ENGINEERINGr DESIGN, A N D PROCESS DEVELOPMENT (3) Freeman, H. A., "Industrial Statistics," pp. 1-61, New York, John Wiley 6: Sons, Inc., 1947. (4) Groggins, P. H., "Unit Processes in Organic Synthesis," 4th ed., PP. 75-7, New York, McGraw-Hill Book Co., Inc., 1952. (5) Irick, P;, Purdue Vniversity, private communication. ( 6 ) Lesser, R,, Glaser, A,, and Aczel, G., Ann., 402, 1-5 (1914). (7) hladinaveitia, A., and DeBuruaga, 3. S.,Andes. soc. espni2. !is. gutm., 27, 647-58 (1929). (8) Sah, P. T., Rec. trau. chim., 59, 461-70 (1940). (9) Shulze, K. E., Ber., 17, 842 (1884).
(10) Veldstra, H., and Wiardi, P. W., Rec. traa. chim., 62, 75-84 (1943). (11) Vesely, V., and Kapp, J., Ibid., 44, 360-75 (1925). (12) Vesely, V., and Pac, J.,Collection Czeciioslou. Cheni. Communs.,2, 471-85 (1930). (13) Vesely, V., and Stursa, F., Chem. Listy, 29, 361-3 (1935). (14) Vesely, V.,and Stursa, F., Collection Czechoslor. Chem. Cominuns., 6, 137-44 (1934). RECEIVED for review September 14, 1053.
ACCEPTEDDeoembcr 16, 1453,
l o w Temperature Catalytic Oxidation of Ammonia H. F. JOHNSTONE, E. T. HOUVOURASl, w. R. SCHOWALTER
AND
University of Illinois, Urbana, 111.
The oxidation of ammonia at low temperatures is of interest because of the types of products that are formed and because the reactions, which are easily followed, give fundamental information on the mechanism of heterogeneous catalysis. The reactions were studied in the range 180" to 430" C. in fused salt media and in a porous tube impregnated with a rare earth oxide or a manganese oxide-bismuth oxide catalyst. Only those fused salts whose cations form coordination complexes with ammonia act as catalysts. Cuprous chloride-potassium chloride melts rapidly catalyze the oxidation to nitrogen and water at a temperature of 225" C. An equilibrium i s established between cuprous and cupric ions that depends on the ratio of ammonia to oxygen in the gases. With praseodymium oxide as a catalyst, the overall reaction is zero order with respect to ammonia at low oxygen to ammonia ratios and low temperatures. With the more active manganese oxide-bismuth oxide catalyst, the reaction i s first order with respect to ammonia. A mechanism for nitrous oxide formation in the presence of defect oxide catalysts i s proposed. Semiconductors of the Type P defect lattice are more active catalysts than those with Type N lattices.
S
TUDIES on the mechanism of ammonia oxidation a t high
temperatures using a tubular reactor have been reported by Andrussow ( 2 ) . I n this laboratory, the tubular reactor has been used for studies of the oxidation of sulfur dioxide (S), catalytic hydrogenation of ethylene (6, 2 4 , the reaction of carbon monoxide and hydrogen ( 7 ) , and of the reaction of steam with carbon ( I S ) . It has several advantages for studying mechanisnis of heterogeneous reactions, such as ease of temperature control, choice of sampling positions, reproducibility of results, and amenability to mathematical treatment. I n this xork, a study v a s made of the reaction of ammonia with oxygen a t relatively low temperatures in two types of reactors-a fused salt media and a poroua tubular reactor impregnated nith oxide catalysts of two types. The n-ork n i t h fused salts is an extension of the general study of catalysis in these systems made by Norman and Johnstone (19).
A4ndrussow (1) postulates an initial reaction between aninionin and molecular oxygen in which nitroxyl (HKO) is formed. Schlecht and von Sagel (21) oxidized ammonia with e x m s oxygen a t 275' to 300" C. with a mixed catalyst of ferric. oxide-bismuth oxide-manganese dioxide (Fez03-Bi203-Mn02). Postnikov (20) obtained 84% yield of nitrous oxide on a niaiiganese-iron-bismuth catalyst a t 200' t o 300" C. Krauss and Neuhaus (17) used oxides of manganese, bismuth, barium, iron, and nickel, and obtained nitrous oxide, nitric oxide, and nitrogc7n. The relationship between the activity of the catalyst and the amount of adsorbed active oxygen was further shown by K r m s (16). Kobe and Hosman (15) also studied the reaction with the manganese oxide-bismuth oxide catalyst. I n general, it appears that crystalline oxide catalysts of the defect-oxide lattice t y l x are necessary for the formation of nitrous oxide.
