|
Atomistry » Nitrogen » Ammonia » Ammonia Equilibrium | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
Atomistry » Nitrogen » Ammonia » Ammonia Equilibrium » |
Ammonia Equilibrium
As already mentioned, the passage of electric sparks, which effects an almost complete decomposition of preformed ammonia, also causes the gases to combine to a small, but definite, extent. The same is true of catalytic agents, such as finely divided iron, which accelerate not only the decomposition, but also the formation of ammonia and lead to a true equilibrium. This equilibrium may be set up at ordinary temperatures in the presence of platinum black or colloidal noble metals, when the mixture of gases is passed through acidulated water at 90° C., containing these catalysts. Or hydrogen may be passed through solutions from which nitrogen is being evolved.
It has long been known that nitrogen and hydrogen will combine in the presence of suitable catalysts, and it was suggested by Clark in 1874, that ammonia might be made by passing them over chromium, manganese, iron, or cobalt. The systematic investigation of the equilibrium proportions of ammonia formed by relatively low temperature catalysis was taken up by Haber, Perman, and their co-workers from 1904 onwards. The affinity of the formation of ammonia from its elements is positive. The reversible potential between platinum electrodes traversed by nitrogen and hydrogen respectively, which are immersed in a solution of ammonium nitrate and ammonia, is about 0.6 volts. The decomposition potential of 25 per cent, aqueous ammonia is +0.63 volts. These results indicate that the equilibrium amounts of ammonia which should be formed under reversible conditions at ordinary temperatures should be high. Heat of Formation
Ammonia is an exothermic compound, and the heat of formation under various conditions has been carefully determined, both directly and also from the alteration of the equilibrium constants, with the temperature.
The heat of formation, as determined thermochemically for ordinary temperatures, is given as follows: - N2 + 3H2 = 2NH3 + 2×11,890 calories. The heat of formation from its elements at various temperatures has been obtained by the catalytic decomposition of ammonia in a calorimeter. The heat absorbed was just compensated by electrically heating the catalyst so that the temperature remained constant. The electrical energy supplied per second is equal to the heat of formation of that amount of ammonia which is decomposed per second, viz.:
Omitting the third value, which is discordant, the heat can be expressed as a function of the temperature by the equation Q = 10,950+4.54t-0.001822t2 calories, or Q = 9,575+5.535T-0001822T2 calories, 10,950 and 9,575 calories being the heats of formation at zero centigrade and zero on the absolute scale respectively. Another equation proposed is: Q=9,465+5.819T-0.001261T2+0.06471T3 calories. The variation in the heat of the formation with the temperature is also given by the Kirchhoff Law, Qt-Q0 = ΣnCt in which ΣnC refers to the molar heats of the reacting gases (positive) and resulting gases (negative), each multiplied by the number of molecules participating in the reaction. The values of C, the molar heats, are: Nitrogen . . . 6.58+0.00053t. Hydrogen . . . 6.8+0.0003t. Ammonia . . . 8.62+0.00175t+1.7×10-6t2. In the formation of 1 mol. of ammonia, ⅔H2 + ½N2 = NH3 Hence ΣnC = 4.87 – 0.001035t – 1.7×10-62, and Q = 10,950+4.87t – 0.001035t2 - 1.7×10-63 calories. The values of Q, calculated from this, agree fairly well with those found by thermochemical methods. The variation in the constants of the reaction, viz., ![]() ![]() and, of course, also of the affinity of the reaction A = RT log KC, may be obtained as a function of the heat of reaction and specific heats, by integration of the van't Hoff isochore: ![]() Since +QV, the heat of the reaction at constant volume, is positive, the constant K must diminish with rising temperature, as is seen by the integrated form: ![]() Since the right-hand side of the equation is negative (T2 is greater than T1), there is less ammonia in the equilibrium mixture at the higher temperature. Actually the values of K determined at each temperature are those at constant pressure Kp, and log Kp = log Kc + Σnlog RT. The Q values observed are also those at constant pressure, and Qp = Qc + log RT. Q is a function of the temperature (as shown above), and its temperature coefficients must be included in the integration, giving the equation ![]() The integration constant may be evaluated by a determination of the equilibrium at one temperature, or approximately by means of the Nernst heat theorem. By the first method the constant is 2.6253. (Kp = 0.000273 at t=880° C.) Another equation, used by Haber, is - ![]() A shortened form of this equation, ![]() is sufficiently accurate for many purposes. By the use of the equations representing the most probable specific heat of ammonia at each temperature, and also those representing the variation of the equilibrium constant with the temperature, another equation has been obtained, viz.: ![]() which gives results in good agreement with those of Table. The Equilibrium Constants at Different Temperatures
Constants of the nitrogen, hydrogen, ammonia, equilibrium Experimental and Calculated Values of 104Kp at Total Pressure of 1 Atmosphere.
