REACTION BETWEEN GASEOUS AMMONIA AND H2SO4 IN AEROSOL DROPLETS
April, 1968
469
The results are shown in Table I.
presumed t o be adequate for orbital contributions. From these normal moments, one can conclude TABLE I that copper(I1) acetate dissociates into single THE MOLARMAGNETICSUSCEPTIBILITIES A N D THE EFmolecules in water and pyridine. The molecules FECTIVE MAGNETIC MOMENTSI N BOHR MAGNETONS OF of water and pyridine capable of being readily coCOPPER(II)ACETATE IN VARIOUS SOLVENTS ordinated on metal atoms as ligands are considered Solvents 1, ~(7%) XM X 106 fi(9.h.I.) Color to occupy appropriate positions about copper Water 23.5 0 . 3 - 1 . 4 1491 1.93 Blue atoms, thereby breaking the weak copper-copper Pyridine 23 .2- . 5 1330 1.80 Blue bonds in the dimer molecules. I n methanol soluMethanol 28 .1- . 3 1019 1 . 5 8 Green tion, the magnetic moment assumes an intermediate Ethanol 22 .I- . G 826.5 1 . 4 3 Green value between those in water and ethanol. Dioxane 26 .1- . 3 740 G 1.37 Green The color of the solutions of copper(I1) acetate (Solid) 22 ..... 807.5 1 . 4 1 Green in those solvents in which the copper salt shows a normal magnetic moment was blue,. while the Discussion solutions in which subnormal magnetlc moments The molar magnetic susceptibility of copper(I1) were observed assumed a green color. Tsuchida acetate depends to a considerable extent upon the and Yamadag have found that copper(I1) acetate kind of solvents in which it is dissolved. The and propionate in ethanol or chloroform showed a values of effective magnetic moments in ethanol new absorption band at 80 X l O I 3 sec.-l in addiand dioxane are practically identical with that in tion t o that a t 43 X 10’3 sec.-l commonly ob.the solid state, indicating the presence of the solute served for copper salts. They attributed the new in dimer molecules as in crystals. On the other band t o the presence of a copper-copper link in hand, the moments in aqueous and pyridine solu- such entities as Cuz(CH3COO)4X2or Cuz(CzH6tions are slightly greater than the spin-only value COO),X2, where X denotes HzO, C2H50Hor CHC13. 1.73 Bohr magnetons for a single odd electron. The present magnetic measurements provide a Since a cupric ion has a more than half filled 3d more direct and definite evidence in support of shell, the orbital contribution to the effective mo- their conclusion. ment adds to rather than subtracts from the spin (9) R. Tsuohida and S. Yamada, Nature, 176, 1171 (1955); R. moment and the differences between the observed Tauchida, S. Yamada and H. Nakamura, i b i d , , 178, 1192 (1956); moments and the theoretical spin-only moment are Bull. Chem. S O C .J a p a n , SO, 953 (1957). O C .
KINETICS OF THE REACTION BETWEEN GASEOUS AMMONIA AND SULFURIC ACID DROPLETS IN AN AEROSOL’ BY R. C. ROBBINSAND R. D. CADLE Contribution from Stanford Research Institute, Menlo Park, California Received December 10, 1967
The rates of reaction in the system nitrogen, gaseous ammonia and sulfuric acid droplets have been studied a t 28 and 8”. Initial rates are first order with respect to ammonia and droplet surface. The over-all reaction appears t o be limited by the diffusion of product into the droplets, and an empirical rate equation for the over-all reaction has been obtained. The results are discussed in terms of collision and absolute rate theory.
Introduction I n spite of the large body of literature concerned with heterogeneous reactions, little information is available concerning chemical reactions in aerosol systems. Perhaps the most pertinent studies are those by Lewis and co-workers on reactions in fluidized.beds2v3and studies of aerosol growth, such as those by La Mer and Cotson4 and La Mer and This paper reports a study of the reaction between gaseous ammonia and sulfuric acid droplets in an aerosol. This system was chosen because the reaction would be expected to be rapid (1) This work was supported by Grant 5-30 of the Public Health Service. (2) W. K. Lewis, E. R. Gilliland and W. A. Reed, Ind. Eng. Chern., 41, 1227 (1949). (3) W. IC. Lewis, E. R . Gilliland and G. T. MoBride, Jr., ibid., 41, 1213 (1949). (4) V. K. La Mer and Sidney Cotson, Science. 118, 516 (1953). (5) V. K. La Mer and R . Gruen, Trans. Faraday S O C . ,48, 410 (1952).
