L i t e r a t u r e Cited Alexander, P., Stacey, K. A.. Ann. N.Y. Acad. Sci., 68, 1225 (1958). Bronsted, J. N., Kilpatrick, M., Kilpatrick, M., J. Am. Chem. Soc., 51, 428 (1929). FraenkeCConrat, H. L., J. Biol. Chem., 154, 227 (1944). Gilbert, G. L., et at., Appl. Microbial., 12, 496 (1964). Lichtenstein, H. J., Twigg, G. H., Trans. Faraday Soc., 44, 905 (1948). McCarthy, T. E., Sullivan, M. X., J. Bid. Chem., 141, 871 (1941). Marletta. J., Stumbo. C. R.. J. food. Sci.. 35. 627 (19701. Phillips, C. R , Bact. Rev., 16, 135 (1952). I
,
Phillips, C. R., Kaye, S., Am. J. Hyg., 50, 270 (1949). Weast, R. C.. Ed., "Handbook of Chemistry and Physics", 49th ed,pp D-78, E-37, Chemical Rubber Co., Cleveland, Ohio, 1968. Wheeler, G. p., CancerRes., 22, 651 (1962). Windmueller. H. G., Ackerman, C. J., Engel, R. W., J. Biol. Chem., 234, 895 (1959).
Receiued for review March 12, 1976 Accepted August 4,1976
,
Absorption of Nitrogen Dioxide into Water, Sulfuric Acid, Sodium Hydroxide, and Alkaline Sodium Sulfite Aqueous Solutions Yohji Kameoka and Robert L. Pigford'' Department of Chemical Engineering, University of California, Berkeley, California 94720
Nitrogen dioxide, NO2, was absorbed into water and aqueous solutions using a wetted sphere absorber. The results showed that N2O4, one of the two forms in the gaseous equilibrium mixture, dissolved and reacted more rapidly with water and the solutions than did NO2. The reaction in solution was first order with respect to dissolved N2O4. Hydrogen and hydroxyl ions had little effect on the absorption rate but absorption into 0.1 M alkaline sodium sulfite solution was about 2.5 times faster than in water.
Introduction Absorption of nitrogen oxides into water and aqueous solutions has been studied for a long time due to its importance in the industrial manufacture of nitric acid. Recently, there has been another interest in the phenomenon: air pollution control for stationary combustion facilities, because nitrogen oxides have been recognized to be one of the major pollutants in the atmosphere around us. In this case the concentration of nitrogen oxides in the gaseous effluent of a power plant is generally very low-several hundred ppmv (parts per million, volumetric) or less and most of the oxidized nitrogen is nitric oxide (NO),which is slightly soluble in water. For the removal of NO from flue gases from stationary sources, one can imagine a process in which nitric oxide is oxidized to the more soluble equilibrium mixture of nitrogen dioxide and nitrogen tetroxide and then is absorbed into water or an aqueous solution. In this paper, the rates of absorption of nitrogen dioxide into water, 0.09 N sulfuric acid, 0.2 N sodium hydroxide, and 0.042-0.153 M alkaline sodium sulfite solutions were measured a t low gas-phase concentrations of nitrogen peroxide, using a wetted-sphere absorber. Absorption into sodium sulfite solutions was studied to simulate simultaneous removal of sulfur dioxide and nitrogen peroxide by alkali scrubbing, which might be used to treat flue gases from stationary sources. Theory The following chemical reactions occur during absorption of nitrogen dioxide into water: 2NO2 = N204 (gas-phase equilibrium)
(1)
I Department of Chemical Engineering, University of Delaware, Newark, Del. 19711.
