I
E. J. KOVALl and M. S. PETERS University of Illinois, Urbana, Ill.
How Does Nitric Oxide AfFect.
..
Reactions of Aqueous Nitrogen Dioxide?
When water is the absorber, adding nitric oxide to nitrogen dioxide b Increases nitrous acid
b Removes nitric oxide
b Decreases nitric acid
b Changes only slightly the re-
b Decreases total acid
moval rate of nitrogen dioxide
NITROGEN
OXIDES are absorbed in aqueous solutions industrially to form nitric acid, nitrates, and nitrites. These oxides are composed of NOz, N204, NO, and small amounts of Nz03 and NzOs. The absorption process is carried out in large absorption towers where air is introduced between trays to oxidize the NO to NO?. The important effect of this oxidation step on yield has been recognized for many years, but only one study (7) is reported in the literature on the effect of N O on the kinetics involved in the aqueous absorption of mixtures containing NO, NOS, and Nz04. Three earlier studies (5, 77, 72) are available on the absorption of nitrogen oxides into either sodium hydroxide or strong nitric acid solutions. Very recently, two researchers ( 9 , 78) have suggested that nitric oxide influences markedly the rate of reaction of nitrogen dioxide with water even when N O is not present in the inlet gas stream. The study presented in this paper was initiated to obtain a clearer picture on
Present address, California Research Corp., La Habra, Calif.
the influence of NO on the production of nitric acid and on the kinetics and mechanisms involved in the absorption of nitrogen oxides into water. Mixtures of nitrogen oxides and nitrogen were absorbed in demineralized water in a long wetted-wall column a t room temperature and atmospheric pressure. Material balances based on total active nitrogen ranged from 98 to 10201,. The controlling mechanism is the rate of the chemical reactions, and reactions of Nz04 with HzO and the forward and reverse reactions of N203 with HzO must all be considered. When nitric oxide is added with eNO2 in the inlet gas stream, the nitrous acid production increases, the nitric acid production decreases, the removal rate of eNO9 changes only slightly, except a t very high NO/eN02 ratios, total acid production measured per mole of entering active nitrogen decreases, and nitric oxide disappears rather than appears as is the case without the addition of NO. Also, the hypothesis that nitrous acid concentration is negligibly small in a long wetted-wall column or in a bubbledispersing unit, with or without agitation, is not always correct.
Equipment and Procedure The wetted-wall column consisted of a borosilicate glass tube, 73 cm. long and 2.15 cm. in diameter. A 500-cc. glass bulb, which was sealed to the upper part of the column, acted as the water feed bulb. The lower section of the column penetrated vertically into a 300-cc. borosilicate glass bulb just beyond its middle plane. The gas mixture entered the bottom of the column through a 1.58cm. diameter borosilicate glass tube which overlapped the lower part of the column from the inside, providing a water seal. Demineralized water was fed to the column a t a steady measured rate from a constant-head tank. Gaseous nitrogen dioxide and nitric oxide (both Matheson Co., 99% purity) were mixed with nitrogen as the diluent. The gas mixture was sampled and entered the bottom of the column. Nitrogen flow rates were measured by a calibrated Venturi meter, while the nitric oxide flow rate was measured with a calibrated capillary tube. A calibrated capillary tube was also used to set the approximate flow rate of nitrogen dioxide. The exact amount of nitrogen dioxide was determined from the nitrogm and nitric oxide flow rates and a chemical analysis of the entering gas mixture. After steady state was attained, samples of the inlet and outlet gas and the exit liquid were taken. Evacuated Gaillard bulbs were used for taking the gas samples. For the inlet gas sample, the bulb contained hydrogen peroxide. For the outlet gas sample, two bulbs were necessary; one contained hydrogen peroxide, the other contained potassium permanganate and sulfuric acid. The amount of gas was found by weighing. The hydrogen peroxide converted the nitrogen oxides to nitric acid which was VOL. 52. NO. 12
DECEMBER 1960
101 1
0 501
/ 00 0
I
I
I
1
TMAN -MOLES NO, SMOLES NOzI t MOLES NaOa,
08
16 24 MOLE ?6 e NO2 I
32
40
