Ammonium nitrate decomposition temperature

According to Rosser, Inami and Wise [741 ammonium bichromate, soluble in molten ammonium nitrate, was found to bo a catalyst of the decomposition of ammonium nitrate at temperatures 460-520 K. For low concentration of the catalyst the principal products of decomposition were Nj. N2O, HjO. HN03. The catalysed decomposition was inlubited by ammonia and water, promoted by nitric acid. 

Write formula unit equations for the following (a) thermal decomposition of potassium azide, KN3 (b) reaction of gaseous ammonia with gaseous HCl (c) reaction of aqueous ammonia with aqueous HCl (d) thermal decomposition of ammonium nitrate at temperature above 260°C (e) reaction of ammonia with oxygen in the presence of red hot platinum catalyst (f) thermal decomposition of nitrous oxide (dinitrogen oxide), N2O (g) reaction of NO2 with water. 

Ammonium nitrate decomposes into nitrous oxide and water. In the solid phase, decomposition begins at about I50°C (302°F) but becomes extensive only above the melting point (I70°C) (338°F). The reaction is first-order, with activation energy about 40 kcal/g mol (72,000 Btii/lb mol). Traces of moisture and Cr lower the decomposition temperature thoroughly dried material has been kept at 300°C (572°F). All oxides of nitrogen, as well as oxygen and nitrogen, have been detected in decompositions of nitrates. 

In the context of safety of the process of neutralisation of nitric acid with ammonia, the effects of temperature (160-230°C), pressure (2.3-9.8 bar), and concentrations of ammonium nitrate (86-94%) and of nitric acid (0-4%) upon decomposition rate were studied. 

A violent explosion in an ammonium nitrate store with sawdust-covered floors (to absorb spillage and prevent sparks) was attributed to local decomposition of the moist nitrate-containing sawdust, leading to temperature rise and spontaneous ignition. The observation of red-brown fumes just before the explosion supports this hypothesis. 

A lire disaster costing 67 million occurred in Texas City, Texas, on the SS Grandcamp (16 April 1987) due to spontaneous ignition of stored fertilizer in the ship s hold. A release of steam from an engine leak caused the atmosphere of the ammonium nitrate fertilizer to be exposed to temperatures of 100 °C. Ammonium nitrate (NH4NO3) decomposes exothermically releasing 378 kJ/g mol. Its rate of decomposition can be described by the Arrhenius equation ... 

The reported decomposition of ammonium nitrate indicates that the reaction is unimolecular and that the rate constant has an A factor of 1013 8 and an activation energy of 170kJ/mol. Using this information, determine the critical storage radius at 160°C. Report the calculation so that a plot of rcrit versus T0 can be obtained. Take a temperature range from 80°C to 320°C. 

DSC and DTA measurements show melting of ADN, NH4N(N02)2, at 328 K, the onset of decomposition at 421 K, and an exothermic peak at 457 K.l l Gasification of 30% of the mass of ADN occurs helow the exothermic peak temperature, and the remaining 70% decomposes after the peak temperature. The decomposition is initiated by dissociation into ammonia and hydrogen dinitramide. The hydrogen dinitramide further decomposes to ammonium nitrate and NjO. The final decomposition products in the temperature range 400-500 K are NH3, HjO, NO,... 

Aqueous solutions of ammonium nitrate undergo a double decomposition reaction with metal salts. NH4NO3 acts as an oxidizing agent in aqueous solutions and is reduced by various metals at ambient temperatures. 

The decomposition of aqueous ammonium nitrate at elevated temperatures and pressures occurs and is a function of chloride, nitrate, and total acidity. Catalysis requiring both chloride and acid was observed in solutions containing 20% (w/w)... 

It is worth emphasizing that the reaction scheme above is able to explain not only the stoichiometry of the fast SCR reaction, and specifically the optimal equimolar NO to N02 feed ratio, but also the selectivity to all of the observed products, namely N2, NH4NO3 and N20, which derives from thermal decomposition of ammonium nitrate (Ciardelli et al., 2004b, 2007a Nova et al., 2006b) furthermore it is in agreement with the observed kinetics of the fast SCR reactions, which at low temperature is limited by the rate of the reaction between nitrate and NO. 