+
2N02 (or N204) H20 = HNOz
+ HNO:3 (reaction in solution) (2)
+ NO:;- (fast reaction in solution) 4HN02 = 2 N 0 + N204 + 2H20 N 2 0 4= NO+
(slow reaction in solution)
(3a) (3b)
Decomposition of nitrous acid does not take place rapidly if its concentration is very low. Previous studies (Caudle and Denbigh (1953), Chilton and Knell (1972), Corriveau and Pigford (1971), Denbigh and Prince (1947), Kramers et al. (1961),Peters et al. (1955),Wendel (1956, Wendel and Pigford (1958)) indicated that the rate-determining step is reactioc 2, the slow hydrolysis of nitrogen tetroxide. If reaction 2 is considered to be first order with respect to nitrogen tetroxide, solution of the unsteady-state diffusion equation in the liquid phase yields (4)
where 6 is the total rate of absorption in moles of N204 per unit time for a spherical surface of radius R. Cpi is the interfacial concentration of dissolved but unreacted N204. The time of exposure of the liquid surface to the gas can be calculated from the interfacial velocity of the fluid on the sphere. The result can be expressed by the two equations, t = 2.56R/Ui~and
is the equatorial surface velocity of the liquid layer a t the interface. If k t >> 1 in eq 4 the rate of absorption is nearly independent o f t and, therefore, of the speed of the fluid as it flows over the sphere. Equation 4 must be applied to each sphere. The total rate is found by multiplying by the number Ind. Eng. Chem., Fundam., Vol. 16, No. 1, 1977
163
Gas Out
Five-Sphere Absorber
Constant - head
Gas Out Cylinder Rubber Stopper
m
Lucite Cylinder
Liquid
Nazzle
Spacers Spacer Spheres
Tii
L J N2 Cylinder
Reservoir Water Bath
Water Bath
Pump
Figure 2. Flow diagram of the apparatus.
Gas in
k!jp!!i/
-Liquid Out
Figure 1. The wetted-sphere absorber. Each of the five stainless steel spheres was 3.81 crn in diameter. Each of the small cylindrical spacers between the spheres was 0.40 crn high and 0.48 crn 0.d.
of spheres in series. Since liquid mixing a t the junction between adjacent spheres is very rapid (Corriveau, 1971), t in eq 4 should be the value for a single sphere. Experimental Section
The five-sphere absorber, shown in Figure 1, consisted of five 1.5-in. diameter stainless steel spheres supported by a stainless steel rod and spacers in a 6-in. diameter Lucite cylinder, which provided adequate gas and liquid flow around the spheres. The flow diagram of the apparatus is shown in Figure 2. Nitrogen dioxide was used which had a minimum purity of 99.5 mol %. The storage cylinder was heated to about 30 "C to keep the nitrogen dioxide a t constant pressure (about 22 lb/in.2 absolute). After regulating the pressure, the flow rate of nitrogen dioxide was controlled by the needle valve and was measred by the rotameter. The nitrogen dioxide was then diluted by carrier nitrogen gas. To calibrate the nitrogen dioxide rotameter, a freezing-flask method was used. The unknown flow of nitrogen dioxide and a known flow of carrier nitrogen were collected in a flask, the temperature and pressure of which were measured. Then this flask was immersed in a dry ice-acetone mixture to solidify the nitrogen dioxide. After measuring the temperature and the new pressure and assuming ideal gas behavior, the flow rate of nitrogen dioxide could be calculated. The absorption rate was measured by acid-base or oxidation-reduction titrations. For absorption into water, the rate was measured by the change of the acidity between the absorber inlet and outlet liquids using acid-base titration. When nitrogen dioxide was absorbed into 0.09 N sulfuric acid or 0.2 N sodium hydroxide, the change of acidity or alkalinity of the absorbents was small compared to the original liquid concentration. Therefore, oxidation-reduction titration for nitrite ions rather than acid-base titration for both nitrate and nitrite ions was used to measure the rate of absorption of nitrogen dioxide. Based on the known rate of reaction 3b, it was estimated that nitrous acid, formed according to eq 2, did not decompose within the experimental range of liquid exposure times. When nitrogen dioxide was absorbed into alkaline sodium sulfite solutions 3% hydrogen peroxide was added to the samples to oxidize both sulfite and nitrite, after which the 164
