Question

Ammonium nitrate has been used as a high explosive because it is unstable and decomposes into several gaseous substances. The rapid expansion of the gaseous substances produces the explosive force. NH4NO3(s) + N2(g) + O2(g) + H2O(9) Calculate the mass of each product gas if 84.6 g of ammonium nitrate reacts. g N₂ 802 gH,0

Answer 1

The balanced decomposition reaction of ammonium nitrate is as follows:

2NH₄NO₃(s) $\rightarrow$ 2N₂(g) + O₂(g) + 4H₂O(g)

Mass of NH₄NO₃ = 84.6 g

Determine the number of moles of NH₄NO₃ from the given mass and molar mass as follows:

The molar mass of NH₄NO₃ = 80.043 g/mol

= 84.6 g NH₄NO₃ x ( 1 mol NH₄NO₃ / 80.043 g NH₄NO₃)

= 1.0569 mol NH₄NO₃

Use the moles of NH₄NO₃ and the mole ratio from the balanced chemical reaction and determine the number of moles of N₂, O₂, and H₂O as follows:

= 1.0569 mol NH₄NO₃ x ( 2 mol N₂ / 2 mol NH₄NO₃)

= 1.0569 mol N₂

Similarly,

= 1.0569 mol NH₄NO₃ x ( 1 mol O₂ / 2 mol NH₄NO₃)

= 0.5285 mol O₂

Similarly,

= 1.0569 mol NH₄NO₃ x ( 4 mol H₂O / 2 mol NH₄NO₃)

= 2.114 mol H₂O

Convert the moles of each product to grams by using the molar masses as follows:

The molar mass of N₂ = 28.0134 g/mol

The molar mass of O₂ = 31.9988 g/mol

The molar mass of H₂O = 18.02 g/mol

Now,

= 1.0569 mol N₂ x (28.0134 g N₂ /1 mol N₂)

= 29.6 g N₂

Similarly,

= 0.5285 mol O₂ x (31.9988 g O₂ /1 mol O₂)

= 16.9 g O₂

Similarly,

= 2.114 mol H₂O x (18.02 g H₂O /1 mol H₂O)

= 38.1 g H₂O

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