and the system can be prepared readily and reproducibly. Experimental The aerosols were prepared from 98.3y0 sulfuric acid, using a condensation-type generator; particle sizes were determined with an “ 0 ~ 1 . ” 6 Particle concentrations were monitored with a Sinclair-Phoenix forward-scattering Tyndallometer. The nitrogen was dried by passing it through a tube containing phosphorus pentoxide. Nitrogenammonia mixtures were prepared in a conventional gashandling system. Reaction rates were measured using a flow system. The reaction cells were vertical glass U-tubes, 20 mm. inside diameter, and varying in length from a few centimeters to approximately two meters. The aerosol and the nitrogenammonia mixture were introduced into the bottom of one arm of the cell from opposite directions so that turbulent mixing occurred. The Reynolds number in !he cell was about 250 and the flow became laminar immediately above the mixing zone. The mixture flowing from the other end (6)
V. K. La
471 (19.50).
Mer, E. C. Y. Inn and I. B. Wilson, J . Coll. Sci., ti,
470
R. C. ROBBINS AND R. D. CADLE
Vol. 62
to ammonia and 0.99 with respect to droplet surface. Thus the order of the initial reaction was essentially two, and one with respect to each reactant. Initial rates can be represented by the equation _ -da= k [",lo dt
TIME - m c . Fig. 1.-Reaction between ammonia (394 pmoles/m.s) and sulfuric acid droplets of various diameters (19.7 pmolesl m.3) a t 28'. All reactions were in nitrogen except for the reaction of 0.6 p diameter droplets, which was studied in nitrogen ( A ) and in the nitrogen-helium mixture (v). Additional points, which are not shown, were obtained for later times.
of the cell passed through a bed of pelleted molecular sieve material,7 about 5 cm. deep, which removed the ammonia without removing the aerosol. The particulate material was then collected from the aerosol with a Mine Safety Appliance electrostatic precipitator and analyzed for ammonium ion by titration with a dilute aqueous solution of methyl purple. The analytical method was corroborated by the colorimetric Nessler method for ammonia.8 The gas velocity in the reaction cells was 19 cm./sec. The variation of fraction completion of the reaction with time was investigated by using several cells of different lengths for each set of conditions. A few experiments were made in which most of the nitrogen (all except that used for generating the aerosol) was replaced with helium. The resulting concentration of helium in the reaction cells was 62.5 mole %. Particle diameters of the sulfuric acid droplets ranged from 0.2 to 0.9 p . Concentrations were of the order of micromoles of sulfuric acid and ammonia per mole of total aerosol. Most of the experiments were undertaken a t 28 f 1'. A few experiments were made at 8 f 0.1' by placing the entire apparatus in a large thermostated refrigerator equipped with blowers.
