the alkaline medium. After the oxidation is completed in an alkaline medium, the cobaltic complex is not easily dccomposed by acid; howewr, the acidity affects the absorption niasinium as shown in Figures 1 and 2 . K h e n sodium carbonate was used for the adjustment of pH (pH 8.5 t'o 9), the color was violet rather t'han reddish purple, as was t'he case n-hen ammonium hydroxide was used (pH about 11). This violet color changrd to reddish purple after addition of acid. For spectrophotometric dcterniination, t,he color of thr conipl~xshould h,developed
0 8
0
01 02 03 MILLIMOLE COBALT PER 5 0 M L
04
in the pH r : q e of 7 . 5 to 9. The higher the pH of the solution, the slower will be the color development. Time. Naxinium color development Tvas reasonably rapid. It is recommended t h a t the solution stand for a t least 10 minutes before the measurement of absorbance. The rate of oxidation of the cobalt(I1) complex and its stability are greatly influenced by pH. Between pH 7 and 9, the ovidation occurred almost instantaneously; over pH 9, the speed of oxidation was considerably slon-er. Thirty minutes to one hour may be needed for complete oxidation a t pH above 10. Oxidant. Sodium perborate offers the advantage of preventing interference from iron. When perborate is dissolved in water, it liberates hydrogen peioxide. Because of low solubility of sodium perborate, it cannot supply enough hydrogen peroxide to oxidize the nitrilotriacetic acid of iron(II1) t o the violet-colored compound, An experiment confirmed that a 0.02% hydrogen peroxide solution failed to oxidize the nitrilotriacetic acid complex of iron(II1). Crystalline sodium perborate n as found t o be better than its saturated solution. One-tenth gram of sodium perborate tetrahydrate was enough for 0.1 mmole of cobalt. Stability and Beer's Law. The color of the nitrilotriacetic acid coinplex of cobalt(II1) was not stable in the presence of relatively large amounts of manganese(I1). The solutions containing 0.1 mmole of the violet cobalt(II1) complev and 0.1, 0.5, and 1.0 mmole of manganese faded after standing overnight. The one containing 0.1 mmole of manganese became very light violet; 0.5 mniole of manganese gave a very light
1
/
4 MILLIMOLf
05
Figure 3. Effect of concentration of phosphoric acid on calibration curves for cobalt
0
0 05
0.10
MILLIMOLE
0 15 COBALT
0 20
0 25
PER 5 0 M L .
Figure 4. Effect of NTA on calibration curves for cobalt VOL. 30, NO. 6, JUNE 1958
1037
pink color; 1.0 mmole of manganese became yellow without a trace of violet or pink color. The addition of 0.1 gram of sodium perborate did not redevelop the maximum color. The violet cobalt(II1) complex followed Beer's law at the range of 0.01 t o 0.4 mmole of cobalt per 50 ml. in phosphoric acid medium (see Figure 3) and also a t the range of 0.01 to 0.1 mmole of cobalt per 50 ml. in sodium bicarbonate medium (pH about 8 to 9). The concentrations over 0.1 mmole of cobalt in 50 ml. did not follow Beer's law (see Figure 4). Temperature. Upon boiling, t h e complex decomposes with t h e liberation of oxygen, and the yiolet cobalt(111) complex gradually changes t o a light pink color nhich is the color of nitrilotriacetic acid complex of cobalt(11). Specificity. S o other metals tested gave t h e yiolet color when perborate was added t o their nitrilotriacetic acid complexes. Kitrilotriacetic acid is used here not only as a colorforming agent, but also as a masking agent. The desirability of measuring the absorbance in acid or alkaline medium depends upon the other metals present in t h e mixture. Pribil and Svestka (9) found that nickel and manganese affected the oxidation of cobalt in the presence of (ethylenedinitri1o)tetraacetic acid. I n the present work, neither nickel nor manganese had significant influence on the oxidation of nitrilotriacetic acid complex of cobalt(I1) by perborate (see Tables I11 and IT). The qualitative test of the reaction of all common elements with nitrilotriacetic acid and hydrogen peroxide has been reported ( 3 ) . The determinations of cobalt in the presence of many possible interfering metals were made (Table IV). It may be concluded that the color reaction of nitrilotriacetic acid and perborate is highly selective for cobalt and that cobalt may be determined in the presence of a number of metals spectrophotometrically.
Table 111.