Oxidation in Fused Salt Media
Previous Work on Ammonia Oxidation The mechanism of the oxidation of ammonia has been discussed by Zawadski (28). There are a t least three theories of the course of the reaction, Bodenstein (5) and Krauss (16) propose that the initial step is the reaction of ammonia with adsorbed oxygen [ O ]sdS. t o form hydroxylamine, while Zawadski is of the opinion that imide (NH) is formed in the initial step. 1 Present address,
E. I. du Pont de Nemoura & Co., Pna., Benger Labora-
tory, Waynesboro, YE.
702
While the similarity between the catalytic action of finely divided metals and their respective cation melts has been noted in the previous paper (19), t,here are other mechanisms of catalysiP by ionic melts, such as the one involving intermediate compound formation t h a t takes place in the chlorination of light paraffin hydrocarbons with air and hydrogen chloride in the presence of fused copper chloride (11). Thus in the low temperature oxidation of ammonia, the formation of a loose coordinat>ion compound, either with ammonia or with oxygen, might b e t'he mechanism by vhich a fused salt could act as a c:italyst. 4c-
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
Vol. 46, No. 4
UNIT PROCESSES cordingly, several types of fused salt systems were studied for catalytic activity a t various temperatures in ammonia concentrations. The 68 mole % cuprous chloride-32yo potassium chloride eutectic melt was found to he the only active catalyst. Khen pure ammonia was passed through this melt ( M ) , it was first absorbed as a loose compound, presumably M[(XH3)2,] +*> from which it could be recovered quantitatively. The deformation of the ammonia molecule in the complex rendered it susceptible to oxidation. This took place when oxygen was passed through the melt simultaneously with or subsequent to the ammonia. The oxidation products were always nitrogen and water. The formation of cupric ions could be observed in the change in color of the melt from colorless to dark and opaque as the oxidation proceeded. Reconversion to the colorless cuprous form was readily accomplished by simply shutting off the oxygen. Melts of ferric chloride-sodium chloride and manganous chloride in sodium chloride did not catalyze the ammonia oxidation. The fused salt reactor consisted of a 22-mm. inside diameter glass tube, 15 inches in length, which was heated externally. The preheated gaseous reactants were introduced near the bottom of the tube below the surface of the salt. Analysis of the products was made after steady state conditions were obtained. A study was made of the significant variables in the oxidation reaction, such as mole ratio of oxygen to ammonia, space velocity, and temperature. The conversion increased with increasing ratio of oxygen to ammonia in the range from 0.18 to 1.06. It was 18.2% a t a mole ratio of 1.02 and a space velocity of 7.5 min-1 at 220" C. Higher conversions were obtained a t low space velocities; the maximum was 21.6% a t a space velocity of 2.1 min.-l and a t 220' C. The conversion decreased with increasing temperature from 218' to 273" C. The decreased conversion is attributed to the lower stability of the ammine complex and lower ammonia adsorption a t the higher temperatures ( 1 3 ) . Since no decomposition occurred in the absence of oxygen, the reaction of ammonia and oxygen in the melt appears to be a true oxidation reaction and not a decomposition followed by the oxidation of hydrogen. Because of the nature of the products of the reaction, no further studies were made of the reaction in the fused salt media.
where
L n Pn
Jo
J1
= length of reactor, em. = running index, 1, 2, 3,. . Jl(P7L) =
g
.
J a ) ZD. = Bessel function of zero order and first kind = Bessel function of first order and first kind
The dimensionless group, kd/2D, is the ratio of the mean time required for diffusion of the reactants to the tube surface to the mean time required for the heterogeneous reaction. If k r d C represents the number of moles reacting per unit time and unit length of tube, the time required for the total number of moles in this volume to react will be 0 reaction =
~
TdZC 4
,
1 k r d C - 4k
(4)
where C is average concentration, and 0 is contact time In seconds. But the time required for diffusion is (d/2)2 = d2 0 diffusion = -
8D
20
(5)
therefore
e diffusion _ -kd 0 reaction
20
Equation 4 also shows that k is equivalent to the ordinary firstorder reaction rate constant multiplied by the hydraulic radius. For small values of kd/2D, say less than 0.1, Equation 3 can be simplified to give
C, = ca
[g +
.9-2(kd/2D)r
(7)
21
where r = 4DL - a dimensionless group.