Ind. Eng. Chem., Fundam., Vol. 16, No. 1, 1977
difference of alkalinity of the samples before and after the absorption was measured by acid-base titration. The inlet gas nitrogen dioxide concentration was obtained from the readings of the rotameters, while the outlet gas concentration was calculated from the material balance on nitrogen oxides around the absorber. The flow rate of nitrogen dioxide was adjusted to keep the nitrogen dioxide concentration of the inlet gas less than 2%. This avoided undesired side effects, such as fog formation, which might be encountered if higher nitrogen dioxide concentration had been used. No fog was observed. Based on previous determinations of the rate coefficient for reaction 2 in solution and using the experimentally selected liquid flow rate of 1.84 cm3-s,the product k t was equal to (194 s-l) (0.35 s) = 69. Thus, a t most the second term in the parenthetical expression of eq 4 was only about 0.7% of the first term. Under these conditions the liquid flow rate was not expected to affect the rate of absorption appreciably. The carrier nitrogen gas flow rate was fixed a t 6.48 X lop3 g-mol/s. All experiments were carried out a t 25 "C and atmospheric pressure. Gas-Phase Mass Transfer and Interfacial Partial Pressures
This work was concerned primarily with the liquid-phase mass transfer and chemical reactions of nitrogen dioxide. T o analyze the data, therefore, the effects of gas-phase mass transfer resistance had to be eliminated. Steady-state diffusion could be assumed in the gas phase. The equations of gas-phase mass transfer for nitrogen dioxide (1)and nitrogen tetroxide (2) through nitrogen are, respectively (6)
(7) where 1is the fractional distance through the gas film. Assuming chemical equilibrium between nitrogen dioxide and nitrogen tetroxide anywhere in the gas film, we have
From eq 6,7, and 8, and using the boundary conditions p1 = P lh at 3 = 0
Table 11. Comparison of Values of Hd\/kD
Table I. Absorption of Nitrogen Dioxide into Water
H 4 Z , g-mol of N:!O.+/cm'?-s-atm x lo5 20°C 25°C 30°C 35°C 1 0.992 2 0.992
3 0.993 4 0.993 5 0.995 6 0.995 7 0.995 8 0.995 9 0.985 10 0.985 11 0.985 12 0.985 13 0.985 14 0.985 15 0.985 16 0.985 17 0.985 18 0.985 19 0.985
0.0180 0.0136 0.0155 0.0195 0.0136 0.0172 0.0101 0.0119 0.0194 0.0123 0.0150 0.0082 0.0169 0.0100 0.0118 0.0144 0.01 10 0.0127 0.0189
0.0145 0.0112 0.0129 0.0161 0.0115 0.0139 0.0086 0.0100 0.0159 0.0103 0.0124 0.0071 0.0139 0.0087 0.0099 0.0119 0.0094 0.0107 0.0155
0.0162 0.0124 0.0142 0.0178 0.0125 0.0155 0.0093 0.0110 0.0176 0.0113 0.0137 0.0077 0.0154 0.0094 0.0109 0.0132 0.0102 0.0117 0.0172
0.986 0.622 0.846 1.29 0.683 0.917 0.390 0.528 1.21
0.550 0.749 0.276 0.951 0.421 0.515 0.707 0.465 0.586 1.16
Caudle and Denbigh (1953) Dekker et al. (1959) Wendel and Pigford (1958) Kramers et al. (1961) Corriveau and Pigford (1971) This work
0.763 0.478 0.543 0.757 0.429 0.701 0.283 0.356 0.797 0.390 0.557 0.200 0.639 0.240 0.358 0.513 0.314 0.409 0.765
11.0
25.0
11.0 5.8
10.0
7.7
8.9 5.71 6.85
Table 111. Absorption of Nitrogen Peroxide into Sulfuric Acid ( C H ~ S O=~0.0907 N)
No.
20 21 22 23
24
PI I l l n ,
P 1 t,'Jut,
p 1hdV,
atm
atm
atm
0.986 0.986 0.986 0.986 0.986
0.0150 0.0100 0.0169 0.0180 0.0118
0.0126 0.0084 0.0140 0.0148
0.0138 0.0092 0.0155
-
0.0164 0.0100 0.0109
N'
x 103, atm
X lo', g-moll cm'-s
0.799 0.365 0.985 1.08 0.526
0.519 0.301 0.622 0.696 0.349
Pni
P, atm
Table IV. Absorption of Nitrogen Peroxide into Sodium Hydroxide ( C N ~ O H= 0.207 N) pnl
No.