- --
THEORETICAL VALUE.5 I8 ML
I
2
3 4 ML O.04OM
Table IV. Determination of 0.1 Mmole of Cobalt in Presence of Other Metals
Metal Added
0.5M NTA, AbsorbMmole hI1. ance
Phosphoric Acid None Copper (I1j ZinciII) TroniIII) _ . . ~iciL~I1) Molybdenum(F'1j Tungsten(V1) Vanadium (V)
Medium, 560 hIp .., 2 0.390 0.1 2 0.387 0.2 2 0.390 0 2 2 0 395 0 2 0 390 0 2 2 0 390 0 2 2 0 390 0 1 2 0 387
Sulfuric Acid Medium, 560 M p 0.1 2 0.390 Titanium(1V) 0.1 2 0.388 Chromium(II1) 0.1 2 0,388 Uranium(V1) 0 1 2 0 388 Antimony(Vj Sodium Carbonate Medium, 580 NnnP - _.._
Bismuth (111) Mercury (I1j Zirconium(1V) Cadmium(I1) Lead (11) Titanium (IV)
...
0.2 0.2
0.2 0.2 0.2 0.2
4
M p
0 0 0 0
343
345 345 343 0 343 0 345 0 3-1.3
(50-ml. volume) Sodium Bicarbonate 1Iediuni Phosphoric LieidMedium (PH 8.7) XTA added, hbsorbance, Added. KTA added, Absorbance, 1Iniole mmole 560 m p a mmole 580 mpb Sone 1 0 0 398 2 0 0.338 0 01 1 0 0 398 2.0 0.338 0 02 1 0 0 397 2.0 0.338 0 05 1 5 0 398 2.0 0,338 0.1OC 2 0 0 397 2.0 0.338 0.50c 2 5 0 360 2.0 0.337 1 ooc 5 0 0 350 2.0 0.332 a Allowed to stand 60 minutes before absorbance readings were made. Allowed to stand 10 minutes before absorbance readings were made. Nitrilotriacetic arid precipitate caused by addition of phosphoric acid was centrifuged 1
Off.
1038
ANALYTICAL CHEMISTRY
8
9
1
Figure 5. Spectrophotometric titration of cobalt with hydrogen peroxide
Determination of 0.100 Mmole of Cobalt in Presence of Manganese
Mn( I1
5 6 7 HYDROGEN PEROXIDE
As the cobaltic complex is not extracted by organic solvents, the colored salt solution would be expected to interfere slightly with the absorbance measurement. Such interference may be eliminated by using a sample blank n i t h no addition of perborate. Titanium did not cause any interference because it did not show much intense coloration (yellow) in the presence of nitrilotriacetic acid a t p H 9. Vanadium nnd uranium gave an intense yellon. color as they were oxidized by perborate. The perborate was added after the addition of phosphoric acid to the blank; thus, in acid medium cobalt(I1) did not form a complex with nitrilotriacetic acid and could not be oxidized by perborate. On the other hand, vanadium and uranium were still oxidized by perborate in acid medium. For small amounts of vanadium, uranium, and titanium, the yellow color of their peroxy compounds does not affect appreciably the absorbance measurement a t 580 mp (in carbonate medium). Chromium(II1) formed a relatively light purple color with nitrilotriacetic acid upon standing. The blank composed of the sample solution and the reagents corrected the interference caused by the nitrilotriacetic acid complex of chromium(II1). Khen difficulty is encountered in a deeply colored sample, the cobalt may be determined after it is separated as potassium cobalt nitrite. The nitrite precipitate is easily dissolved by nitrilotriacetic acid. Determination of Peroxide. The formation of the violet-colored nitriloacetic acid complex of cobalt(II1) was caused by the peroxide oxidation; hence, this color reaction niay be used for the determination of hydrogen peroxide. Both sodium peroxide and sodium perborate contain hydrogen peiouide; the color reaction
m a y aIso be used in the determination of their peroxide content. Determination of Nitrilotriacetic Acid. T h e color reaction also offers a method of determining nitrilotriacetic acid. The color must be developed a t p H above 7 in t h e presence of excess cobalt and citrate. T h e cobalt which is not complexed b y nitrilotriacetic acid is complexed by citrate. The stable violet nitrilotriacetic acid complex or (ethylenedinitri1o)tetraacetic acid complex of cobalt(II1) may be determined spectrophotometrically in acid or alkaline medium. Color Reaction for Tartaric Acid. Tartaric acid reacts with cobalt(I1) t o give a bright red color, which changed t o yellow and finally to green upon addition of perborate a t 8.5 t o 10. This colored complex was decomposed after addition of phosphoric acid or other acid. Citrate gave no color change upon addition of perborate.
Determining nitrilotriacetic acid, (ethylenedinitri1o)tetraacetic acid, and tartaric acid by using the formation of cobaltic complexes will be reported in another paper ( 7 ) .
x. L., -4XAL. C H E M . 27, 1165-6 (~1955). (3) Cheng, K. L., Lott, P. F., Zbid., 28,462 (1956). (4) Hanker, J. S., Master, I., Mattison, L. E., Witten, B., Zbid., 29, 82 (2) Cheng,
(1957).
