Open Question
Nitrogen dioxide dimerizes to give dinitrogen tetroxide: 2 NO2(g) ⇌ N2O4(g). At 298 K, 9.66 g of an NO2 and N2O4 mixture exerts a pressure of 0.487 atm in a volume of 6.51 L. What are the mole fractions of the two gases in the mixture?
Ch.10 - Gases: Their Properties & Behavior
Chapter 10, Problem 145c
An empty 4.00-L steel vessel is filled with 1.00 atm of CH4(g) and 4.00 atm of O2(g) at 300 °C. A spark causes the CH4 to burn completely, according to the equation
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) ΔH° = -802 kJ
(c) What is the partial pressure of CO2(g) in the vessel after combustion?
Verified step by step guidance1
Identify the balanced chemical equation for the combustion of methane: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g). This equation shows that one mole of CH4 reacts with two moles of O2 to produce one mole of CO2 and two moles of H2O.
Determine the initial moles of CH4 and O2 using the ideal gas law, PV = nRT. Given that the pressure (P) of CH4 is 1.00 atm and the volume (V) is 4.00 L, calculate the moles of CH4. Similarly, calculate the moles of O2 using its initial pressure of 4.00 atm.
Use stoichiometry to find the limiting reactant. Compare the mole ratio of CH4 to O2 from the balanced equation with the initial moles calculated. The reactant that produces the least amount of product is the limiting reactant.
Calculate the moles of CO2 produced using the stoichiometry of the reaction. Since the limiting reactant determines the amount of product formed, use the mole ratio from the balanced equation to find the moles of CO2.
Determine the partial pressure of CO2 using the ideal gas law. With the moles of CO2 calculated and the volume and temperature of the vessel known, apply the ideal gas law to find the partial pressure of CO2 in the vessel after combustion.
Verified Solution
Video duration:
6m
This video solution was recommended by our tutors as helpful for the problem above.Was this helpful?
Related Practice
Textbook Question
An empty 4.00-L steel vessel is filled with 1.00 atm of CH4(g) and 4.00 atm of O2(g) at 300 °C. A spark causes the CH4 to burn completely, according to the equation
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) ΔH° = -802 kJ
(a) What mass of CO2(g) is produced in the reaction?
575
views
Textbook Question
An empty 4.00-L steel vessel is filled with 1.00 atm of CH4(g) and 4.00 atm of O2(g) at 300 °C. A spark causes the CH4 to burn completely, according to the equation
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) ΔH° = -802 kJ
(b) What is the final temperature inside the vessel after combustion, assuming that the steel vessel has a mass of 14.500 kg, the mixture of gases has an average molar heat capacity of 21 J/(mol·°C), and the heat capacity of steel is 0.449 J/(g·°C)?
685
views
Textbook Question
When a gaseous compound X containing only C, H, and O is burned in O2, 1 volume of the unknown gas reacts with 3 volumes of O2 to give 2 volumes of CO2 and 3 volumes of gaseous H2O. Assume all volumes are measured at the same temperature and pressure. (a) Calculate a formula for the unknown gas, and write a balanced equation for the combustion reaction.
793
views
Textbook Question
Isooctane, C8H18, is the component of gasoline from which the term octane rating derives. (a) Write a balanced equation for the combustion of isooctane to yield CO2 and H2O.
518
views
Textbook Question
Isooctane, C8H18, is the component of gasoline from which the term octane rating derives. (b) Assuming that gasoline is 100% isooctane, that isooctane burns to produce only CO2 and H2O, and that the density of isooctane is 0.792 g/mL, what mass of CO2 in kilograms is produced each year by the annual U.S. gasoline consumption of 4.6⨉1010 L?
1135
views