vd2 Oxidation in Tubular Reaction Theory of Tubular Reactor. Damkohler (9) has developed a mathematical analysis for first-order constant volume heterogeneous reactions in a tubular reactor. He assumed that: 1, The temperature along the length of the tube is constant. 2. There is no velocity gradient in the radial direction. 3. Steady state conditions prevail. 4. The pressure along the length of tube is constant. A material balance can be written for the limiting reactant in cylindrical coordinates
When the reaction rate is of the same order as the rate of diffusion the following boundary condition is valid
where C = ammonia concentration; subscript refers to entering gas; subscript refers to leaving gas, g. moles/cc. d = inside diameter of tubular reactor, cm. D = diffusivity of ammonia, sq. cm./sec. 12 = reaction velocity constant expressed in dimensions of length/time T = coordinate in radial direction, cm. z = distance in axial direction, cm. The solution of Equation 1 for the average exit concentration with this boundary condition is m
April 1954
Oxidation with Rare Earth Oxide Catalyst in Tubular Reactor. The apparatus was similar to that used in the previous work (3). The housing for the catalyst tube consisted of a I1/8-inch inside diameter stainless steel tube, 24 inches long. An additional 11inch length of the same tubing was used a t the exit end of the reactor to ensure a uniform longitudinal temperature. The catalyst tube was sealed in the housing with a layer of alundum cement to provide good heat transmission. The preheater consisted of a 12-inch length of the stainless steel tubing packed with 5-mm. glass beads to provide a large heat transfer area. The catalyst support tube was of l/Z-inch inside diameter, 24 inches long, l/s-inch wall Norton Alundum RA 98. This material has a porosity of 47% and a thermal conductivity of 6 to 7 B.t.u./(hr.)(sq. ft.)(" F./ft.). The tube was soaked in concentrated nitric acid, washed with distilled water, dried, and weighed. It was then impregnated with the catalyst by soaking in a concentrated solution of mixed rare earth nitrates a t 90" C. for 2 hours. The solution was prepared by dissolving 30 grams of a mixture of approximately equal proportions of praseodymium and neodymium oxides containing a trace of ceric oxide in 100 ml. of dilute nitric acid and concentrating the resulting solution by evaporation. After impregnation, the nitrates were converted to the oxides by gradual1 raising the temperature from 100" to 450' C. over a period o f 1 2 hours and holding a t 450" C. for several days. The catalyst deposited in the tube wall amounted to 3.68% of the total tube weight. The temperature along the tube was measured a t five points by Chromel-Alumel thermocouples silver-soldered in small wells drilled within 1 / 3 t inch of the inner wall of the stainless steel tube. Indirect heating was provided by a tubular furnace around the catalyst tube housing. Four heating elements, one wound over the entire length, two over the end sections of the main core, and the fourth over the exit end of the core, permitted temperature control within 1 2 ' C. of the desired value. The flow rates of oxygen and ammonia were measured by capillary flowmeters. Mixtures of ammonia and oxygen were passed over the catalyst a t measured rates ranging from 0.27 X 10-2 to 2.87 X
INDUSTRIAL A N D ENGINEERING CHEMISTRY
703
ENGINEERING, DESIGN, AND PROCESS DEVELOPMENT
I
I
0
I
5
I
I
10 I5 20 CONTACT TIME-0,SECONDS
0 1.00
I
25
30
1.05
j I ID
kZ ,
(A-rM
1.15
I20
I
,-"z
Figure 1. Effect of Contact Time on Fractional Conversion of Ammonia with Rare Earth Oxide Catalyst
Figure 2. Effect of Oxygen Pressure on Rate of Oxidation of Ammonia with Rare Earth Oxide Catalyst
gram moles per minute (19). The product gases were anelyzed for ammonia and IT-ater vapor by gravimetric adsorption on Bnhydrone. After dissolving the mixture in water, the ammonia was determined by titration Kith standard hydrochloric acid; il-ater was determined by difference. One sample of the unadsorbed gas mixture Tyas analyzed for oxygen by means of Oxsorbent solution; nitrous oxide and nitrogen dioxide m r e determined in a second sample by infrared analysis; nitrogen v-a3 found by difference. Sitrogen and oxygen balances n-ere c d culated from the analyses and flow rates. Sitrogen dioxide was seldom found in more than trace concentrations and it x a s assumed, therefore, t,hat nitric oxide was not present. Comparison of x-ray diffraction pat,terns of the catalyst, the alundum catalyst support and of the mixture of the catalyst aiid mpport shox-ed that the crystal structure of the catalyst iTas not altered by the alundum. The properties of the praeeodpniium oxide used indicated that it was the same as the dark brown, nonstoichiometric PrOi,sa described by Xartin (18), which is apparently an electron, or N-type semiconductor. Since the activity of t,he catalyst was substantially constant over a long period of time,. it was assumed that its structure was not changed by the reaction.