Figure 1. Removal efficiency decreases as eN02 in the entering gases i s reduced, or as the ratio of NO to total moles of active nitrogen i s increased
I 0
TOTAL ACID HNO. F R A C T I O N
I? I
~
-MAN,
Figure 2. Adding NO has a marked effect on acid-product distribution. Over-aII effect at 2 mole of eNO2
7 0
then titrated. Since the amount of S O in the inlet gas sample was known, the nitrogen dioxide was obtained by difference. In the outlet-gas sample, the permanganate-sulfuric acid bulb was necessary to give the SO-SO? split. The amount of K M n 0 4 consumed in the reaction with the nitrogen oxides was determined by titrating the excess K M n 0 4 with sodium thiosulfate. Relative amounts of nitrogen dioxide, dinitrogen tetroxide. and dinitrogen trioxide were calculated from the known equilibrium constants for the dimerization of SO2 (22) and the formation of N 2 0 3 from N O and NO2 ( 7 3 , 22). The liquid samples were collected under a sodium hydroxide solution. The resultant sodium nitrate, sodium nitrite, and excess sodium hydroxide mixture was back-titrated with standard HC1 to determine the original total acid content. Then the sodium nitrite was determined by the Saltzman method (20) to determine the nitrous-nitric acid split. Material balances were made by comparing the moles of active nitrogen entering the system as K O or eSOa (SO7 f 2 N 2 0 4 ) with the moles of active nitrogen leaving the system as NO, eSOz, HS02,and “03. The amount of dinitrogen trioxide was negligible at all times. All runs were made a t a constant water flow rate of 270 ml. per minute and a constant nitrogen flow rate ol 0.032 pound per minute (0.46 cubic foot). The entering eNO2 concentration was maintained approximately constant a t one of four levels--0.6. 1, 2, and 4 mole yowhile the concentration of N O was being varied from 0 to as high as 8.2 mole Room temperature averaged about
88OF.
Removal efficiency, defined as the percentage of nitrogen oxides removed from the gases, decreases with reduction in eNO2 content of the gases and also decreases when the ratio of NO to the total moles of active nitrogen (TMAN)
3. Because the ratio of to H N 0 2 in the acid product changes, the stoichiometry of the reactions involved in the absorption process changes as the ratio of N O t o e N O l increases Figure
“03
1012
Previous investigators dealing with aqueous absorption of nitrogen oxides ( 9 , 17-79, 23) have applied the following rate equation to give good correlation of their experimental data when little or no N O was present in the inlet gas stream.
The basic assumptions involved in the theoretical development of Equation 1 are: Rate of eKOz removal is controlled by the rate of the reaction between water and 1 - 9 0 4 and NO?; the chemical reactions occur under irreversible conditions ; constant temperature and constant gas rate prevail; instantaneous equilibrium exists between NO2 and NzOi; no oxidation of S O occurs ; and the nitrous acid concentration is negligible or unchanging. Under these conditions for a steady-state flow reactor, Equation 1 can be integrated to give (19) : (1/Pix,o,)1’2 = (1 / P o h , ~ 4 ) 1-’ 22 K p In
+ At
(PoX2OpI~IN*o4)
I /A I
I
I
I
I
I
+OO,.li__l 0
Figure 4. As NO i s added, the overall removal rate of eNOz first decreases and then increases
Interpretation of Res u Its
YG.
Experimental Results
NO, eN021
entering the system is increased (Figure 1). The decline in removal efficiency? as the entering concentration of eNO? is reduced: is characteristic for this system when the rate of nitrogen dioxide removal is controlled by the rate of the chemical reactions ( 79). Because NO does not react directly with water, it is more difficult to remove; thus, removal efficiency is adversely affected by KO. Although the presence of NO in the reacting gases can have a dramatic effect on the H N 0 3 - H N 0 2split (Figures 2 and 3), the over-all rate of removal of eSO2 is only slightly affected (Figure 4). The average rate of disappearance of eNO? declines initially as Ti0 is increased, but then increases to a level higher than that obtainable with no Ti0 in the inlet gas stream. Delta quantities are a reasonable measure of the average rate, because gas and liquid rates are constant, thus keeping contact time constant. NO is first produced but then disappears in large quantities as the SO/eN02 ratio is increased (Figure 5). This disappearance is brought about by the sharp increase in the formation of nitrous acid.
