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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.:

t° C0466503554659
Heat of formation, Q10,95012,67012,70012,90013,150

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,
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,

H2 + ½N2 = NH3
ΣnC = 4.87 – 0.001035t – 1.7×10-62,
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

equilibrium constant of the N2, H2, NH3
Variation in the equilibrium constant of the N2, H2, NH3, reactions with temperature. Total pressure = 1 atm.
The following tables contain the most important of the experimentally determined values of Kp, which are compared with those calculated by equations (1) and (2).

Constants of the nitrogen, hydrogen, ammonia, equilibrium Experimental and Calculated Values of 104Kp at Total Pressure of 1 Atmosphere.

Temperature, °C104Kp (Haber and le Rossignol.)104Kp (Haber and Maschke.)Calculated by Equation (2)Calculated by Equation (1)
600. . .13.815.214.9
7504.68. . .4.984.88
8003.34. . .3.623.59
9301.89. . .1.841.81
1100. . .0.930.930.89

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

% NH30.70.320. 0290.0250.0230.0190.0160.0130 00960 00650.0036

Values of 104Kp at a Total Pressure of 30 Atmospheres (Measurements of Haber and le Rossignol, Haber, Tamaru, and Pommaz) (loc. cit., infra).

t° C.561620631700704710722801812901914952974

equilibrium N2, H2, NH3
Equilibrium per cent, NH3 in N2 + 3H2 at different pressures
The value of the constant should be independent of the total pressures if these are not so high that the participating gases depart widely from the gas laws. Actually there is a fair agreement between Tables C and A. These equilibria were attained in the presence of iron, manganese, nickel, or chromium.

The calculated yields corresponding to various pressures and temperatures are summarised.

Calculated Percentage of NH3 under various Temperatures and Pressures

Temperature, ° CPercentage of NH3.
1 Atm.30 Atm.100 Atm.200 Atm.

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 must increase at the expense of and . This relation is shown by Table, which is based on the constancy of Kp. As this increase is advantageous, the values of the constants at high pressures have been investigated experimentally and the relation has been verified.

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.

Temperature, °C300 Atm600 Atm1000 Atm.
KPer cent. NH3KPer cent. NH3KPer cent. NH3
5000.0049826.20.0065142.1. . .. . .

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:

p in Atmαβγ×104δ×107c

From these results the percentages of ammonia formed at 500° C. and at different pressures have been calculated. These are as follows: -

p in atm1030501003006001000
NH3 per cent.1.213.495.5610.6126.4442.1557.47

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: -

Time of contact (secs.)86.026.710.14.3
NH3 per cent3.933.382.772.0

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.


Osmium 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.


The 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.


The oxide is reduced with magnesium in an atmosphere of hydrogen.


The commercial metal is heated in chlorine, and the chloride is reduced by gaseous ammonia.


Manganese 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 Cyanamides

Magnesium and Beryllium Cyanamides may be made by heating the carbonates to 500° C. in ammonia. They are improved by an admixture of iron.


This 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

synthesis of ammonia
Cycle of operations in the synthesis of ammonia (Halber and le Rossingol)
The simplified sketch of the apparatus used by Haber and le Rossignol will give a general idea of the manner in which the combination is effected. The mixture N2+3H2 enters at A, is dried by the soda-lime tube B, and passes over the outside of a bundle of steel capillary tubes C. It is thus heated up to 400°-500° C. by the reacted gases which are flowing inside from the contact mass D, and have a temperature of 500° to 600° C. The entering gases are further heated to 600°-1000° C. by the electrically heated coil E, then pass through the catalyst 1), the capillary tubes C, the connecting tube G, to the pump P, in which they are compressed to 200 atmospheres, and then pass through the inside of the small bore metal tubes H, which are cooled outside by the gas escaping from the liquid ammonia. The ammonia in the reacted gases then liquefies in the vessel I, the outside of which is kept at - 60° to - 70° C. by a freezing mixture. The uncombined nitrogen and hydrogen are circulated as before.

Catalyst bombs
Catalyst bombs, Haber, Claude and Casale.
The Haber-Bosch bomb as used at Oppau consists of two forgings of tungsten steel held together by bolts (centre of fig). The flanged ends are closed by covers, top and bottom. The over-all dimensions of this bomb are: length 42 ft. 8 in., largest diameter 5 ft. 5 in., and the weight is 74½ tons. It is lined with electrolytic iron; inside this is refractory material, and inside this another lining. The internal space, of diameter 2 ft. 7½ in., is occupied by the catalyst. An electrical resistance heater raises the temperature initially to 600° C., and then this temperature is largely maintained by the heat of the reaction, conserved by good heat insulation. There is a heat interchanger to reduce the temperature of the issuing gases so that the ammonia at 200 atmospheres can be absorbed in water, giving a 20 per cent, solution. The production is 20 tons of ammonia per day.

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.

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