Explosives and Fertilizers

James Richard Fromm

Explosives, fertilizers, and matches all constitute an essentially similar group of products based upon the elements nitrogen and phosphorus. The first explosives known to the Western world date from medieval times; gunpowder and cannon were known by 1350 AD. Gunpowder consists of a mixture of potassium nitrate (saltpeter, KNO3), charcoal, and sulfur in different proportions, the modern proportions being about 7.5:1.5:1 by mass. Charcoal was obtained from wood, and sulfur was obtained from volcanic areas or by heating of pyrites in the absence of air. Potassium nitrate was at first obtained from cave efflorescences but after its use in gunpowder was known it was extracted from nitrogenous earth from stables, pigsties, pigeon lofts, and similar places. The source of the nitrogen is the excretion of animals or birds, which contains nitrogen in the form of ammonium salts, urea, or uric acid. In the presence of atmospheric oxygen, all of these are slowly oxidized to soluble nitrates. Separation was by leaching followed by evaporation; treatment with salt, lime, and potash helped to separate out other soluble salts. Nitrogenous earth remained the major source of nitrates for all purposes until the exploitation of the Chile sodium nitrate beds (1700).

The direct combination of atmospheric nitrogen and hydrogen first became practical with the development of the Haber process in 1914:

N2(g) + H2(g) rarrow.gif (63 bytes) 2NH3(g)

The Haber process uses an iron-metal oxide catalyst, temperatures of 400o-600oC, and pressures of 200-400 atm.

Ammonia dissolves readily in water, and advantage can often be taken of this to remove ammonia from gas process streams by washing. Solutions of ammonia in water are weakly basic and, as "household ammonia", find use in domestic cleaning and disinfecting. More concentrated solutions are used in the chemical laboratory. Reagent-grade ammonia solutions are colorless. The usual laboratory concentrated reagent solutions are 28% NH3 by mass; their density is 0.90 kg/L, which corresponds to a 14.8 molar aqueous solution of ammonia.

Ammonia gas, sometimes called anhydrous ammonia, and ammonia solutions have a characteristic strong irritating odor; the vapor pressure of ammonia above its concentrated solutions is high. Heating of ammonia solutions in closed containers may produce explosions due to the increase in vapor pressure. Ammonia gas is toxic and extremely irritating to the respiratory system, especially the eyes. Concentrated ammonia solutions cause skin burns and very severe eye damage and should be handled with care.

Ammonia or urea or uric acid excreted by animals is only slowly converted to nitrate by atmospheric oxygen. The commercial Ostwald process for ammonia oxidation, however, is run above 500oC in the presence of a platinum catalyst. In the absence of the catalyst, nitrogen rather than nitrogen monoxide is produced.

4NH3 + 5O2 rarrow.gif (63 bytes) 4NO + 6H2O

Nitrates can be manufactured from nitrogen monoxide by further direct oxidation:

2NO + O2 larrow.GIF (55 bytes)rarrow.gif (63 bytes) 2NO2; 2NO2 + H2O rarrow.gif (63 bytes) HNO3 + HNO2

3HNO2 rarrow.gif (63 bytes) HNO3 + 2NO + H2O

Nitric acid, like hydrochloric acid, forms an azeotrope; the nitric acid-water azeotrope contains 68.4% HNO3 and boils at 121.9oC. This azeotrope is the normal commercial form of concentrated nitric acid. Reagent-grade concentrated nitric acid is colorless. Its density is 1.41 kg/L which corresponds to a 15.4 molar aqueous solution. It is extremely corrosive, reacting to nitrate the protein of human flesh and destroy most natural and synthetic clothing fibers, and must be handled and used with great care.

Nitric acid can be prepared by reaction of sulfuric acid with a nitrate salt:

H2SO4 + 2NaNO3 rarrow.gif (63 bytes) Na2SO4 + 2HNO3

The process proceeds on gentle heating to either Na2SO4 or a mixture of NaHSO4 and Na2SO4. Very little nitric acid is still prepared by this process, because nitrogen fixation by the Haber process is much less costly.

One of the most important industrial nitrogen compounds, ammonium nitrate, is prepared by reaction of nitric acid with ammonia in an acid-base reaction:

HNO3(aq) + NH3(aq) rarrow.gif (63 bytes) NH4NO3(aq)

Ammonium nitrate, used widely as a fertilizer, is also a powerful explosive. An ammonium nitrate explosion in 1923 killed 600 people in Oppau, Germany. The detonation of a cargo of ammonium nitrate was responsible for the demolition of the port of Texas City, Texas, with 492 deaths and $50 million in damage in 1947. In these cases the ammonium nitrate acted as both oxidizing agent and reducing agent in the explosive oxidation-reduction reaction. A mixture of ammonium nitrate and a substance which can be oxidized such as fuel oil is also a powerful explosive.


Explosives other than gunpowder were not of industrial importance prior to 1846, at which time guncotton was introduced by Schonbein and nitroglycerine by the Italian chemist Ascanio Sobrero. These two explosives are similar in manufacture, being obtained upon treatment of cotton or of glycerol with a mixture of concentrated nitric and sulfuric acids. Both will easily detonate if impure, and nitroglycerine is sufficiently sensitive to shock that its use was uncommon until the Swedish chemist Alfred Nobel devised the earth-nitroglycerine mixture dynamite in 1860. Modern explosives for almost all purposes are derived from guncotton and nitroglycerine.

Concentrated nitric acid, often used together with concentrated sulfuric acid, can effect the nitration of many organic compounds; in this process, the nitro group -NO2 is added to the organic compound at a carbon atom, replacing a hydrogen. The first such reaction discovered was used for the preparation of nitrobenzene by Mitscherlich in 1834:

C6H6(l) + HNO3 rarrow.gif (63 bytes) C6H5NO2(l) + H2O

Organic nitro compounds can be powerful explosives and highly energetic, if somewhat unsafe, fuels; the military explosive trinitrotoluene, more commonly known as "TNT", is but one example. Nitration of organic compounds is a common first step in the industrial synthesis of most organic compounds containing nitrogen because the nitro group can be reduced to other nitrogen-containing groups such as amines.

Copyright 1997 James R. Fromm