Chretien and Woringer [34] described the preparation of silver cyanamide from calcium cyanamide by the action of silver nitrate and also described its explosive properties. Montagu-Pollock [35] described a method for growing large crystals of the salt from its aqueous solution in the presence of ammonium nitrate, ammonia and a surface active agent. Bowden and Montagu-Pollock [36] and Montagu-Pollock [35] studied the slow decomposition of the crystals when heated at temperatures from 150 to 360°C. The course of decomposition was studied by electron microscope. 

The effect of solid particles. In his studies of the detonation products of explosives stemmed in different ways, Audibert [29, 32] paid attention to the fact that in charges insufficiently stemmed a certain amount of the explosive remains in the form of small particles which may undergo further explosive decomposition according to conditions, i.e., temperature and ambient pressure. If these particles are ejected into a space filled with a methane-air mixture they may lead to the explosion of this mixture. The possibility of the existence of particles of undecomposed explosive in ejected detonation products has been disputed by some authors (Segay [33]), but many others have proved that it can occur. T. Urbanski [34] found that a thin layer on the periphery of a cylindrical charge of ammonium nitrate explosive is... 

Attention was paid to the explosive properties of ammonium nitrate as early as 1883 by Berthelot [10] who first formulated the equation of decomposition, and gave the numerical data for the heat of explosion, heat of formation, the volume of gases evolved and the temperature of explosion. 

Wood and Wise [13] examined the rate of decomposition of ammonium nitrate enclosed in sealed ampoules and kept at temperatures between 200 and 300°C. They found the energy of activation E — 31.4 kcal/mole. 

Cook and A. Taylor [14] as well as Guiochon and L. Jacque [15] used a thermo-gravimetric method to study the decomposition of ammonium nitrate between the temperatures 217-267 and 180-280°C respectively. They found the energy of activation was 38.3 and 36.5 kcal/mole respectively. 

The increase in the rate of decomposition under the influence of added Cr203 and K2Cr207 is shown on Figs. 171 and 172 respectively.lt has been found that an increase of the rate occurred only at temperatures near to the melting point of the samples. Guiochon [17] also found that cobalt salts increase the rate of thermolysis of ammonium nitrate. Other mineral salts (of manganese, nickel and copper) have a similar but much weaker action. A large number of salts of other metals are without any noticeable action. 

Thermal decomposition of ammonium nitrate can also be facilitated by adding organic compounds. Thus, ammonium nitrate mixed with cellulose begins to decompose at 100°C and decomposition becomes distinctly perceptible at 120°C. Also salts of some organic bases (e.g. pyridine nitrate) considerably lower the temperature of decomposition of ammonium nitrate. 

Pure ammonium nitrate when hermetically confined can be made to explode by rapid heating to a temperature above 200°C. Thus, for example Herguet [34] reports that ammonium nitrate confined hermetically undergoes an explosive decomposition on being heated to a temperature of 260-280°C. On the other hand attempts to bring about an explosion of ammonium nitrate that was not perfectly confined failed. Sherrick [35] has established that ammonium nitrate non-hermetically confined does not decompose explosively due to a thermal reaction, but will do so if brought in contact with molten iron. 

We have investigated the humification of straw, as an example, under constant conditions of humidity and temperature in a climatic chamber (4,18) and separated different fractions according to a modified method of Waksman s proximate analysis. The amount of nitrogen present in these processes is the factor limiting the rate of decomposition. Therefore we added nitrogen in form of ammonium nitrate in a quantity of 1% of straw dry weight to a nutrient solution in one experiment while the nutrient solution had no nitrogen in a parallel experiment. Table I shows the calculated data. 

Dissolving is often an endothermic process. For example, when ammonium nitrate dissolves in water the temperature of the water falls, indicating that energy is being taken from the surroundings. Photosynthesis and thermal decomposition are other examples of endothermic processes. 

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