08
16
24
32
NO~ANO~~
1 40 4 4
Figure 5 . Because of the large relative amounts of HNOz formed at high ratios of NO to eNO2, NO i s first produced, but then i t i s consumed as the ratio increases
INDUSTRIAL AND ENGINEERING CHEMISTRY
(2)
where Pih20pis partial pressure of S 2 0 4 in exit gas; PoT204 is its partial pressure in entering gas; K , is equilibrium constant for dimerization of NOg, Px,o,/ (PNo2)2.atm.-l; A is a constant at any given temperature ; and t is contact time. The fit of the experimental data obtained in this study with Equation 1 or 2 was unsatisfactory which is expected because of the high concentrations of NO. There was a decided trend of the rate constant k with the concentration of
N I T R O G E N DIOXIDE R E A C T I O N S eNO2, and an effect of NO concentration, although small, was apparent. Because some investigators (8, 70) have proposed that gas-phase diffusion is the controlling step in this type of operation, attempts were made to correlate the data by using the diffusion-controlling equation previously proposed (8, 75). No agreement was found. When a wetted-wall column is used with water as an absorbent, the following rate equation is applicable a t high concentrations of NO (7) :
*
Thus some form of mechanism other than diffusion or direct reaction between dinitrogen tetroxide or nitrogen dioxide and water must be involved to account for the change in the rate of removal of eNOz as a function of the NO/eNOp ratio and for the marked change in the HNOs-HNOz split. Clearly, the hypothesis upon which Equations 1 and 3 are based, that the amount of nitrous acid is negligible or is unchanging, is not valid (Figure 2). With these facts in mind, a mechanism is postulated which appears to fit the experimental data and to agree well with all prior observations. The absorption process is broken down into the following steps :
The second term was small and could be considered a correction term. Equation 3 would be satisfactory to explain the results observed in this work when the NO/eNOz ratio was greater than about 3, but it would not explain the apparent decrease in the rate of removal of eNO:! a t lower NO/eN02 ratios nor would it suggest the observed acid product distribution. Attempts to explain the results by including the effect of H N O ? diffusion were unsuccessful
+
Nz04(g) = NzOd(1) = N O f NOa(physical equilibrium) ( a ) ~
~) =0N ~ ~ o ~ (=(~ ) NO$ NOzNz04 Hz0 HNOz
+ +
1 reactions
I
involved ( b ) in determining "03 N ~ O H ~ z o = 2HN02 rate
+
+ 2HN02 = Nz08 + HzO N O + NO2 = Nz03 (equilibrium) 2N02
Nz04
(e)
(f) (equilibrium) ( g )
The rate of disappearance of eNOz then is governed by the rate of Reaction c and the difference in rates of Reactions d and e, because the latter two reactions are not in equilibrium for the reasons cited previously. This leads to Equation 4 after including the appropriate Henry's Law and equilibrium constants for steps a, b, f, and g :
-
d(eNoz)g
dt
+
= ,+,(NOz)g2
kd(NO)g(NO~)g - ke("0z)i2
(4)
The only difference between Equation 4 and Equation l is a n allowance for a significant and varying amount of "02. The production rate of " 0 2 did not equal its disappearance rate, except a t certain operating conditions. Equation 1 was derived by assuming that only a small amount of HNOz was present a t equilibrium. This idea was based on the principle that HNOz was in equilibrium with NO, Nz04, and water (7-3). However, this principle applied to a batch system and time was measured in minutes. I n the present work, time was measured in seconds or fractions thereof.
Experimental and Calculated Data for W e t t e d - W a l l Columns Agree W e l l
(Total press., 1 atm.; water rate, 270 cc./min.; Nz rate, about 0.032 lb./min.) Enterine Mole Mole Ratio, Entering Ratio, Total HNOz in NO/ id/ i id, NO, Mole yo TMAN TMAN Mole yo I
Entering
Av.