25 26 27 28
29 Figure 3. Measured rate of absorption of nitrogen tetroxide into water vs. interfacial partial pressure of nitrogen dioxide or nitrogen tetroxide. where
N
= NI
+ 2N2 g-mol NJ+/cm2-s
(10)
A=v'l+4KP
(11)
Using eq 9 and 8, p I 1and p2i could be calculated. Verhoek and Daniels (1931) gave the following equation for the equilibrium constant in eq 8: 1
= 0.1426 - 0.7588C2, atm a t 298.2 K K
(12)
k(;1 and kc:! were estimated by Schmidt number correction from the gas-phase mass transfer coefficient of ammonia through nitrogen, which was measured using the five-sphere absorber a t the same flow conditions as in the absorption experiments. These values were hc,l = 2.70 X lov5 and k ( : 2 = 2.52 X g-mol/cm2-s-atm,respectively. Results a n d Discussion The experimental data for the rate of absorption of tetravalent nitrogen into water a t 25 "C are shown in Table I. Recalling eq 4 and assuming Henry's law a t the gas-liquid interface, the rate of the simultaneous absorption and the quick first-order chemical reaction in a wetted-sphere absorber is $/4aR2 = Np = H m p 2 ,
(13)
p, atm
P 1 bin,
0.989 0.989 0.989 0.989 0.989
0.0204 0.0122 0.0104 0.0181 0.0166
atm
P II,I'U',
Pli,"',
atm
atm
0.0164 0.0184 0.0101 0.0111 0.0088 0.0096 0.0147 0.0164 0.0135 0.0150
x
lo,',
Nn X
lo',
atm
g-mol/ cm%
1.26 0.521 0.405 1.03 0.878
0.900 0.396 0.307 0.753 0.656
According to this equation, if nitrogen tetroxide were absorbed much more rapidly than nitrogen dioxide the absorption rate of tetravalent nitrogen atoms would be proportional to the interfacial partial pressure of nitrogen tetroxide or to the square of the interfacial partial pressure of nitrogen dioxide. Figure 3 shows the relations between the rate of absorption, expressed as moles of nitrogen tetroxide, and the interfacial partial pressures of nitrogen dioxide and nitrogen tetroxide. Most of the data agree with eq 13. I t is concluded, therefore, that the reaction with water of nitrogen tetroxide rather than nitrogen dioxide governs the rate of absorption of tetravalent nitrogen. A least-squares method gave the slope of the straight line which passed through the origin in Figure 3 as 6.85 X lop5 g-mol of N ~ O ~ / c r n ~ - s - athis t m ;is equal to H Several workers have also measured this product of physical constants. Table I1 shows a comparison of the values of H As can be seen in this table, the value obtained by this work was 17% lower than the interpolated value from Kramers et al. (1961), probably the best value available in the literature. Tables I11 and IV show the data for the absorption of nitrogen dioxide into 0.09 N sulfuric acid and 0.2 N sodium hydroxide. In Figures 4 and 5 , these data were plotted as the absorption rate of nitrogen tetroxide vs. its interfacial partial pressure. The dashed lines in these figures are for the absorption into water. As shown in Figure 4, the data for absorption into 0.09 N sulfuric acid were on the line, indicating that the absorption rate of nitrogen dioxide was not affected by the addition of sulfuric acid within the experimental con-
m.
a.
Ind. Eng. Chem., Fundam., Vol. 16, No. i , 1977
165
Table V. Values of Gas Solubility Coefficients (van Krevelen and Hoftiizer (1954)) h+ and h-
l./g-ion
Na+ Hf OH-
0.091 0.000 0.066 0.022
sop-
PZI
[ u t m l x 103
Figure 4. Comparison of rate of absorption into 0.09 N sulfuric acid with that into water.
, 05 PZI
250 t I O
[ o t m l x 103
Figure 5. Effect of 0.2 N sodium hydroxide on rate of absorption of nitrogen tetroxide.