( 1 2 ) . These varied from 0.439 sq. em. per second a t 180" C. to 0.512 sq. em. per second a t 220" C.
The results obtained Ti-ith the tubular reactors arc showi i n Tables I to 111. The values listed under Fraction Ammonia
Oxidation with Manganese Oxide-Bismuth Oxide Catalyst. The reactor rias similar to that described above (28). The catalyst was prepared by impregnating another '/*-inch alundum tube mith a concentrated solution of manganous and bismuth nitrates having a molecular ratio of 1.2. The tube was then immersed in ammonium hydroxide. The deposited metal hydroxides were converted t o their respective oxides by heating a t 375" C. for 2 days. The catalyst, deposited in the tube was 13.4 grams or 5,3y0of the total tube weight,. Operating conditions xere chosen so that the approximations of Damkohler were valid. Isothermal operation and steady state conditions vere attained prior to sampling. The preesure drop through the reactor vias small in comparison with t'he absolute static pressure. The maximum Reynolds No. v I w 150, and the maximum increase in the molar gas rate due to the reaction was 3%. The assumption of rodlike flovi is justified in view of previous experience with flon- a t low velocities in wetted-wall columns ( 2 3 ) . The results for rodlike and parabolic flow are actually the same RThen the diffusional resistance is small in comparison mith the chemical resistance, as in the present case. Values for the diffusivity of ammonia in the gas mixtures xere calculated by the method recommended by Hirschfelder et al.
2'04
xidafion of Ammonia with Manganese Oxide-Bismuth Oxide Catalyst Oxygen to ammonia ratio, 9:1
INDUSTRIAL A N D ENGINEERING CHEMISTRY
Vol. 46,No. 4
UNIT PROCESSES
Table
I.
Table II.
Oxidation of Ammonia with Rare Earth Oxide Catalyst Fraction
Time Contact (B), Seo.
T J ~ ~ ~Ammonia ~ $ Converted, & Yield, % N20
Na"
(cS/cO)
Average Initial Gas Composition, 33.3%
"I,
66.7% Oa
(330' C.) 7.0 26.3 32.1
0.96 0.86 0.80
0.77 0.64 0.61 0.50 0.38 0.27
92.7 91 .o 91.2 91.0 91.0 90.1
7.3 9.0 8.8 9.0 9.0 9.9
(330' C.)
0.90 0.86 0.70 0.58
5.2 7.6 8.5 10.4 15 0
0.61 0 49 0 40 0 36 0.09
88.2 89.8 89.8 91.7
11.8 10.2 10.2 8.3
89.2 89 0 90.2 88.2 84.5
10.8 11 .o 9.8 11.8 15.5
Contact Time ( B ) , Seo.
10.8 12.4 16.0
89.2 87.6 84.0
Average Initial Gas Composition, 10.0%
1.7 3.3 5.8 8.6 10.8
1.7 3.4 5.9 8.6 11.2
"3,
90.0%
01
Ammonia Converted,
SzO, Yield,
%
(Ce/Ca)
C.;kd/2D
0 2
= 0.018) 0.92
...
0.84
t . .
0.72 0.65 0.62
93.5 69.6 53.3
2.04 4 03 7.00 10.20 13.40
I .60 3.50 5.95 6.70 14.10
0.75 0.63 0.46 0.30 0.20 =
14.3 14.2 22.0 19.2 16.2
0.085) 0.75 0.40 0.28 0.25 0.05
45.6 50.8 34.2 50.7 32.5
Average Initial Gas Composition, 20.0% NHs, 80.0% 02 (180' C.; kd/2D = 0.012)
89.8 89.7 88.6 85.3
10.2 10.3 11.4 14.7
3.4 5.8 8.6 11.0
0.23 85.3 0.19 82.6 0.11 82.4 in small quantities.