where a is the concentration of acid in the aerosol expressed in weight or volume units, r is the droplet radius, and the subscript 0 refers to initial conditions. The average second-order velocity constant, k, was 7.6 X lo4 mole-' set.-' with a standard deviation of 0.92 X lo4. Changing from an all-nitrogen atmosphere to one containing 62.5 mole yc of helium had no measurable effect on the reaction rates (Fig. 1). Decreasing the temperature from 28 to 8" slightly decreased the velocity constant. The usual Arrhenius-type plot led to an activation energy of about 3 kcal. mole-'. Discussion Two facts suggest that gas-phase diffusion of ammonia did not control the over-all rates. These are (1) that even when a large excess of ammonia was present the rates were not constant over an appreciable period of time, and ( 2 ) that substituting helium for nitrogen had no appreciable effect on the rates. The fact that the over-all rates had no particular order suggests that these rates were controlled by diffusion of reaction product in the droplets. The over-all reaction rate can be represented by the equation
- da
=k
3ao
7(1 - Fx) [NH,]
where x is the fraction of acid reacted and F is a dimensionless multiplier which allows for the rate of Results diffusion of products and for the surface area Two sets of runs were made to determine the or- represented by each unit of product concentration. der of the reaction. One set involved maintaining F was calculated from the rate data for various the initial concentration of acid in the aerosol con-. values of x and various initial conditions, and was stant a t 19.7 pmoles/m.a, and the droplet size con- found to be a single-valued function of x within the stant at 0.6 p while varying the initial stoichio- range of conditions studied. By substituting this metric ratio (equiv. H2S04/equiv. NH3) between function of x in the above equation, the empirical 0.48 and 0.1. The other set involved maintaining rate equation representing the over-all reaction was the acid concentration a t 19.7 p m o l e ~ / m .and ~ the found to be initial stoichiometric ratio at 0.1 while varying 1-x - da the particle diameter between 0.2 and 0.9 p . = k 3 3 (0.18 0.18 + 0.822 ["a1 Typical results are shown in Fig. 1 as yo aerosol It is of interest to consider the results for the reacted plotted against time. The slopes of these curves were measured a t several points and plotted initial reaction in terms of collision and absolute on log paper against corresponding concentrations rate theory. Calculations based on the kinetic of gases and the empirical value of at 28" of sulfuric acid. Curved rather than straight lines theory indicate that the fraction of collisions between amwere obtained, which showed that the over-all remonia molecules and acid surface resulting in reaction had no particular order. The order of the initial reaction was e3timated by plotting log initial action is about 0.1. If the velocity constant is rates against log initial concentrations of one re- written as h: = PZ exp(-EIRT), the steric factor E = 0, and has a maximum value of 1 actant when the initial concentration of the other P is 0.1 for The empirical value of for E = 1.4 kc.al. mole-I. reactant was held constant. The slopes of the straight lines fitted to the points by the method of E (3 kcal. mole-I) is in fair agreement with these least squares were the orders with respect to the results, considering uncertainties in the calculation variable reactant. These were 0.96 with respect of the collision number Z and the experimental determination of E. Order of magnitude values for (7) An artificial zeolite nianufactured by the Linde Company. the pre-exponential factor PZ were calculated by (8) M. B. Jacobs, "Analytical Chemistry of Industrial Poisons, Partition functions Hazards, and Solvents," Interscience Publishers, New York, N. Y., using absolute rate theory. for gaseous ammonia, liquid sulfuric acid, and an 1941, pp. 289-291.
)
.
April, 1958
ADSORPTION O F
HYDROGEN ON NICKEL-KIESELGUHR
/H-N
activated complex of skeleton structure S--0 were calculated, using the method of Wilson and Johnstong for estimating the parameters of the activated complex. The resulting value of lo2 mole-l see.-' is too low by at least two orders of magnitude. This suggests the interesting pos(9) D. J. Wilson and H. S. Johnston, J . A m . Chem. Soc., 79, 29 (1Y57).
47 1
sibility that the ammonia is physically dissolved in or adsorbed on the sulfuric acid droplet before reaction occurs, restricting the rotational and translational motion of the ammonia molecule, thus reducing the magnitude of the partition function and increasing the value of PZ. Present knowledge of the theory of liquids is inadequate to permit calculation of the magnitude of the partition function of ammonia in such states.
THE ADSORPTION OF HYDROGEN ON NICKEL-KIESELGUHR BY L. LEIBOWITZ,' M.J. D. Low'
H. AUSTINTAYLOR Nichols Laboratory, New Yo& University, New York 55, N . Y. AND
Received December 1 1 , 1957
The rates of chemisorption of hydrogen on a nickel-kieselguhr surface have been studied over a pressure range from about 1 cm. to 1 atm. in the temperature range 160 t o 325'. The parameters LY and -log to of the Elovich equation are approximately linearly proportional to the initial pressure and linear in the reciprocal temperature. New techniques are described which demonstrate the existence of site production and of site decay, proving the latter to be independent of the prevailing pressure during adsorption. The results are interpreted on the basis of the Taylor-Thon mechanism of chemisorption.