ditions used. Figure 5 indicates, on the other hand, that the absorption rate of nitrogen peroxide into 0.2 N sodium hydroxide was about 7% greater than that into water. The data for the absorption of nitrogen dioxide into aqueous solutions of 5.7-69.8% nitric acid and 2.7-34.1% sodium hydroxide measured by Chambers and Sherwood (1937) and Chilton and Knell (1972) showed that the absorption rate of nitrogen dioxide into such solutions was smaller than that into water; it decreased with the increase of acid or alkali concentrations. Chilton and Knell explained this evidence by the higher viscosities, the smaller activities of water, and the smaller solubilities of nitrogen tetroxide in the aqueous solutions than in pure water. Formation of an acid fog by gasphase reaction of N 2 0 4and HzO followed by condensation with HzO and impact of fog particles on the liquid surface may have contributed to measured absorption rates (Chambers and Sherwood (1937)),causing the observed rates to be greater than the values for the liquid-phase reaction alone. At the experimental conditions of this work, the activities of water in the dilute solutions should have been constant and equal to that of pure water; the partial pressures of N204 were so low that fog formation was not visible. Whether invisible particles were present could not be determined. I t is known that diffusivities in liquids are inversely proportional to viscosities of the liquids. The viscosities of 0.09 N sulfuric acid and 0.2 N sodium hydroxide are about 1%and 5% larger than that of water, which could be responsible for only 0.5% and 2.2% decreases of the absorption rates, respectively. The solubilities of gases in solutions of electrolytes were studied by van Krevelen and Hoftijzer (1954). They found that available data could be correlated by the empirical formula
166
Ind. Eng. Chem., Fundam., Vol. 16, No. 1, 1977
h=h++h-+hG
(16)
The values of h+, h-, and h G for some species for which measurements are possible are listed by Danckwerts (1970). Some of them are shown in Tables V and VI. We have no way to guess the value of hc for nitrogen tetroxide from Table VI, but probably it lies between 0.022 and -0.103. The sums of h+ and h- for sulfuric and sodium hydroxide are 0.022 and 0.157 l./g-ion, respectively. Even if the extreme values of hc from Table VI apply to N204 the value of HIH" according to eq 14 is very near unity a t the acid concentrations used here. Therefore, it seems reasonable to conclude that for the absorption of nitrogen dioxide into 0.09 N sulfuric acid the presence of the acid caused only small changes in the activity of water, the diffusivity, and the solubility of nitrogen textroxide. Consequently, the hydrogen ion concentration in the absorbent should not have affected the absorption rate appreciably, in agreement with the data. On the other hand, h for sodium hydroxide is always positive and large, which means that the solubility should decrease appreciably with an increase of the concentration. For the absorption of nitrogen peroxide into 0.2 N sodium hydroxide, the data showed that the rate of absorption was about 7% greater than that into water, even though the diffusivity and the solubility should have decreased a little. The increase in rate may have come from an additional chemical reaction in this system. That is to say, a reaction between nitrogen tetroxide and hydroxyl ions may have taken place in parallel with the reaction of nitrogen tetroxide with water molecules. In this case, assuming both the parallel reactions to be pseudo-first order, eq 13 can be rewritten as
Nz = H d/D(kl[H2O]
+ k2[OH-])Pzi
(17)
where [H20]and [OH-] stand for the concentrations of water and hydroxyl ions, respectively. One should note that there is no direct evidence concerning this reaction mechanism. In view of the small effect of hydroxyl ion observed there is no way to test the assumptions using the data. The assumption of first-order behavior is possible but not proved. For absorption into water hn[OH-] was negligible since [OH-] in water is only lo-' g-ion/l. For the absorption into 0.2 N sodium hydroxide, on the other hand, [OH-] was 2 X lo-' g-ionfl., 2 x IO6 times larger, and hn[OH-] apparently became large enough to affect the measured absorption rate slightly. Data for the more concentrated sodium hydroxide aqueous solutions which were studied by Chambers and Sherwood and by Chilton and Knell can be explained as follows. Although the term h*[OH-] increased proportionally with the increase of the alkali concentration, the diffusivity and the activity of water decreased somewhat, and the solubility decreased exponentially with the increase of the concentration, according to eq 14. Therefore, a t high concentrations the overall absorption rate decreased with the increase of the sodium hydroxide concentration. Table VI1 shows the experimental data for the absorption of nitrogen tetroxide into alkaline sulfite solutions a t 25 "C;
Table VI. Values of Gas Solubility Coefficients (van Krevelen and Hoftijzer (1954)) ( hat~25 " C ,l./g-ion)
-0.002
0.022
-0.019
0.000
-0.033
-0.054
-0.009
-0.103
Table VII. Absorption of Nitrogen Peroxide into Sodium Sulfite Solutions ( C N ~ O H= 0.1 N) N~ x 107, g-mol/ cm% h = 0.0420 ~ M~ 0.0105 0.0120 0.0146 0.0168 0.0116 0.0133 0.0080 0.0090 0.0136 0.0156 0.0125 0.0143 Cr\a2SO3 = 0.0877 M 0.0135 0.0101 0.0118 0.0172 0.0127 0.0149 0.0150 0.0113 0.0131 0.0100 0.0077 0.0088 0.0161 0.0121 0.0141 0.0184 0.0135 0.0160 0.0118 0.0088 0.0103 cha,~o= 0.153 M 0.0168 0.0121 0.0144 0.0125 0.0093 0.0109 0.0148 0.0109 0.0128 0.0105 0.0079 0.0092 0.0135 0.0098 0.0117 0.0184 0.0133 0.0159 0.0157 0.0113 0.0135 ~
05
0 PPI
IO
[ a t m l x lo3
Figure 6. Rate of absorption of nitrogen tetroxide into alkaline sodium sulfite solutions.