14.7 17.4 17.6
3.2 6.2 8.8 11.4
3.67 6.42 9.36 12.00
0.89 0.83 0.77 0.72
79.3 67.0 47.8 54.4
(200' C.; kd/2D = 0.028)
(380' C.) 5.3 6.3 7.6 Includes NOz present
Fraction Ammonia Unconverted
1.86 3.62 6.37 9.28 11.80
(220° C.; k d / 2 D 1.2 2.6 4.7 5.3 11.2
(330' C.) 0.60 0.44 0.27 0.09
T
(200' C.; M/2D = 0.047)
(380' C.)
11.8 17.5 21.0 24.6
4DL vdz (180'
Average Initial Gas Composition. 15.0% NHs, 85.0% Oz
0.56 0.25 0.09
Oxidation of Ammonia with Manganese OxideBismuth Oxide Catalyst
Average Initial Gas Composition, 10.0% NHr, 90.0%
(380' C.)
5.2 8.7 11.9
Manganese Oxide-Bismuth Oxide Catalyst 10s(Liter)(Sec.)(Atm31/21 [k: X lOZ(Sec:-')(Atm.)-1/2] 0.59 1.50 ... 3.09 1.5 7.3
... ...
Table 111.
Average Initial Gas Composition, 20.0% KH3, 80.0% Oz
7.0 11.7 23.4 31.7
Specific Reaction Rate Constants
Eg$e:$&%&bz\
15.4 8.6 7.3
84.6 91.4 92.7 (380' C.)
6.1 9.6 11.0 14.8 18.9 24.5
Ek'
Temp., O C. 180 200 220 330 380
Rare
3.82 7.38 10.50 13.70
0.73 0.58 0.49 0.41
8.2 3.5 16.0 I1 6
(220' C ; k d / 2 D = 0 . 0 5 2 )
Unconverted are averages of analyses of several samples of the product gases after steady state was attained. These analyses agreed within 1 3 % . The ammonia conversions obtained with the rare earth oxide catalyst are shown in Table I and Figure 1. For high concentrations of ammonia and low temperatures, the straight lines show that the reaction is zero order with respect to ammonia. -4t 380" C., the oxidation approaches a firstrorder reaction. The apparent activation energies based on the zero-order rate constants for the 3 mole ratios of 9 9 , 4:1, and 2:l are 23.7, 24.0, and 24.2 kcal. per mole, respectively. There is evidence from previous investigators (5,17')to support the theory that oxygen is adsorbed on the surface in an activated state, presumably as atomic oxygen. If the controlling step is the reaction of ammonia with adsorbed oxygen, which may be expressed by a Langmuir-type adsorption isotherm, the reaction rate equation can be written
where b = adsorption equilibrium constant p = partial pressure, atm. The low activity of the rare earth oxide catalyst indicates a low concentration of active centers. At low temperatures and high
April 1954
2 4 4.8 7.5 10.1
3.09 6.12 9.44 12.70
0.66 0.45 0.29 0.20
Average Initial Gas Composition, 33.3%
15.2 16,1
11.2 9.6 "1,
66.7% 02
(180' C.; k d / 2 D = 0.0074) 3.3 5.0 8.1 10.8
3.62 5.46 8.85 11.80
0.93 0.91 0.85 0.83
10 10 10 11.3
(200' C.; kd/2D = 0.020) 3.6 5.4 8.9 12.0
4.24 6.36 10.60 14.20
0.79 0.71 0.60 0.52
10.9 25.5 28.3 20 0
(220' C.; k d / 2 D = 0.034)
3.4
5,l
8.6 11.2
4.34 6.46 10.80 14.20
0.69 0.55 0.40 0.32
...
15.1 15.5 8.1
ammonia concentrations, therefore, the small amount of adsorbed oxygen becomes the limiting reactant. This means that the ammonia concentration a t the surface reaches saturation and remains a constant. With a large initial excess of oxygen the change in oxygen partial pressure is small and pol is also constant. If a new constant ko is used to designate the right hand side of Equation 8, integration will then give
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
705
ENGINEERING. DESIGN. AND PROCESS DEVELOPMENT The first-order reaction rate constant may be calculated from the value of k d / 2 D , using Equation 4, or from integration of Equation 8 as follows:
where
li.:
= first-order reaction rate constant (sec.-l)(atm.)-l/2
eince
e
= L =
d"
v
40
'10' D
_
.
.
-
-
vd2
- dz 40'
where v = average velocity of gases in reactor in cm./sec.
A plot of
should give a straight line with a slope of l/ki from which the value of k; may be found. Such a plot is shown in Figure 6. Values for the reaction rate constants are in Table 11. The activation energy for the first-order reaction is 18.6 kcal. per mole of ammonia or per atom of oxygen.
Figure 4.
