ammonium nitrate

Ammonium nitrate
IUPAC name	Ammonium nitrate
Identifiers
CAS number	6484-52-2 Y
UN number	0222 – with > 0.2% combustible substances 1942 – with <= 0.2% combustible substances 2067 – fertilizers 2426 – liquid
RTECS number	BR9050000
Properties
Molecular formula	(NH4)(NO3)
Molar mass	80.043 g/mol
Appearance	white solid
Density	1.725 g/cm3 (20 °C)
Melting point	169.6 °C
Boiling point	approx. 210 °C decomp.
Solubility in water	118 g/100 ml (0 °C) 150 g/100 ml (20 °C) 297 g/100 ml (40 °C) 410 g/100 ml (60 °C) 576 g/100 ml (80 °C) 1024 g/100 ml (100 °C) [1]
Explosive data
Shock sensitivity	very low
Friction sensitivity	very low
Explosive velocity	5270 m/s
Hazards
MSDS	ICSC 0216
EU Index	not listed
Main hazards	Explosive
NFPA 704	0 2 3
LD50	2085–5300 mg/kg (oral in rats, mice)[2]
Related compounds
Other anions	Ammonium nitrite
Other cations	Sodium nitrate Potassium nitrate Hydroxylammonium nitrate
Related compounds	Ammonium perchlorate
Y (what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

The chemical compound ammonium nitrate, the nitrate of ammonia with the chemical formula NH₄NO₃, is a white crystalline solid 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.

Ammonium nitrate is used in cold packs, as hydrating the salt is an endothermic process.

Production

The processes involved in the production of ammonium nitrate in industry, although chemically simple, are technologically challenging. The acid-base reaction of ammonia with nitric acid gives a solution of ammonium nitrate:^[3] HNO₃(aq) + NH₃(g) → NH₄NO₃(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.

Ammonium nitrate is also manufactured by amateur explosive enthusiasts by metathesis reactions:

(NH₄)₂SO₄ + 2 NaNO₃ → Na₂SO₄ + 2 NH₄NO₃

Ca(NO₃)₂ + (NH₄)₂SO₄ → 2 NH₄NO₃ + CaSO₄

Sodium sulfate is removed by lowering the temperature of the mixture. Since sodium sulfate is much less water-soluble than ammonium nitrate, it precipitates, and may be filtered off. For the reaction with calcium nitrate, the calcium sulfate generated is quite insoluble, even at room temperature.

Crystalline phases

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:

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.

Disasters

Ammonium nitrate decomposes into gases including oxygen 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, and there are numerous cases when even impure AN did not explode in a fire.

Ammonium nitrate based explosives were used in the Oklahoma City bombing.

Ammonium nitrate decomposes in temperatures normally well above 200°C. However the presence of impurities (organic and/or inorganic) will often reduce the temperature point when heat is being generated. Once the AN has started to decompose then a runaway reaction will normally occur as the heat of decomposition is very large. AN evolves so much heat that this runaway reaction is not normally possible to stop. This is a well-known hazard with some types of N-P-K Fertilizers, and is responsible for the loss of several cargo ships.

References

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External links

• International Chemical Safety Card 0216
• "Storing and Handling Ammonium Nitrate", UK Health and Safety Executive publication INDG230 (1986)
• Specific Heat Capacity as function of temperature - on-line calculation