Figure 6 is the plot of the absorption rate vs. the interfacial partial pressure of nitrogen tetroxide with a parameter showing the sulfite concentration. The straight lines in this figure were drawn using a least-squares method. As can be seen in the figure, there clearly was an effect of the sulfite concentration. The diffusivity of nitrogen tetroxide in the solutions should decrease a little with the increase of the sulfite concentration, since the viscosity of the solutions increased with the increase in concentration. From Table V, the sum of h+ and h- for sodium sulfate is equal to 0.113 l./g-ion, which is smaller than the value for sodium hydroxide but is larger than that for sulfuric acid. Comparing this value with hc, for sulfur dioxide, which is the smallest in Table VI, one can assume that h of nitrogen tetroxide in sodium sulfite solutions is probably always positive. Further, one can expect that h of nitrogen tetroxide in alkaline sulfite solutions is positive too. If so, the solubility of nitrogen tetroxide in the solutions should decrease with an increase of sulfite concentration. Nevertheless, the absorption rate increased, probably because of an additional chemical reaction between nitrogen tetroxide and sulfite ions. Assuming another parallel pseudo-first-order reaction, the absorption rate is expressed as
(As the solutions were strongly alkaline, very few bisulfite ions could have been present.) A test of the assumed dependence of the rate of the homogeneous reaction on the concentration of sulfite ion can be made using the data in Table VII. By squaring the slopes of the lines in Figure 6, subtracting the squared slope for 0.1 N sodium hydroxide solution, and dividing by H22D we should obtain a quantity proportional to [SO:32-].Table VI11 lists the results arranged in this manner. Using the data obtained in this work one can evaluate k , ( = h,[HZO]),hz, and h:i. Kramers and co-workers (1961) found the solubility and the diffusivity of nitrogen tetroxide in water to be H = 1.31 X g-mol of Np04/cm:'-atm and D = 1.41 X cm2/s, respectively, a t 25 "C. Neglecting the changes of the diffusivity and the solubility of nitrogen tetroxide in the alkaline sulfite solutions, we have
hl[HzO] = k , = 194 s-' k ? = 147 I./g-ion-s h:j = 8690 l./g-ion-s
30 31 32 33 34 35
0.987 0.987 0.987 0.987 0.987 0.987
36 37 38 39 40 41 42
0.983 0.983 0.983 0.983 0.983 0.983 0.983
43 44 45 46 47 48 49
0.985 0.985 0.985 0.985 0.985 0.985 0.985
b
0.0135 0.0189 0.0151 0.0100 0.0176 0.0162
~
~
0.473 0.885 0.582 0.295 0.786 0.661
0.610 0.972 0.704 0.382 0.868 0.777
0.417 0.609 0.511 0.258 0.574 0.681 0.311
0.662 0.959 0.764 0.424 0.849 1.07 0.569
0.507 0.337 0.443 0.259 0.360 0.613 0.444
1.00 0.634 0.807
0.478 0.725 1.13 0.924
Table VI11
Ion concn, g-ionll. so {?OH0 0 0
0
0.207 0.1 0.1 0.1 0.1
0.0420 0.0877 0.153
N~P?,, g-mol/ cm"-s-atm
0.685 X 0.742 X 0.710 X 1.153 X 1.562 X 1.920 X
Note b
"
0
341 800
1315
By interpolation using values for [OH-] = 0 and 0.207. The values listed represent (N2/pl# minus this quantity for [SOi2-] = 0. The difference is divided by HeaD = 2.42 X lo-" g-molL/ cm4-atm2-s.The slope of the straight line relating the result to sulfite ion concentration is k 3 = 8690 l./g-mol-s. Reaction Mechanism of Nitrogen Tetroxide with Water and Sulfite Ion Although a number of workers have studied the structure of nitrogen tetroxide based on some physical and chemical evidence, the structure has not been established with certainty. Three structures are possible:
I
I1
I11