Oxidation of Ammonia with Manganese Oxide-Bismuth Oxide Catalyst Oxygen to ammonia ratio,
4:l
where k o = zero-order reaction rate constant in moles NH, oxidized/(liter)( see.)
4 s shown in Figure 1, the data agree with Equation 9. Constant k~ depends on the oxygen partial pressure. This may be shown by substituting IC; for k'C in Equation 8 and rewriting
where k ; = zero-order reaction rate constant in moles KH3
R
oxidized/(liter)(sec,>(atm.)'/2 =
rs-2 4DL
rate of oxidation of ammonia in moles/(liter)(sec.)
Rearranging
vd Figure
5.
Oxidation of Ammonia with Manganese Oxide-Bismuth Oxide Catalyst Oxygen to ammonia ratio, 2:l
The straight lines in Figure 2 show that the reaction rate with the rare earth oxide catalyst is represented by Equation 11. Values of rC6 calculated from the slopes of these lines are listed in Table 11. The calculated activation energy is 24.8 kcal. per mole. The results obtained with the manganese oxide-bismuth oxide catalyst are shown in Table I11 and in Figures 3 t o 6. The graphs show that the data are compatible with Damkohler's analysis for a first-order heterogeneous reaction. The catalyst is much more active than the rare earth oxide catalyst. The values of kd/2D can be obtained from the slopes of the straight lines of log C,/C, vs. 7 according t o Equation 7, and are shown in Table 111. These satisfy the prescribed condition. They indicate t h a t the diffusional resistance is less than 6% of the over-all reaction resistance for all of the conditions investigated with this catalyst.
706
The difference between the activation energy found for the apparent first-order reaction on the manganese oxide-bismuth oxide and that for the apparent zero-order reaction on the rare earth oxide may be due to the difference in active centers. As shown later, the catalysts have different types of defective lattices. The controlling reaction in both cases, however, is t h e formation of active adsorbed oxygen. Values of the activation energy of adsorbed oxygen have been reported between 20 and 80 kcal. per atom of oxygen on platinum catalysts (26)Role of Defect Oxides in Catalysis. Recently the theory of semiconductors has been applied with some success to the interpretation of catalyst activity. It is of interest to consider t h e results of the present work in terms of this theory. Some oxides-e.g., those of zinc and cadmium-can lose n
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
Vol. 46, No. 4
UNIT PROCESSES number of defects and the partial pressure of oxygen and ammonia a t the catalyst surface. According t o Equation 15, on an N-type catalyst an increase in oxygen partial pressure tends to decreaee the number of defects and thus decreases the reaction rate. This has been observed in the decomposition of nitrous oxide on a zinc oxide catalyst; and Bevan and Anderson (4) have found that the conductance of zinc oxide changes with the oxygen pressure in the range 0.01 to 1000 mm. as PO:.^'. If the same relationship applies to the present study, the variation of oxygen pressure in the range 500 t o 685 mm. pill scarcely affect the number of defects. On the other hand, h y g e n is a reactant in this case and an increase in its partial pressure will increase the reaction rate. With a nearly constant number of defects, the net effect of an increase in oxygen will be an increase in rate. Manganous oxide readily loses electrons to form cations of high valency or to combine with excess oxygen-e.g., Mn2+to Mn3+. I n this way i t may be a P-type catalyst. Oxygen adsorption on the manganous oxide surface may be written as
+).(CV).O-z 1.00
1.05
2, p
1.15
1.20
where (CV), = a vacant cation sites per mole and (f), missing electrons, or a positive carriers per mole.
small fraction of their oxygen atoms, and the resulting defective structures contain an excess of metal ions; whereas some oxides -e.g., those of nickel and copper-can gain extra oxygen ions above their stoichiometric composition. The former, with available electrons, are called N-type semiconductors while the latter, with missing electrons, are called P-type semiconductors. Wagner (96) has demonstrated that the catalytic activity may be related to the number of defects per mole, which in turn may be estimated from the electrical conductance. The rate of material transport within a defective lattice apparently determines the catalytic activity and the chemical reaction rate. The addition of another component to a catalyst, however, may increase the conductance without causing a corresponding increase in activity. Cremer (8) has shown, for example, that the addition of gallium oxide to zinc oxide results in an increase in the activation energy without a corresponding increase in the activity of the catalyst. By means of thermoelectric measurements, Martin (18) has ahown that, a t low temperatures, PrOt.sa is effectively Pro$ deficient in oxygen and not PrzOa with excess oxygen. Using the nomenclature of Weyl ( d 7 ) to describe the various defects i n a lattice structure, the chemical reactions a t the surface of praseodymium oxide can be expressed as
+ zN+6(e);H:f
+
Pr+4(e);(.+,)(AV).+,0;_2,_.