(5) Kdlthoff, I. M., Sandell, E. B.,
“Quantitative Inorganic Analysis”
ACKNOWLEDGMENT
The author is grateful t o the Alrose Chemical Co. for supplying diethylenetriaminepentascetic acid, N-hydroxyethylethylenediaminetriacetic acid, and diaminocyclohexanetetraacetic acid, to J. R. Geigy S. A. for providing ethyleneglycol-bis-(@-aminoethyl ether) - iV,N’tetraacetic acid, and to the Becco Chemical Division, Food hfachinery and Chemical Corp., for providing sodium perborate used in this investigation. LITERATURE CITED
(1) Bailar, J. C., Jr., “Chemistry,pf the Coordination Compounds, pp.
398-415, Reinhold, Kex York, 1956.
p. 600, Macmillan, Sew York,
1952. (6) Lott, P. F., Cheng, K. L., A s . 4 ~ . CHEW29, 1777 (1957). (7) Zbid., submitted. (8) Pribil, R., hlalicky, V., Collection Czechoslov. Chem. Communs. 14,
413 (1949). (9) Pribil, R., Svestka, L., Ibid., 14, 31 (1949). (10) Schwarzenbach, G., “Die Komplexometrische Titration,” pp. 87-9, Ferdinand Enke, Berlin, 1955. (11) Schwarzenbach, G., Heller, J., Helv. Chim. Acta 31, 1039 (1948). (12) Ibid., 34, 1889-1900 (1951). RECEIVEDfor review March 12, 1957. Accepted December 6, 1957. Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, March 1957.
Determination of Nitrogen Dioxide Gas-Solid Chromatography 5.A. GREENE’ and H. PUST Aerojet-General Corp., Azusa, Calif.
b Nitrogen dioxide i s rapidly converted to nitric oxide by the water absorbed on a Linde Molecular Sieve column. The resultant wave form i s indistinguishable from that of pure nitric oxide introduced to the column. Thus, by conversion to the low boiling nitric oxide, the complication of oxidation of organic partitioning liquids i s obviated.
B
nitrogen dioxide boils a t approxiniately room temperature, attempts were made t o separate i t on gas-liquid hydrocarbon columns. Sharp peaks followed by a long tailing peak seemed t o indicate coinplications involving oxidation of the colunin liquid. Use of column liquids such as fluorinated hydrocarbons and silicone oils gave essentially the same phenomena; long tailing peaks were observed, and the retention time of the leading edge mas not predictable. The wave forms did not yield data amenable for quantitative analysis. If nitrogen dioxide reacted a t room temperature with \vater, the following reactions would take place: 2S02 L HzO + HKOz HXO3 (1) Sitrous acid, being unstable, m-ould decompose : * Present address, Rocketdyne, 6633 Canoga Ave.. Canoga Park, Calif. ECAL-SE
+
3IIS02 -+ HXO3
+ 2x0 + H20
RESULTS AND DISCUSSION
(2)
liberating nitric oxide. The theoretical T-ield-Le., nitric oxide to nitrogen &oxide-is 1 to 3 . It rrould also be convenient if the water was adsorbed on a n adsorption column and t h e reaction and separation of mixtures took place on a single adsorption column. Nitric acid from the reaction would be strongly adsorbed on the column and would not be eluted. Because the reaction product would be nitric oxide (a low boiling liquid), a n adsorption column, such as Molecular Sieve, alumina, or silica gel, would be needed. The separation of nitric oxide by gas chromatography has been described (3). A Molecular Sieve column was chosen for evaluation because methane, another low boiling gas, and nitric oxide could be conveniently separated a t room temperature.
The retention volumes of some gases of the column were: 180 cc. for oxygen, 330 cc. for nitrogen, and 480 cc. for nitric oxide. The upper curve of Figure 1 shows the calibration obtained with pure nitric oxide. K h e n equal pressures of nitrogen dioxide are admitted t o the column, a wave form of the same
EXPERIMENTAL
The apparatus has been described ( 2 ) . Linde Molecular Sieve, 5 8 , was crushed to 20 to 40 mesh and packed into a 10-foot length of aluminum tubing, l/4 inch in outside diameter. The column was not activated, and 2 nil. of water was pipetted onto the anterior end of the adsorbent. Helium carrier gas-flow rate was 60 ml. per minute, and column temperature was 23” C
Figure 1. Calibrations obtained with nitric oxide and nitrogen dioxide 0
0
+
Nitric oxide Nitric oxide yield from nitrogen dioxidewater reaction Theoretical yield of nitric oxide from nitrogen dioxide-water reaction VOL. 30, NO. 6, JUNE 1958
1039
