Philip L. Bayless Wilmington College Wilmington. Ohio 45177
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Ammonia Synthesis A simulated laboratory research project
The usefulness of simulated experiments to extend the laboratory experience is well understood, and documented with a number of excellent examples (1). We have been seeking examples of simulated experiments which (1) are short enough to he programmed on calculators such as the Hewlett-Packard HP-65, (2) require the student to interpret the data for further interaction with the program, and (3) lead to conclusions of economic, environmental, or historical significance. This paper describes such a simulation (100 steps) of the ammonia synthesis, combined with a narrative of leadine statements and kev references. resultine in a "research" ~ r i j e c of t one to two Leek's duration. By a d justing the assignment to the student's level of accomplishment the simulation is suitable for General Chemistry, Physical Science for non-scientists, or Physical Chemistry. Students are first given some history of the ammonia svnthesis, then asked to write a research report, choosing &om a list of suggested topics. Narrative Ammonia occu~iesthird d a c e in the ~roductionof the United States chemical indukry. In 1974-31 billion pounds of liauid ammonia were svnthesized from the elements ( 2 ) . Only sulfuric acid and oxygen were produced on a larger scale. While some of the liquid ammonia is directly applied as fertilizer, most is converted into urea, ammonium nitrate, ammonium sulfate. and ammonium uhosuhate. Some . . of the ammonia is oxidized t o nitric acid (31, primarily for USQ i n the manufacture of fertilizer and ex~loaiver. Before the turn of this century Chile saltpeter was the world's principal source of nitrogen fertilizer. Chemists of the last century recognized the unlimited potential of atmospheric nitrogen for this purpose, if it could he inexpensively converted into a water soluble form, i.e., fixed (4). The efforts of Fritz Haher (5) and Walter Nernst (6) to investigate and understand the rate and equilibrium laws governing the direct synthesis of ammonia is a fascinating bit of intellectual history, all but unknown to nonchemists. Haber, who had barely passed his PhD orals in 1891, was by 1911 Director of the Kaiser Wilhelm Institute (Max Planck Institute, after WWII) in Berlin, and had been able to synthesize three ounces of ammonia per hour (7). Badische Anilin-und Soda-Fabrik organized a team of chemical engineers under the direction of Carl Bosch to scale up the production to 30 tons per day by 1913. The new high pressure technology developed by BASF has made possible a variety of other important commercial processes (8). Because the British naval blockade of Germany in 1914 cut off the supply of Chile saltpeter, Haber's ammonia process was critical to the continued production of food and explosives. It is very likely that World War I would not have lasted for four years, had Germany not possessed the world's most advanced chemical industry. Haber was appreciated in his own time; a national hero, he was awarded the Nobel Prize in 1918, as well. Yet, he later regretted his direction of chemical warfare, including battlefield supervision of the first use of poison gas. Harsh Armistice terms, fostered bv the bitterness of the drawn-out war. and incredible inflation in the post-war period, certainli contributed to the rise of Nazism. In 1933 the "Jew Haher" was 318 1 Journal of Chemical Education
exiled from Germany, and died a year later in Switzerland. Hydrogen is the economicaHy limiting reagent in the formation of ammonia, since it must be produced for mixing with atmospheric nitrogen. All the possible methods for manufacturing hydrogen have given way to "reforming" natural gas, because of the traditional cheapness of natural gas and petroleum. Current practice combines the proper proportions of natural gas, steam, and air according to CH4 + Hz0 = C0 + 3Hz CO + Hz0 = CO. + Ha is hefore the usynthesis enters the ~ h con , as liquid carsynthesis converter, and is recovered and honic acid or converted to urea. The residual mixture contains the proper 3:l ratio of hydrogen to nitrogen 3H2 + N2 = 2NH3 Modern North American agriculture is dependent upon ammonia ~roduction.and is extremelv enerev intensive (9). It has heen estimated (10) that ove; six f k o r i e s of fossil fuel energy are required to place one Calorie of food on the table. The high-yield hybrid grains of the Green Revolution are also energy-demanding, requiring large applications of fertilizers, herbicides, and irrigation. So far ammonia production has been increasing faster than the world's populatinu, but recent changes in petroleum economics will require reexamination of many pet priorities and strategies. Research Report Background of the Research (the Importance of the Ammonia Synthesis): Choose One 1) Nitrogen fixation methods 2) Nitroeen fertilizers : I , Ammmia ns an intermediate in the chemiral industy 4 , Suurvea of hydrogm for the nmmonia synthesis 5) Chemical explosives World hunger The Green Revolution Energy requirements for food production Commercial high pressure syntheses 10) History of the Haber-Bosch process 11) Equilibria of gas phase reactions Experimental Data (by HP-65 Simulation) 1) Determine the effect of pressure on yield and time to equilib. rium 2) Determine the effect of temperature on yield and time to equilibrium 3) Determine the effect of catalyst on yield and time to equiliblillrn ~ -
.
