Question

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 670. liters per second of dioxygen are consumed when the reaction is run at 203.°C and the dioxygen is supplied at 0.23atm. Calculate the rate at which nitrogen monoxide is being produced. Give your answer in kilograms per second. Round your answer to 2 significant digits.

Answer 1

To determine the rate at which nitrogen monoxide (NO) is being produced in the reaction of ammonia (NH3) with dioxygen (O2), we need to use the stoichiometry of the reaction and the given information.

Step-by-step explanation:

To determine the rate at which nitrogen monoxide (NO) is being produced in the reaction of ammonia (NH3) with dioxygen (O2), we need to use the stoichiometry of the reaction and the given information. The balanced equation for the reaction is 4NH3(g) + 5O2(g) →4NO(g) + 6H2O(g).

From the balanced equation, we can see that for every 4 moles of ammonia consumed, 4 moles of nitrogen monoxide are produced. Therefore, the rate of formation of NO is equal to the rate of consumption of ammonia.

Given that 670 liters per second of O2 are consumed and considering the stoichiometry of the reaction, we can calculate the rate of formation of NO as follows:

670 L/s of O2 * (4 mol NO / 5 mol O2) * (30.02 g / mol) * (1 kg / 1000 g) = 16.8112 kg/s