Question

Nitric acid is a key industrial chemical, largely used to make fertilizers and explosives, the first step in its synthesis is the oxidiation of ammonia. In this reaction, gaseous ammonia reacts with dioxygen gas to produce nitrogen monoxide gas and water.

Suppose a chemical engineer studying a new catalyst for the oxidation of ammonia reactions finds that 613. liters per second of dioxygen are consumed when the reaction is run at 175. C and 0.16 atn=m. Calculate the rate at which nitrogen monoxide is being produced. Give your answer in Kg/second. Round your answer to 2 significant digits.

Answer 1

Sol . As Reaction :

4NH3(g) + 5O2(g) ---> 4NO(g) + 6H2O(g)

Now , for O2 gas :

Pressure = P = 0.16 atm

Volume = V = 613 L

Temperature = T = 175°C = 448 K

R = Gas constant = 0.0821 L atm / K mol

So , moles of O2 = n = PV /RT

= (0.16×613) / (0.0821 × 448)

= 2.666 mol

As from reaction , 5 moles of O2 gives 4 moles of NO.

So, 2.666 moles of O2 gives = (4×2.666) / 5 = 2.1328 moles of NO.

As molar mass of NO = 30 g/mol

So, mass of NO produced = 2.1328 × 30 = 63.984 g

= 63.984 /1000 Kg = 0.064 Kg

Therefore , Rate at which NO produced is  0.064 Kg/s