Nitric acid is a key industrial chemical, largely used to make fertilizers and explosives. the first step in its synthesis is the oxidation 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 reaction finds that liters per second of dioxygen are consumed when the reaction is run at and . calculate the rate at which nitrogen monoxide is being produced. give your answer in kilograms per second. round your answer to significant digits.

When using 645 L /s  of O2 in a temperature and pressure of  195°C,  0.88 atm  respectively, we will get 0.355Kg /s NO

Explanation:

4NH3(g) + 5O2(g) ⟶4 NO(g) +6 H2O(l)  

O2 Volume Rate = 645 L /s

Pressure = 0.88 atm

Temperature = 195°C + 273 = 468K

NO molecular weight = 30.01 g/mol  

Let´s keep in mind that using this equation our constant R is 0.08205Lxatm/Kxmol

PV =n RT

n= PV / RT

n= [ 0.88atm x 645L/s] / [ (0.08205 Lxatm/Kxmol) x 468K]

n= 14.781 moles /s of O2

14.781 moles /s of O2 x 4moles NO / 5 moles O2 x  30.01g NO / 1 mol NO x 1Kg NO /1000g NO = 0.355Kg /s NO

the answer is c i took the test