exchemist
Valued Senior Member
Seems topical to post the two competing decomposition reactions that can lead to explosion:
1) NH4NO3 -> N2O + 2H2O , predominant at lower temperatures
2) 2NH4NO3 -> 2N2 + 4H2O + O2 , at higher temperatures
Both are exothermic.
In the second mode, it generates free oxygen, capable or reacting with another oxidisable substance, as well. Hence its use as an ingredient of explosives when mixed with other things, most notoriously, in the case of the IRA, sugar.
The Beirut explosion was notable for its orange cloud, which is due to NO2, a secondary reaction product which absorbs mainly in the blue/green region of the spectrum, leading to a red/orange colour.
Unrelated to its explosive potential, ammonium nitrate is also interesting because it dissolves in water very endothermically. You stir it in and the water becomes instantly cold! It is a nice demonstration of how a reaction can occur spontaneously, even though it requires significant energy input, due to the increase in entropy associated with dissolution. Reactions take place when there is a reduction in the free energy, G, i.e. a negative value for ΔG = ΔH - TΔS. In this case even though the enthalpy, H change is +ve (i.e. you have to put energy in), ΔS is sufficiently +ve that the -TΔS term wins.
As to why the entropy change is so +ve in the case of ammonium nitrate, I am don't have the answer. It may be that less of a solvent cage is formed when these ions dissolve than is the case with most ionic compounds, so there is not much of an ordering process among the solute molecules to offset again the disordering due to the break up of the crystal structure. But I have never seen this explained so I am speculating.
1) NH4NO3 -> N2O + 2H2O , predominant at lower temperatures
2) 2NH4NO3 -> 2N2 + 4H2O + O2 , at higher temperatures
Both are exothermic.
In the second mode, it generates free oxygen, capable or reacting with another oxidisable substance, as well. Hence its use as an ingredient of explosives when mixed with other things, most notoriously, in the case of the IRA, sugar.
The Beirut explosion was notable for its orange cloud, which is due to NO2, a secondary reaction product which absorbs mainly in the blue/green region of the spectrum, leading to a red/orange colour.
Unrelated to its explosive potential, ammonium nitrate is also interesting because it dissolves in water very endothermically. You stir it in and the water becomes instantly cold! It is a nice demonstration of how a reaction can occur spontaneously, even though it requires significant energy input, due to the increase in entropy associated with dissolution. Reactions take place when there is a reduction in the free energy, G, i.e. a negative value for ΔG = ΔH - TΔS. In this case even though the enthalpy, H change is +ve (i.e. you have to put energy in), ΔS is sufficiently +ve that the -TΔS term wins.
As to why the entropy change is so +ve in the case of ammonium nitrate, I am don't have the answer. It may be that less of a solvent cage is formed when these ions dissolve than is the case with most ionic compounds, so there is not much of an ordering process among the solute molecules to offset again the disordering due to the break up of the crystal structure. But I have never seen this explained so I am speculating.
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