+
~ ~ n ~ ~ ~ + z ~ ( + ) ~ t ~ (16) ~ ( C V ) * * z ~ O
1.25
CATMS-~
Figure 6. Effect of Oxygen Pressure on Rate of Oxidation of Ammonia with Manganese OxideBismuth Oxide Catalyst
Prt4(e);,(AV),0;2.
+ x/202
I
+ X N + ~ H ~ ~ ( ~ ) (14) ;O-*
and
= a
According to Equation 16 on a P-type catalyst the cation vacancies increase with the adsorption of oxygen, thus resulting in higher catalyst activity and faster reaction rate. This agrees with the results of the present studies. A change in temperature can bring about a shift from one type of semiconductor to another. At high temperatures, praseodymium oxide becomes a P-type catalyst (18). Thus, the reaction rate is more rapid, and the mechanism approaches first order in respect to ammonia.
Conclusions The oxidation of ammonia to nitrogen and water is catalyzed by fused salts of which the cations form complex ammines. Cuprous chloride-potassium chloride melts are active a t 225' C. An equilibrium is established between cuprous and cupric ions that depends on the ratio of ammonia to oxygen in the gases. In catalysis of ammonia oxidation by metallic oxides with defect lattices, the rate of oxidation is determined by the amount of active oxygen on the catalyst surface, which in turn is fixed by the number of active centers available. When a rare earth oxide catalyst is used, the over-all reaction is zero order with respect to ammonia a t high ammonia concentrations. When a manganese oxide-bismuth oxide catalyst is used, the reaction is first order with respect to ammonia under all of the operating conditions investigated. The different results obtained with the two catalysts can be explained in terms of the catalyst activity. The effect of oxygen partial pressure on the reaction rate can be predicted by the Langmuir adsorption isotherm. Oxides which are semiconductors of the P-type are more active catalysts than those which are N-type.
Acknowledgment where IAV), = a vacant anion site8 per mole and (e), available electrons per mole.
= k ~ ao
The authors express their appreciation to L. F. Budrieth and C. Y . Shen for their suggestions concerning the catalyst mechanisms, and to J. D. Stroupe and H.S. Gutowskyfortheirassistance on infrared analysis.
The catalyst in the two quasi reactions acts as a medium for bhe transfer of electrons. Expressed in terms of Fajan's quanticule theory (IO) by Equation 14, ammonia removes oxygen ions from the catalyst and produces more anion defects. By Equation 15, oxygen combines with electrons and thus removes the defects. In other words, there is a dynamic equilibrium between the .April 1954
Literature Cited (1) Andrussow, L., Angew. Chem., 63, 21 (1951). f2) Andrussow. L.. Ber.. 60B. 2005 (1927).
(31 Banon, T., 'Makning, W. 'R., and Johnstone, H. F., Chem. Eng. Pi-ogr., 48, 125 (1952).
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ENGINEERING, DESIGN, AND PROCESS DEVELOPMENT (4j Bevan, D. J. hI., and Anderson, J. S., Dlscussions Faraday SOC., No. 8 (1950), 238. (5) Bodenstein, AI., Trans. Electwchenz. SOC.,71, 353 (1937). (6) Byerly, J., 1\13, thesis, University of Illinois, 1948. (7) Campbell, W. AI., and Johnstone, H. F., IND.ENG.CHmf., 44, 1570 (1952). (8) Cremer, E , , and Marshall, E., Monafsh,., 82, 841 (1951). (9) Damkohler, G., 2. Eleklrochem., 42, 846 (1936). (10) Fajans, X., Chem. Eng. News, 27, 900 (1949). (11) Gorin, E., Fontana, C. AI., and Kidder, 0. d.,IND. ENG. CHEW,40, 2128 (1948). (12) Hirsohfelder, J. O., Bird, R. B., and Spots, E. L., J . Chem,. Phys., 16, 968 (1948). (13) Houvouras, E. T.,Ph.D. thesis, University of Illinois, 1953. (14) Johnstone, H. F., Chen, C. Y., arid Scott, D. S., ISD. EXG. CHEM., 44, 1564 (1952). (15) Kobe, K. A , , and Hosman, P. D., Ihid., 40, 397 (1948). (16) Krauss, UT.,2. Elektrochem., 53, 320 (1949); 54, 264 (1950). (17) Krauss, IT’..and Seuhaus,-4,, 2. p h y s i k . Chem.,B50,323 (1941).