~
~
~
Determine the pressure and temperature for best yield with a reaction time between 0.01 hr and 0.1 hr Calculations: Depends on Level of Student 1) K p at several temperatures/pressures 2) Change of K, with temperature (graphical or curve-fit by HP-65) 3) Calculation of standard free energy, enthalpy, and entropy; comparison with literature values 4) Relation of K, to pressure; compressibility of gases 5) Fugaeity and activity coefficients 6) Design of an experimental apparatus to produce the data 4)
Discussion (Relation of Experimental Data to Lecture/Tert Material): Choose One 1) Le Chatelier's principle 2) Non-ideal gases 3) Intermolecular attractions and bond polarity 4) Van't Hoff or Clausius-Clapeyron equation Comments on the Program (for Instructors Only) The calculation of %NH3 a t equilihrium can he found in most physical chemistry texts. I t is assumed that Hz and Nz are introduced in a 3:l mole ratio, and remain in this ratio at equilihrium. Thus
simplified to fit into the 100 program steps of the HP-65 program card. I t does, however, give qualitatively correct results. The rate determining step of the ammonia synthesis seems to he the chemisorption of Nz on the catalyst, while the slow step in ammonia decomposition is the desorption of Nz from the catalyst surface (4). Furthermore, on a tungsten surface the ammonia decomposition is zeroorder initially, because of surface saturation (12).' This does not appear to be true for platinum (12) or iron (13). The simulation assumes zero-order kinetics. Thus ~ P N H = kdt ~
and PNHs = kt where P N H ~and t represent the equilihrium ammonia pressure and time to equilihrium. Then
and since
K,
=
PNHa
PN~'"PH~~"
the final form is
The rate constant k is temperature dependent, and was calculated by
Since the equilibrium calculations are performed using partial gas pressures, the calculations assume ideality of the gases. This is obviously a poor assumption for ammonia, and for all the reactants a t the high pressures used in the synthesis. Indeed attempts to predict accurate Kp values at very high pressures have been unsuccessful (4). The simulation uses an equation obtained by fitting experimental values of Kp (11). For example a t 30 atm the temperature dependence of KO is linear
The variation of K p with total pressure is illustrated hy the following examples
+ 1.50 X 10-3P + 5.39 X 10-8P2 (45OoC)In K, = -5.06 + 1.24 X 10-3P + 1.67 X W 7 P 2
(400°C) In K, = -4.41
The P2 term is not important below 2000 atm, so was omitted from consideration for the simulation. Reconciling the temperature dependence with the pressure dependence (trial and error) gave the simplified equation used in the simulation I n K,,=
6460
+ 1.1P -
(2)
T This equation gives agreement with experimental K p values generally within 10% relative error up to ahout 2500 atm (Table 1).One of Haber's earliest successes produced approximately 0.005% NH3 at 1000°C a t 1 atm. The simulation gives 0.0045% under the same conditions. The simulation of time to equilihrium has been greatly
(4) Ink = (-E,IRT) + 30 where En is the Arrhenius activation enerm in calories, and 30 is the natural logarithm of the pre-e;ponential factor. E. (14) is entered into the program by the student
no catalyst; E, = 80,000 Fe catalyst; E, = 50,000 W catalyst; E, = 40,000 Table 2. Ammonia Synthesis Simulation IHP-651
LBL A ST0 1
RTN LBL B 2 7 3
--
' A n mterestmg addiuon tr, the prugrnm way ruggesred hy 0. Hurherr. Culurndu (‘dirge. I f more progrnm stpps are available, t h e dcpendenrr of the z r n r - w d e r rate constant on catalyst surface area