In Table B the temperatures are those corresponding to the stated values of K, the total pressure being 1 atmosphere and the composition of the gas 3HH2 + N2. Percentages of Ammonia Formed
Values of 104Kp at a Total Pressure of 30 Atmospheres (Measurements of Haber and le Rossignol, Haber, Tamaru, and Pommaz) (loc. cit., infra).
The calculated yields corresponding to various pressures and temperatures are summarised. Calculated Percentage of NH3 under various Temperatures and Pressures
The Experimental Determinations of the Ammonia Equilibrium
Many determinations at atmospheric pressure were carried out by Haber and co-workers. The mixture of nitrogen and hydrogen was prepared (in one series) by the combustion of air with the theoretical amount of electrolytic hydrogen. This was passed through one of a pair of tubes, each containing a catalyst plug, laid side by side in an electric furnace which maintained a constant temperature over a length of about 6 cm. The proportion of ammonia formed was estimated by absorption in standard acid. In order to approach the equilibrium from the ammonia side, a percentage, say 0.28, which was greater than the equilibrium percentage of ammonia, was added to the mixture, which was passed back through the other tube. When higher pressures were employed the mixture of gases was contained in a cylinder under a pressure of 100 atmospheres. From this, the pressure was reduced to the standard.
Catalyst bombs have been designed for experimental work of this character as well as for continuous production of ammonia. They are constructed of special materials, as, for example, tungsten steel lined with electrolytic iron, since they have to withstand not only very high pressures but also the corrosive action which ammonia has on some commercial irons. Equilibria at High Pressures
Since ammonia is formed from its constituent elements with contraction, it follows from the law of mobile equilibrium that an increase in the total pressure should increase the proportion of ammonia formed at each temperature. This result also follows from the assumption of an equilibrium constant. If the expression
![]() or what comes to the same thing, its square root, Kp, as stated above, is to remain constant when the pressure is increased, then ![]() ![]() ![]() A mixture of nitrogen and hydrogen was made by burning electrolytic hydrogen and air in the calculated proportions. The gas at 100 atmospheres was purified by a copper deoxidiser with nickel catalyst, and also by passing through soda lime, granular aluminium oxide, and fused potash, and finally through an ammonia catalyst at 500° C. The small amount of ammonia formed carried down with it traces of CO2 and H2O. From the observed percentages of ammonia it was estimated that the equilibrium constant, ![]() varies with the pressure at a single temperature. This result is not necessarily in disagreement with the law of mass, or concentration, action. For it is known that all gases at high pressures deviate markedly from the gas laws, and it is probable that the active masses calculated from the concentrations present are not identical with the real activities of the gases even at such high temperatures as 500° C. Equilibrium Constants under High Pressures.
Hence it is necessary to split up the general equations which give Kp as a function of the temperature into series, each characteristic of certain ranges of pressure. Thus: ![]()
From these results the percentages of ammonia formed at 500° C. and at different pressures have been calculated. These are as follows: -
The Effect of Velocity upon Yield
In the investigation of the time factor, a type of apparatus was used which allows the mixture of gases to be circulated at different rates, with continuous refrigeration of the product in order to separate the ammonia formed. This apparatus is similar to that employed in the large scale synthesis of ammonia. Uranium carbide was used as a catalyst, and, in the experiments of Maxted, iron potash. Although of course the maximum yield is that which corresponds to equilibrium, yet it is not desirable in practice to allow sufficient time for the equilibrium to be nearly attained, but rather to pass the gases at such a rate that the space-time-yield (S.T.Y.) is a maximum. The S.T.Y. is defined as the yield of ammonia expressed in grams per hour and per c.c. of catalyst space. The space-velocity (S.V.) is the rate of flow in litres per hour at room temperature and atmospheric pressure, per litre of catalyst space. In one series of experiments at 515° C. and p=49.6 atmospheres pressure, the following results were obtained: -
Thus, in spite of the lower percentage of ammonia formed, the space- time-yield is increased by a moderate increase in the velocity of the current of gases. The Effect of Velocity
An increase of pressure increases both the equilibrium percentage and the reaction velocity. Thus at a pressure of 113.6 atmospheres and a space-velocity of 28.5, the percentage of ammonia is 6.42 and the S.T.Y. Is 1.3, which is considerably higher than the comparable second S.T.Y. in the table. In these experiments about 50 per cent, of the total possible percentage of ammonia was formed by a contact of one to two seconds.