There are two possible forms for structure I (planar and nonplanar) and three forms for structure I11 (cis, trans, and gauche). Infrared and Raman spectral data indicate that the molecule should have a center of symmetry (Gray and Joffe (1955)). As can be seen, structures I (planar) and I1 do have Ind. Eng. Chem., Fundam., Vol. 16, No. 1, 1977 167
centers of symmetry but structures I (nonplanar) and I11 do not. Thermodynamic data for gaseous nitrogen tetroxide show hindered internal rotation, which makes structure I1 unlikely. Much chemical evidence indicates that nitrogen tetroxide ionizes in solutions to NO+ and NOa- as in eq 3a but not to N02+ and N02-, except in the presence of a strong acid or electron acceptor such as BFs. For example, the reaction between isotopically labeled alcohols and nitrogen tetroxide shows that the reaction involves the formation of NO+ (Anber et al. (1954, 1955)): RO”H
+ NO+N03-
= R0l8NO
+ H+ + NO3-
3. ONO-NO:! = NO+ ible) 4.
5. vant)
(fast, reversible)
2N02 = O N 0 = NO2
2.
+ N03-
(fast, reversible) (rate-limiting, irrevers-
+ HOH = H*ONO+ (fast, irreversible) H20NO+ = H+ + HONO (fast, reversibility irreleNO+
Steps 1 and 2 together effect an isomerization of nitrogen tetroxide from the symmetric form to the asymmetric form. Moll stated that a direct isomerization within a single molecule by the breaking of the N-N bond and the immediate formation of a N-0-N bond appeared to be impossible because the energy barrier to “rocking” an oxygen atom into the N-N bond was probably much greater than 13 cal/mol, the dissociation energy of nitrogen tetroxide. Assuming that 02N-NO2 is stable and only ONO-NO2 is the reacting species, the following modification of Moll’s mechanism is proposed: 1. 0,N-N02 = 2N02 (fast, reversible) 2. 2N0, = ONO-NO, (fast, reversible)
0
/ \H (rate-controlling) H
H
/o\
H
H/
o
The reasons are: (1)Ionization reactions are generally very fast, contrary to Moll’s assumption about his step 3. (2) The reaction between nitrogen tetroxide and water is first order for nitrogen tetroxide and probably also for water. (3) The 168
Ind. Eng. Chem., Fundam., Vol. 16, No. 1, 1977
\O
O \
(19)
This evidence favors structures I1 and 111;for structure I, the extra assumption of easy oxygen atom transfer is required. Recent infrared spectral studies a t liquid hydrogen temperatures (Baldeschweiler (1959), Fately et al. (1959), St. Louis and Crawford (1965)) indicate the existence of both stable symmetric and unstable asymmetric structures. Therefore, it can be considered that structure I (planar) is predominant in solid, liquid, and gas phases, and that unstable structure 111, even though present in a small amount, plays an important role for chemical reaction, yielding NO and NO3 radicals in a gas phase, and NO+ and NOa- ions in a liquid phase. Based on the concept of Bent (1963), Moll (1966) suggested that unstable asymmetric structure I11 is an intermediate in the reaction by which stable symmetric structure I forms the ions NO+ and NO:?-: 1. 02N-NO2 = 2N02
variation of the rate of absorption observed in this work implies that the reaction between nitrogen tetroxide and sulfite ions was kinetically first order for both nitrogen tetroxide and sulfite. The reaction of nitrogen tetroxide with a sulfite ion may be \O /o-soJ2O\ N-0-N /O + S O , ’ - = N-0-N
An alternative mechanism is 0 \N-O-N/O + so:= O\
so,-0
\
/02-
N-0--N ‘ 0
