The chemical compound ammonium nitrate, the nitrate of ammonia with the chemical formula NH4NO3, is a white powder at room temperature and standard pressure. It is commonly used in agriculture as a high-nitrogen fertilizer, and it has also been used as an oxidizing agent in explosives, including improvised explosive devices.
Use in industry
Ammonium nitrate is used for zeolite
modification. In ion-exchangers, UZM
zeolites have their sodium ions exchanged with the proton in NH4+
from ammonium nitrate. This forms zeolite catalysts
which have many uses in various fields, including petroleum
Use in fertilizer
The highly water-soluble salt is the preferred nitrogen source of fertilizers
. Most of the produced ammonium nitrate ends therefore in the production of fertilizers. However, the runoff of excess ammonium nitrate is a leading source of environmental waste.
Use in explosives
As a strong oxidizing agent, ammonium nitrate makes an explosive mixture when combined with a hydrocarbon
, usually diesel fuel
(oil), or sometimes kerosene
or (fine) coal dust. Ammonium nitrate and fuel oil (ANFO
) mixtures have been used for bombs
in terrorist acts such as the Oklahoma City Bombing
, because ammonium nitrate was readily available in bulk. The 34% ammonium nitrate required to make bombs is now difficult to obtain in bulk. During the troubles
in Northern Ireland
, ammonium nitrate was illegal as a result of it being used in car bombs
by paramilitary groups like the IRA
Ammonium nitrate is used in military explosives such as the daisy cutter bomb, and as a component of amatol. Military mixtures are often spiked with ~20% aluminium powder as well, increasing the blast power, but with some loss of brisance. One example of this is ammonal, which contains ammonium nitrate, trinitrotoluene (TNT) and aluminium. Aluminised mixtures are very effective under confinement, as in underwater demolition, torpedoes, and rock blasting. Very cheap water-based blasting slurries tap the power of an aluminium-water reaction with enough ammonium nitrate added to burn off the resulting hydrogen.
Ammonium nitrate is also an explosive in its purest form although it is an unusually insensitive one. Explosive properties become much more evident at elevated temperatures. When ammonium nitrate is fused and "boiled" to generate nitrous oxide, it has been claimed to be as sensitive as dynamite at the ~240 °C operating temperature.
This exothermic reaction can run away and reach detonation velocities (without proper temperature controls). The extent of this possibility has been demonstrated several times, most notably at the Ohio Chemical plant in Montreal in 1966.
Millions of pounds of relatively pure ammonium nitrate have been (accidentally) detonated when subjected to severe heat and/or shocks; see "Disasters" below. Ammonium nitrate has also found use as a solid rocket propellant, but for a while ammonium perchlorate was frequently considered preferable due to higher performance and faster burn rates. Lately, favor has been swinging back towards ammonium nitrate in rocketry, as it delivers almost as much thrust without producing an exhaust jet full of gaseous hydrogen chloride (HCl) and without the extra expense and sensitivity hazards.
Fertilizer-grade ammonium nitrate (FGAN) is manufactured in more compact form, with much lower porosity, in order to achieve more stability and less sensitivity to detonation, whereas technical grade ammonium nitrate (TGAN) prills are made to be porous for better absorption of fuel and higher reactivity.
"Diver's Liquid" is an explosive solution of ammonium nitrate in liquid ammonia. It was considered as a monopropellant for rockets, in Germany before WWII.
Ammonium nitrate is also used in instant cold packs
. In this use, ammonium nitrate is mixed with water
in an endothermic reaction
, which absorbs 25.69 kilojoules
Products of ammonium nitrate reactions are used in airbags. When sodium azide
) is used in airbags, it decomposes to Na (s) and N2
(g), the sodium forms a fine dust composed of sodium salts, which is not preferred by the airbag producers.
Ammonium nitrate is used in the treatment of some titanium ores.
Ammonium nitrate is used in the preparation of nitrous oxide (N2O):
- NH4NO3(aq) → N2O(g) + 2H2O(l)
Ammonium nitrate is used in survival kits mixed with zinc dust and ammonium chloride because it will ignite on contact with water.
Ammonium nitrate can be used to make anhydrous ammonia, a chemical often used in the production of methamphetamine.
The processes involved in the production of ammonium nitrate in industry, although simple in chemistry, challenge technology: The acid-base reaction
with nitric acid
gives a solution of ammonium nitrate: HNO3
(aq) + NH3
(g) → NH4
(aq). For industrial production, this is done using anhydrous ammonia gas and concentrated nitric acid. This reaction is violent and very exothermic. After the solution is formed, typically at about 83% concentration, the excess water is evaporated to an ammonium nitrate (AN) content of 95% to 99.9% concentration (AN melt), depending on grade. The AN melt is then made into "prills" or small beads in a spray tower, or into granules by spraying and tumbling in a rotating drum. The prills or granules may be further dried, cooled, and then coated to prevent caking. These prills or granules are the typical AN products in commerce.
The Haber process combines nitrogen and hydrogen to produce ammonia, part of which can be oxidised to nitric acid and combined with the remaining ammonia to produce the nitrate. Another production method is used in the so-called Odda process.
Transformations of the crystal states due to changing conditions (temperature, pressure) affect the physical properties of ammonium nitrate. The following crystalline states have been identified:
|| Temperature (°C)
|| Volume Change (%) |
|| - |
|| 169.6 to 125.2
|| +2.1 |
|| 125.5 to 84.2
|| -1.3 |
|| 84.2 to 32.3
|| +3.6 |
|| 32.3 to −16.8
|| −2.9 |
|| - |
The type V crystal is a quasi-cubic form which is related to caesium chloride, the nitrogens of the nitrates and the ammoniums are at the sites in a cubic array where Cs and Cl would be in the CsCl lattice. See C.S. Choi and H.J. Prask, Acta Crystallographica B, 1983, 39, 414-420.
Ammonium nitrate decomposes into gases
when heated (non-explosive reaction); however, ammonium nitrate can be induced to decompose explosively by detonation. Large stockpiles of the material can be a major fire risk due to their supporting oxidation
, and may also detonate, as happened in the Texas City disaster
of 1947, which led to major changes in the regulations for storage and handling.
There are two major classes of incidents resulting in explosions:
- In the first case, the explosion happens by the mechanism of shock to detonation transition. The initiation happens by an explosive charge going off in the mass, by the detonation of a shell thrown into the mass, or by detonation of an explosive mixture in contact with the mass. The examples are Kriewald, Morgan (present-day Sayreville, New Jersey) Oppau, Tessenderlo and Traskwood.
- In the second case, the explosion results from a fire that spreads into the ammonium nitrate itself (Texas City, Brest, Oakdale), or from a mixture of ammonium nitrate with a combustible material during the fire (Repauno, Cherokee, Nadadores). The fire must be confined at least to a degree for successful transition from a fire to an explosion (a phenomenon known as "deflagration to detonation transition", or DDT). Pure, compact AN is stable and very difficult to initiate. However, there are numerous cases when even impure AN did not explode in a fire.
Ammonium nitrate decomposes in temperatures above 210 °C. Pure AN is stable and will stop decomposing once the heat source is removed, but when catalysts are present (combustible materials, acids, metal ions, chlorides. ..) the reaction can become self-sustaining (known as self-sustaining decomposition, SSD). This is a well-known hazard with some types of NPK fertilizers, and is responsible for the loss of several cargo ships.
- Properties: UNIDO and International Fertilizer Development Center (1998), Fertilizer Manual, Kluwer Academic Publishers, ISBN 0-7923-5032-4.