The Catalysts
Among the catalysts which induce a sufficient velocity at about 500° C. (or even below this temperature) are osmium, uranium, uranium carbide, iron, and metals of the iron group, or those similar to iron in their physical properties.
OsmiumOsmium was used in the early experiments of Haber and le Rossignol. It may be in the form of grains, or asbestos may be soaked in an osmium salt, which is then reduced to the finely divided metal. It is too expensive for technical use.UraniumThe commercial metal may be broken into small pieces, or it may be prepared by the reduction of green uranium oxide with sugar charcoal. These catalysts are easily oxidised by traces of air or water-vapour. Hence they are unsuitable for continuous use. The same may be said of sodamide,. which was used in some early experiments, including those in connection with the American plant at Muscle Shoals. Uranium may also be used in the form of the carbide, which is prepared in the electric furnace from the oxide and carbon.CeriumThe oxide is reduced with magnesium in an atmosphere of hydrogen.TungstenThe commercial metal is heated in chlorine, and the chloride is reduced by gaseous ammonia.ManganeseManganese in a state of fine division may be prepared from the amalgam. The high initial activity of this catalyst soon falls to a lower constant value.Magnesium and Beryllium CyanamidesMagnesium and Beryllium Cyanamides may be made by heating the carbonates to 500° C. in ammonia. They are improved by an admixture of iron.IronThis is the catalyst which has been most widely adopted, either alone, or with another metal of the iron group, or with Cr, Mo, W, or with promoters (vide infra). Iron may be prepared in a very active form by oxidation in a furnace heated with oxyhydrogen flames. The oxide is broken up and reduced by hydrogen or ammonia at 500° to 600° C. It can be freed from catalyst poisons by alternate oxidation and reduction. Iron with molybdenum has been largely used by the Badische Company.The use of "promoters" with the iron has been found advantageous. These are oxides of the alkaline earth or rare earth metals, alumina, silica, and potassium aluminate. The yield is much improved if two promoters are present, of which one is a basic oxide, such as those of potassium and caesium, and the other a more acidic oxide, such as those of alumina or zirconia. Thus it was found that iron with alumina alone gave a yield of 8 per cent., with potassium oxide alone 5 per cent., but with both together 14 per cent, of ammonia. The manufacture of such a catalyst proceeds in two stages. In the first, iron oxide is fused between water-cooled iron electrodes on a hearth of the oxide. The fused material is mixed with the promoters and reduced in a current of hydrogen. Some of the modern catalysts are so effective that they allow a sufficient velocity of combination at temperatures slightly above 300° C. Catalyst Bombs
In the Claude bomb (fig) the gases are heated by passing from A round the contact mass, then through this and out at C. The tubes M and T are made of nickel-chromium alloy. Each bomb weighs 15 cwt., and twenty-four of these, weighing 18 tons, are required to make 20 tons of ammonia per day. The reacted gas passes through a coil cooled externally by water, and all the ammonia, except 2 or 3 per cent., is liquefied at the enormous pressure, 900 atmospheres. Fig. 17, the Casale bomb,1 illustrates some details of heat exchange. The tube 4, with flanged ends 13, is separated from its thin lining 6 by a heat insulator 8. The flanges are closed by plates, of which the lower carries the electrical resistance heater 7 and its electrical connections 11, and the upper the inflow and outflow tubes. The mixture of gases enters the annular space between the thin inner tube 1 and the thick outer tube 17, thus being heated by the reacted mixture which is passing out through 1. The entering gases pass through 9, 10, and 2 to the electrical heater 7, and thence through the perforated plates 18, the catalyst material, the holes 3, and the passages 16 and 15, and out through 1. |
Last articlesMg in 4Y52Mg in 4Y30 Mg in 4Y2V Mg in 4Y2X Mg in 4Y2Y Mg in 4Y2U Mg in 4Y2T Mg in 4Y2Q Mg in 4Y2R Mg in 4Y2S |
© Copyright 2008-2020 by atomistry.com | ||
Home | Site Map | Copyright | Contact us | Privacy |