According to the chemical bonding theory, breaking a bond by the addition of an ion or a molecule is easier the closer the bond is to the point of addition. Therefore, the former reaction is more the probable one. Consequently, we suggest that the rate-limiting steps for the reactions of dissolved nitrogen tetroxide with water or sulfite ions are the reactions between ONO-NO2 and these substances. Acknowledgment The Chiyoda Chemical Engineering and Construction Co., Ltd., Japan, is gratefully acknowledged for their financial support of Yohji Kameoka’s study. Nomenclature
A = surfacearea, cm2 C = concentration, g-mol/l. c = concentration in liquid, g-mol/cm:’ D = diffusivity, cm2//s H = Henry’s law constagt, g-mol/cm:3-atm h = sum of contributions referring to the species of positive and negative ions present, and to the species of gas H o = Henry’s law constant in pure water, g-mol/cm:’-atm I = ionic strength of the solution, g-ion/l. K = equilibrium constant for the dimerization of N o r , atm-I h = first-order reaction rate constant, s-’ k , = pseudo-first-order rate constant of the reaction between N204 and water, s-I k G = gas-phase mass transfer coefficient, g-mol/cm’-satm h l = second-order rate constant of the reaction between N204 and water, l./g-mol-s k 2 = second-order rate constants of the reaction between N204 and OH-, l./g-ion-s k r 3 = second-order rate constant of the reaction between N204 and S03’-, l./g-ion-s N = molar flux of tetravalent nitrogen, g-mol/cm’-s N1 = molar flux of N02, g-mole/cm2-s N 2 = molar flux of Nz04, g-mol/cm2-s P = total pressure, atm p = partial pressure, atm R = radius of spheres, cm t = time of exposure, s U = liquid velocity, cm/s z, = valency of j t h ion
Greek Letters 7 = distance from interface divided by gas-film thickness 6’ = angle of inclination from the vertical, deg
4 = absorption rate for one sphere, g-mol/s
Subscripts b = bulk E = equator of spheres (0 = 90°) i = interface 1 = NO2 2 = NnO4 Literature Cited Anber, M., Dostrovsky, I., Samuel, D. H.,Yoffe, A. D., J. Chem. SOC., 3603 (1954). Anber, M., Taube, H., J. Am. Chem. SOC.,77, 2993 (1955). Astarita, G., Chem. Eng. Sci., 17, 708 (1962). Baldeschwieler, J. D., Ph.D. Thesis, University of California, Berkeley, 1959. Bent, H. A., Inorg. Chem., 2, 747 (1963). Caudle, P. G., Denbigh, K. G., Trans. Faraday Soc., 47, 39 (1953). Chambers, F. S.,Sherwood, T. K., Ind. Eng. Chem., 29, 1415 (1937). Chilton. T. H.,Knell, E. W., Preprint for PACHEC '72, Session-13, p 75, 1972.
Corriveau, C. E., Pigford, R. L., UCRL-20479 (University of California, Berkeley, Lawrence Radiation Laboratory Report) (1971). Danckwerts, P. V., "Gas-Liquid Reactions", pp 18-20, McGraw-Hili, New York, N.Y., 1970. Dekker, W. A., Snoeck, E., Kramers, H.. Chem. Eng. Sci., 11, 61 (1959). Denbigh, K. G., Prince, A. J., J. Chem. Soc., 790 (1947). Fately, W. G., Bent, H. A., Crawford, B., J. Chem. Phys., 31, 204 (1959). Gray, P.,Yoffe, A. D.. Chem. Rev., 55, 1069 (1955). Kramers, H., Blind, M. P. P.,Snoeck, E., Chem. Eng. Sci., 14, 115 (1961). Moll, A. J., Ph.D. Thesis, University of Washington, 1966. Peters, M. S., Ross, C. P., Klein. J. E., A.I.Ch.€.J., I , 105 (1955). St. Louis, R. V., Crawford, B., J. Chem. Phys., 42, 857 (1965). van Krevelen, D. W., Hoflijzer, P. J., Trans. inst. Chem. Eng., 32, 560 (1954). Verhoek, F., Daniels, F., J. Am. Chem. Soc., 53, 1250 (1931). Wendel, M. M., Ph.D. Thesis, University of Delaware, 1956. Wendel, M. M.. Pigford, R. L., A./.Ch.E.J., 4, 249 (1958).
Received for reuieu January 12, 1976 Accepted October 28, 1976
Ind. Eng. Chem., Fundam., Vol. 16, No. 1, 1977
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