Temp., O F. 82 88 91 92 86 84 83 93 92 88 85 88 90 87 86 90 89 92 88 87 87 87 90 91 92 89 88 88 91 86 85 87 86 88
Matl. Bal., 70 100.2 95.3 100.4 100.0 100.1 101.3
100.0 102.6 97.2 100.4 100.3 99.9 97.8 101.5 99.3 99.7 98.0 100.0 101.2 99.8 100.9 101.2 98.4 95.8 100.9 98.8 99.7 101.3 100.8 100.7
eNO2,
Mole yo 4.04 4.25 4.16 4.24 4.44 4.40 4.46 4.08 3.92 3.79 3.87 2.05 1.96 2.03 2.26 2.20 1.78 1.82 1.81 1.64 1.15 1.14 1.14 1.14 1.14 1.06 0.62 0.64 0.61 0.63 0.66 0.63 0.63 0.64
0 0 0
0 0 0
0.74 1.78 3.21 3.60 5.17 6.22 6.28 8.20 0 0.33 0.85 1.67 1.70 2.37 4.63 5.68 6.45
0.16 0.32 0.46 0.48 0.58 0.64 0.65 0.70 0 0.15 0.31 0.45 0.45 0.59 0.74 0.77 0.81
0
0
0.20 0.37 0.71 1.70 5.80
0.15 0.25 0.38 0.61 0.85 0 0.14 0.26 0.40 0.42 0.56 0.73 0.81
0
0.10 0.21 0.41 0.46 0.76 1.76 2.63
0.46 0.44 0.40 0.33 0.31 0.29 0.26 0.23 0.22 0.20 0.19 0.33 0.28 0.24 0.20 0.19 0.17 0.13 0.13 0.12 0.26 0.22 0.16 0.16 0.13 0.08 0.19 0.18 0.14 0.14 0.14 0.11 0.08 0.07
45.7 45.2 44.7 51.0 59.4 65.5 70.3 70.0 72.0 77.6 77.7 44.4 48.1 52.2 66.5 66.8 69.5 75.3 84.5 84.7 44.6 49.1 50.4 57.0 68.6 91.0 48.7 49.1 52.1 54.3 58.1 66.9 77.6 81.2
Basis, per Min.
Outlet eNOn, Lb.-Moles X 10' Calcd. Observed Diff., Yo 280.0 289.7 299.8 317.0 346.0 318.5 364.5 292.3 279.5 301.5 299.0 156.0 152.9 158.7 192.8 184.5 146.5 131.5 141.5 121.8 98.6 99.9 108.1 101.1 99.8 95.6 58.4 57.2 59.4 57.0 59.9 59.3 60.4 59.0
258.0 261.0 282.0 283.0 321.0
- 7.9 - 6.5
333.0
- 8.7
325.0 339.0 150.5 143.5 152.7 188.0 182.5 135.1 137.0 146.0 96.8 97.1 113.8 96.4 98.2 97.4 34.5 57.9 61.9 52.8 54.4 53.4 45.3 72.5
-
5.9 -10.7 - 7.2
7.8 13.4
- 3.5 - 6.1 - 3.8
-
2.5 1.0 - 7.5
- 3.2 20.0 - 1.8 - 2.8 5.3 - 4.7 - 1.6 1.9 -41.0 1.2 4.2 - 7.4 - 9.1 - 9.9 -25.2 22.9
+
VOL. 5 2 ,
Outlet NO, Lb.-?doles X lo7 Calod. Observed Diff., 70 7.9 9.3 9.2 84.0 195.0 352.5 397.5 577.6 709.5 726.5 965.0 4.4 39.3 96.6 183.8 182.9 262.9 527.5 651.7 750.4 1.9 23.6 43.2 81.0 191.4 700.0 0.4 11.5 24.4 47.5 52.0 86.1 196.9 309.0
NO. 12
e
32.0 38.0 28.5 118.0 222.0
30.5 30.9 21.0 40.5 13.8
429.0
7.8
705.0 925.0 9.9 50.0 104.0 188.0 185.0 274.3
-
3.0 4.2 12.5 27.2 7.7 2.3 1.1 4.3
656.3 725.0 3.6 27.9 37.8 85.3 192.5 700.0 24.1 10.6 22.0
0.7 3.5 8.5 10.2 -12.5 5.3 0.6
51.5 57.1 92.0 212.0 295.8
8.4 9.8 6.8
DECEMBER 1960
-
0
-
8.1
- 10.0
7.7 - 4.3
1013
Steps a and b represent the physical equilibrium that exists between gaseous N204 and aqueous N204 and between gaseous NzO:, and N z 0 3in solution. The ionization steps are the same as those previously proposed (G, 74, 76) and imply that N204 is the reactive species. Reaction c is assumed irreversible as the Hh703concentration is below 0.1 weight per cent. Prior research ( 7 7) has shown that the reverse reaction is important only at H A T 0 3 concentrations above 20 weight per cent. Reactions d and e are both important and both are faster than Reaction c. This fact was observed in experiments in which mixtures of N O and S O ? were absorbed in sodium h>-droxide solution (4, 72). The direction of the over-all effect of Reactions d and e is governed by the concentrations of H S O z and N203 imposed by the operating conditions. Reactions f and g are both essentially instantaneous reactions with well established equilibrium data available (22). Consider first what occurs lvhen eNOz is added alone. The mixture of NO2 and NzOl diffuses quickly to the liquid film in Jvhich N 2 0 ~reacts with water to form H N O I and HNO?. Some of this HNOz decomposes to NO, SOZ: and HzO. A fraction of this released KOz reacts to make H N 0 3 and “ 0 2 , after first polymerizing to KgO,. Thus, HATOz cannot possibly disappear completely unless the rate of reaction of N z 0 4with water is infinitely fast. This sequence