The generality of the applicability of the Elovich equation to cheinisorption rates2 and its iiiterpretation from several different and mutually exclusive points of vie^,^^^ require an investigation not only of the dependence of the parameters in the equation on the prevailing conditions but, what is even more important, a departure from the classical chemisorption procedures to discover factors of significance previously overlooked. The system hydrogen on nickel-kieselguhr was chosen since there are no data in the literature reporting rate measurements at high temperatures. Sadek and Taylor4have measured some adsorption rates in the temperature interval - 126 to -78". The adsorption isobar shows a very rapid decrease with increasing temperature between - 195 and -126" where the adsorption is reversible. Between -126 and -78" the volume adsorbed increases slightly, by less than 20%, reaching a flat maximum around 0", decreasing steadily thereafter. Previous determinations of the energy of activation of adsorption have usually been made in the region where the isobar for the particular system is increasing with increasing temperature. The measurements reported here cover the temperature range 160 to 325" wherein the isobar is markedly decreasing. To investigate the dependence of the Elovich parameters on pressure, a pressure range from about 1 em. to one atmosphere has been studied. Subsequently in the paper, a number of departures from the usual procedure are described which must cast doubt on the classical interpretations of chemisorption. (1) Abstracts from dissertations ia the Department of Chemistry Rubmitted t o the faculty of the Graduate School of Arts and Science in partial fulfillment of the requirements for the degree of Doctor of Philosophy a t New York University, 1956. ( 2 ) H. A. Taylor and Pi. Thon, J. Am. Chem. Soc., 74, 4169 (l95a. (3) G. S. Porter and F. C. Tompkins, Proc. R o y . Soc. (London), 217A, 529, 544 (1953); H. J. Engell and K. Hauffe, 2. Eleklrochem., 67, 762, 733. 776 (1953); P. T. Landsberg, J. Chem. Phys., 23, 1079 (1955); T. J. Jennings and F. 8. Stone, paper presented a t The International Congress on Catalysis, Sept. 1956. (4) H. Sadek and €1. S. Taylor, J . A m . C h s m . S O L ,7 2 , 1168 (1950).
Experimental The measurements of rates of adsorption were carried out by admitting a known amount of hydrogen to a sample of the catalyst which had been suitably evacuated and measuring the rate of change of pressure in the system. Catalyst.-To 75 g. of kieselguhr in 500 ml. of boiling water was added slowly a solution containing 60 g. of Ni(N03)2,6HzO in 300 ml. of water. The solution was boiled for 30 min. and (NH4)?COIsolution was added in excess to precipitate NiC03. The mixture was boiled for 2 hr. and then filtered by suction through sintered glass. The precipitate, washed ten times with 200-ml. portiona of water, was sucked dry, heated overnight a t l l O o , crushed and sieved, the material between 10 and 100 mesh being retained for use. Duplicate analyses showed 11.5% Ni. A 10.17-g. portion of the carbonate preparation was reduced in a 5-10 l./min. stream of tank hydrogen, which had been passed over copper a t 400°, a liquid nitrogen trap, Ascarite and Drierite, a t a pressure of 70 cm. for about 20 hours a t 350". The pressure of the hydrogen stream was reduced t o 10 cm. and reduction continued for 45 hr. A Drierite tube attached to the exit of the catalyst vessel showed no gain in weight in 44 hr. The hydrogen was flushed out with helium and the exit tube sealed off. The catalyst chamber was connected to a source of hydrogen and to two manometers, one containing dibutyl phthalate (D.B.P.), the other, mercury, in such a way that the D.B.P. manometer could be used for measuring the pressure changes in the system a t any total pressure measured on the mercury manometer. A conventional pumping system of a mercury diffusion pump backed by a Megavac pump was used. Procedure.-The catalyst was exhausted prior to a run by evacuation a t 375". After an initial evacuation for 15 or 20 min., helium was introduced to a pressure of 60 cm. and allowed to remain in contact with the catalyst for about 10 min. Evacuation was then resumed and continued for about 10 hr., when the remaining pressure was less than 10-6 mm. The flushing with helium was found t o reduce the time required for evacuation from 25 to 10 hr . The furnace temperature was reduced to the desired value and hydrogen admitted to a pressure exerted on the two arms of the DBP manometer. A stopcock then isolated these two arms and the pressure changes were read from the change in the levels in the two arms. All the data reported were carried out on the same sample of catalyst. However, mechanical failures on three occasions permitted the catalyst to be exposed to air. The catalyst was restored by allowing it to stand for about 24 hr. in contact with hydrogen a t 1 atm. and 375". After such n treatment the catalyst surface was somewhat changed and tho adsorption rates were not identical with those obtained
