Question

Most of the sulfur used in the United States is chemically synthesized from hydrogen sulfide gas recovered from natural gas wells. In the first step of this synthesis, called the Claus process, hydrogen sulfide gas is reacted with dioxygen gas to produce gaseous sulfur dioxide and water.
Suppose a chemical engineer studying a new catalyst for the Claus reaction finds that 994 liters per second of dioxygen are consumed when the reaction is run at 170°C and 0.77 atm. Calculate the rate at which sulfur dioxide is being produced. Give your answer in kilograms per second. Round your answer to 2 significant digits.

Answer 1

Answer:

The rate at which sulfur dioxide is being produced is 0.90 kg/s.

Explanation:

Volume of oxygen gas consumed in second ,V= 994 L

Pressure of the gas = p

Temperature of the gas = T = 170°C= 170 + 273 K=443 K

Moles of oxygen gas consumed in a second = n

$PV=nRT$ ( ideapl gas equation)

$n=\frac{PV}{RT}=\frac{0.77 atm\times 994 L}{0.0821 atm L/mol K\times 443 K}$

n = 21.044 mole

Moles of dioxygen gas consumed per second = 21.044 mol

$2H_2S(g)+3O_2(g)\rightarrow 2SO_2(g)+2H_2O(g)$   (Claus process)

According to reaction, 3 moles of dioxygen gives 2 moles of sulfur dioxide gas.Then 21.044 moles of dioxygen will give;

$\frac{2}{3}\times 21.044 ol=14.029 mol$ of sulfur dioxide

Mass of 14.029 moles of sulfur dioxide gas;

14.029 mol × 64 g/mol = 897.86 g

897.86 g = 0.89786 kg ≈ 0.90 kg

Mass of sulfur dioxide produced per second = 0.90 kg

The rate at which sulfur dioxide is being produced is 0.90 kg/s.