of events in the decomposition The of HNOz is well established (7-3). amount of the NO2 formed from the decomposition of “ 0 2 that reacts to make H N 0 3 and HNOz depends on the existing concentration level of eNOz. Other molecules of NO1 \vi11 react with N O and H20 to reform “ 0 2 . Since HNOz is being generated constantly by the independent reaction of N204 with water, its rate of decomposition will be accelerated and, as a result, there is a net decomposition of Hn’02. Nitric oxide appears in the exit gas (Figure 5)>and the fraction of the acid product which is H N 0 2 is less than 50% (Figure 2). From the proposed mechanism, the rate of disappearance of N O is controlled by Reactions d and e. This leads to Equation 5.
of disappearance of eNO2 decreases as N O content is increased (Figure 4), and the production of HNOB begins its steady decline (Figure 2). This decrease in rate of eSO2 disappearance comes about because of the stoichiometry involved in Reactions c and d. One mole of e S O 2 is required to make 1 mole of acid via Reaction c. Only l / z mole of e S O z is required to make HNOZ via Reaction d. As a consequence, where HNOz production is first building up a t the expense of H N O I production, the moles of eNOz required actually decline (Figure 3 ) . Sitric acid production steadily declines because the formation of HKOa from NO, N02, and water takes place much more rapidly. As seen from Reaction c, the rate of production of HNO:, depends on the second power of 1 \ 0 2 concentration and thus is sharply sensitive to any decline in the concentration of 3 - 0 2 . Nitrous acid production, on the other hand, depends on the first power of both N O and NO, concentrations, so that, a t a relatively fixed concentration of NOz, the rate of production will increase as the N O concentration is increased (Figure 2 ) . As the N O concentration is raised still further, Mhile holding the inlet concentration of eNO2 constant, more HXOz is formed than decomposes. Nitric oxide disappears and more “ 0 2 than “ 0 3 appears in the acid products (Figures 2 and 5). The rate of disappearance of NO2 begins to increase because of the large amount that is disappearing in the formation of “02. Indeed when the NO/eN02 ratio becomes sufficiently large (about 3), the removal rate of eNO2 exceeds that obtainable when no S O is present in the inlet gas stream (Figure 3). This latter effect agrees with that given by Equation 3. Like Equation 1, Equation 3 is a special case of the more general rate expression for removing eNOZ, Equation 4. Note that a t the high ratios of NOjeNOn, almost all the acid product is H K O Z . Sote also (Figure I ) , that the moles of acid product formed per mole of entering active nitrogen steadily decreases with an increase in N O because N O does not react directly with HzO. Because the concentration of H S 0 2 is far from negligible, the over-all stoichiometry cannot be represented by
Equation 5 shows that NO should disappear when the N O concentration is increased sufficiently to cause a n excess of “ 0 2 over and above that produced from the N2Or-water reaction. Thus, the production or disappearance of NO is a measure of the direction in which the HNO2 reaction is moving. When only a small amount of NO is present in the initial gas stream, the rate
Equation 6 has been used by other investigators (78, 79) to relate the instantaneous value of S O with eSO? for an integral reactor system when the outlet gas composition could not be determined directly. Figure 5 and Equation 6 show that the relationship between NO and eSO2 is not that simple, because Equation 6 does not take into account the presence of significant amounts of HNOa.
