In the Haber reaction, patented by German chemist Fritz Haber in 1908, dinitrogen gas combines with dihydrogen gas to produce gaseous ammonia. Thisreaction is now the first step taken to make most of the world's fertilizer.Suppose a chemical engineer studying a new catalyst for the Haber reaction finds that 101. liters per second of dinitrogen are consumed when the reaction isrun at 262. °C and the dinitrogen is supplied at 0.16 atm. Calculate the rate at which ammonia is being produced. Give your answer in kilograms per second.Round your answer to 2 significant digits.kg?

Given: The volume (V) of the dinitrogen gas (N2) is 101 L The pressure (P) of the dinitrogen gas…

Answered: In the Haber reaction, patented by…

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

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An unknown mass of aluminum is added to HCl solution. The aluminum metal is completely dissolved in the solution and generates hydrogen gas. See the balanced equation for the gas evolution reaction: 2 Al(s) + 6 HCl(aq) ---> 2 AlCl3(aq) + 3 H2(g) The following data is collected: Mass of Al sample: unknown g Volume of gas collected 45.38 mL Barometric pressure 754.8 mmHg Temperature 23.5OC Vapor pressure of H2O at 23.5OC 21.65 mmHg Hydrogen gas is collected over water. The reaction proceeds until all the amount of aluminum is consumed in the gas-evolution reaction (aluminum is the limiting reactant). (a) Calculate the mole number of the hydrogen gas generated in the reaction: Convert the volume of the collected gas to L: 45.38 mL = ___________ L Calculate the pressure exerted by H2 gas (= dry hydrogen gas): Pressure of the dry hydrogen gas =…
For the following reaction: C3H8(g) + 3CO₂(g) + 4H₂O(1) 502(g) The change in gaseous moles, An = To maximize the product yield, the volume of the container should ◆ , so that the pressure will be (> →
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 592. liters per second of dioxygen are consumed when the reaction is run at 195. °℃ and the dioxygen is supplied at 0.48 atm. Calculate the rate at which sulfur dioxide is being produced. Give your answer in kilograms per second. Be sure your answer has the correct number of significant digits.
Automobile air bags employ a chemical reaction involving the rapid decomposition of sodium azide (NaN3) into sodium metal and nitrogen gas to inflate in a collision. How much sodium azide is required to decompose to inflate a 20.0 L air bag? Assume that the gas is at standard temperature and pressure, and thus occupies 22.4 L per mole. (AW: Na=23; N=14.007)
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 978. liters per second of dioxygen are consumed when the reaction is run at 272. °℃ and the dioxygen is supplied at 0.91 atm. Calculate the rate at which nitrogen monoxide is being produced. Give your answer in kilograms per second. Be sure your answer has the correct number of significant digits. ៣ kg S x10 X
Ammonia is among the top ten synthesized compounds. Its many uses include the manufacture of ammonium nitrate and other fertilizers. Ammonia decomposes at high temperatures. In an experiment to explore this behavior, 2.00 moles of gaseous NH3 are sealed in a rigid 1-liter vessel. The vessel is heated at 800 K and some of the NH3 decomposes in the following reaction: 2NH3(g) <--> N2(g)+3H2(g) The system eventually reaches equilibrium and is found to contain 1.74 mole of NH3. What is the value of KC for the decomposition reaction at 800K? What is the value of KP for the decomposition reaction at 800 K?
A chemical engineer must calculate the maximum safe operating temperature of a high-pressure gas reaction vessel. The vessel is a stainless-steel cylinder that measures 52.0 cm wide and 62.4 cm high. The maximum safe pressure inside the vessel has been measured to be 7.20 MPa. For a certain reaction the vessel may contain up to 6.16 kg of dinitrogen monoxide gas. Calculate the maximum safe operating temperature the engineer should recommend for this reaction. Write your answer in degrees Celsius. Be sure your answer has the correct number of significant digits. temperature: °C
Ammonia is among the top ten synthesized compounds. Its many uses include the manufacture of ammonium nitrate and other fertilizers. Ammonia decomposes at high temperatures. In an experiment to explore this behavior, 2.00 moles of gaseous NH3 are sealed in a rigid 1-liter vessel. The vessel is heated at 800 K and some of the NH3 decomposes in the following reaction:2NH3(g)<-->N2(g)+3H2(g) The system eventually reaches equilibrium and is found to contain 1.74 mole of NH3. What is the value of KC for the decomposition reaction at 800 K? What is the value of KP for the decomposition reaction at 800 K?
Ff.119.
The reaction of 50 mL of N2 gas with 150 mL of H2 gas to form ammonia via the equation: N2 (g) + 3H2 (g) → 2NH3 (g) will produce __________ mL of ammonia if pressure and temperature are kept constant.
Q7. In our campus, there are approximately 30.000 students use 2.5 x 10° kJ of energy per day. Suppose all of that energy comes from a combustion of a carbon-hydrogen-oxygen compound in the presence of excess O2(g). 2.647 g sample of this gaseous carbon-hydrogen-oxygen compound that occupies a volume of 580 mL at 918,6 Torr and 24.00 °C. The products of the combustion of given amount are 5.059 g CO2(g), 3.106 g H2O(1) and enough heat to raise the temperature of the calorimeter assembly from 24.00 to 38.33 °C. What is the molecular formula of this unknown compound and how many kilograms of this compound are needed to provide enough daily energy to all students in campus? (The heat capacity of the calorimeter is 4.915 kJ/°C.)
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