(18) Martin, R . L., Nature, 165, 202 (1950). (19) Sorman, K~ E., and Johnstone, H. F., IXD.ENG.CHEX, 43, 1553 (1951). (20) Postnikov, V. F., Kus’niin, L. L., and Tsal’m, N.K., J . Chem. I n d . (U.S.S.E.), 13, 13.28 (1936). (21) Schlecht, L., and von Nagel, d.,Ger. Patent 503,200 (1930). (22) Sohowalter, 1‘7. R., bI.6. thesia, University of Illinois, 1953. (23) Sherwood, T. K., and Pigford, R. L., “Absorption and Extraction,” 2nd ed., p. 81, Sev York, hlcGraw-Hill Book Co., 1952. (24) Stocker, D., M.S. thesis, University of Illinois, 1948. (25) Taylor, G. B., Kistiakowsky, G. B., and Perry, J. H., J , P h ~ e , Chem., 34, 799 (1930). (26) Wagner, C., J . Chem. Phys., 18, 62 (1950). (27) Weyl. W. A., Penn. State Cniversity, Mineral Ind. Expt. Sta., Bull. 57, 81 (1951). (25) Zanadski, J., Discussions Faraday SOC.,No. 8 (1950), 140. RECEIVED for review November 27, 1953.
ACCEPTEDFebruary 6, 19.54.
JAMES E. KNAP’, E. W. COMINGS2, and H. G. BRICKAMEW University of Illinois, Urbana, 111.
The gas phase alkylation of isobutane with propylene was studied in a flow system at 4 0 0 ” C, and at pressures from 4000 to 15,000 pounds per square inch, in the presence of 1,2,3-trichloropropane and 1,2-dichIoropropane. The major products of the reaction are 2J-dimethylpentane and 2-methylhexane. Yields of total alkylate und of isoheptanes increase with increasing pressure, The concentration of 2,2-dimethylpentane increases with pressure while that of 2-methylhexane decreases, indicating an increase in the relative reactivity of the tertiary over the primary hydrogen. The reaction is shown to take place by a free-radical, chain mechanism that i s initiated b y the pyrolysis of the chlorinated compounds. The dichloro cornpound is more efficient as an initiator than the trichloro compound. A theory i s proposed to explain and generalize this difference in efficiency.
T
HERE is very little in the literature on flow reactors that have been operated a t 10,000 pounds per square inch or higher (18, 26). Because this type of equipment is valuable in the development of commercial proresses, a coil-type reactor was constructed that can be operated a t 15,000 pounds per square inch and 500’ C. The operation of this reactor, which will be generally useful for the study of homogeneous reactions, was demonstrated by a study of the gas-phase alkylation of isobutane v i t h propylene a t 400” C. and up to 14,900 pounds per square inch. Petroleum alkylations are classically divided into two groupscatalytic (Friedel-Crafts) and noncatalytic (thermal) alkylation. Catalytic alkylation has been studied extensively, and its mechanism depends on carbonium ion intermediates (7, 9, 13-16, 20’24). Thermal alkylation, although not studied as extensively, is conceded to take place by a free-radical, chain mechanism * Present address, Carbide and Carbon Chemicals Go., So. Charleston,
w. Va. * Present address,
School of Chemical and Metallurgical Engineering, Purdue University, Lafayette, Ind.
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(9. 10, 13, 18) $4). The products of the two reactions are different because of the instability of the carbon skeleton of the carbonium ion intermediates ( 2 4 ) . I n 1946, O’Xelly and Sachanen (19) reported that alkylation under the influence of “homogeneous” catalysts occurs a t lower temperature:: but produces higher yields of the Same products than does the thermal reaction. The reports of investigation of the homogeneous catalytic reaction have been confined almost exclusively to the patent literature (8, 17, 26-18). There have been essentially no studies of the nature of the reaction but only investigations directed top-ard finding catalysts other than molecular halogens and halogenated hydrocarbons. Most of the work has been done a t 2000 to 6000 pounds per square inch and 400’ C. in an attempt to find a catalyst that would increase the very small (2,2,3-trimethglbuyields of a side reaction product-triptane tam). This investigation was made to extend the pressure range to 15,000 pounds per square inch and to investigate the nature of the reaction and the catalysts.
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
Vol. 46, No. 4