10 14
INDUSTRIAL AND ENGINEERING CHEMISTRY
The actual stoichiometry is best described by Equation 7
+ $)
+ (2 + ;) 2HN03 f 2 1 + ( 2)
(2
=
N 2 0 4
” 0 2
-
e2 ’KO (7)
@ refers to the molar excess of HSOz over
This is the case Lvhen HNO2 is formed in greater quantities than are decomposed. If more H N O ? is decomposed than is formed, 0 is negative, but Equation 7 Lvould still apply. Snoeck (27j has developed a similar equation for the case irhen p is negative. The experimental data appear to satisfy the proposed mechanism. In addition, the stoichiometry suggesred by the mechanism was checked by coinparing the outlet gas composition as calculated from Equa ion 7 with the observed values. Agreement was well within experimental error and in most cases was within -;t 5%,
“03.
literature Cited
(1) Abel, E., Schmid, H., Z . Physik. Chern. 132, 56, 64 (1928). ( 2 ) Zbid., 134,279 (1928). (3) Zbid., 135,419, 430 (1928). (4) Atroshchenko, V. I., J . L4fiplied Chetn. ( U . 5‘. S.R.) 12, 167 (1939). (5) Bolshakoff, P. E., S.M. thesis, M.I.T., 1934. (6) Carberry, J. J., Chem. Eng. Scz. 189 (Feb. 1959). (7) Caudle, P. G., Denbigh, K. G., T r a n s . Faraday SOC.49, 39 (1953). (8) Chambers, F. S., Sherwood, T. K.: IND.ENG.CHEM.29, 1415 (1937). (9) Chen, J., Ph.D. thesis, University of Illinois; 1959. (10) Dekker, W. A,, Snoeck, E., Kramcrs, H., Chern. Eng. Sci. 61, (.August 1959). (11) Denbigh, K. G., Prince, A . .J.. J . Chem. SOC.1947,p. 790. (12) Eagleton, L. C., Langer, R. M., Pigford, T. H., S.M. thesis, M.I.T., 1948. (13) Gray, P., “Chemistry of Dinitrogen Tetroxide,” Roy. Inst. Chem., London, 1958.
(14) Gray, P., Yoffe, .A. D., Qziari. Recs. (London) 9, 362 (1955). (15) Klein, J. E., M.S. thesis, University of Illinois, 1954. 116) Miller, D. J.. Watson. D. ,J., J. Chem. ’ S o t . 1957, p. 1369. (17) Peters, M. S., Tech. Rept. 14, Eng. ExDt. Sta.. Universitv of Illinois. 1955. (18) Peters, M. s., koval, E. J., IND. E ~ GCHEM. . 5 5 , 577 (1959). (19) Peters, M. S., Ross, C . P., Klein, J. E., A.I.Ch.E. Journal 1, 105 (1955). (20) Saltzman, B. E., Anal. Chern. 25, 1949 (1954). (21) Snoeck, E., “Afstud. Verslag Fys. Techn.,” T. H. Delft, Holland, 1958. (22) Verhoek, F. H., Daniels, F.. J . Am. Chem. SOC.53, 1250 (1931). (23) Wendel, M. M., Pigford, R . I’., A.I.Ch.E. Journal 4, 249 (1958). RECEIVED for review April 4, 1960 ACCEPTED July 25, 1760 Work supported by a fellowship from E. I. du Pont de Nemours 8r Co., Inc., and a grant from the National Science Foundation.
