Final answer:
The FALSE statement among the given options is (B), which claims that each container has the same number of molecules. Due to different molar masses, 16 g CH4, 44.0 g CO2, and 146 g SF6 do not equate to the same number of molecules.
Step-by-step explanation:
You asked which statement is FALSE about three identical steel containers at the same temperature filled with different gases: 16 g of methane (CH4), 44.0 g of carbon dioxide (CO2), and 146 g of sulfur hexafluoride (SF6). First, let's establish some facts based on the ideal gas law and principles from the kinetic molecular theory.
The masses provided correspond to different quantities of moles for each gas since they have different molar masses. However, if these containers are at the same temperature and volume, then according to Avogadro's law, each container would have the same number of moles of gas.
Statement (A) is TRUE: Since each container has the same number of moles, and SF6 has a higher molar mass than CO2 which is higher than CH4, the densities would decrease in the order SF6 > CO2 > CH4. Statement (B) is FALSE. Since the molar masses of the gases are different, 16 g of CH4, 44.0 g of CO2, and 146 g of SF6 do not represent an equal number of molecules.
Statement (C) assumes ideal behavior, so if the volume is held constant, and each container has the same number of moles at the same temperature, they would exert the same pressure.
Lastly, Statement (D) is TRUE because gases at the same temperature have the same average kinetic energy, but they do not have the same average speed due to different molecular masses. Since the speed depends on mass, lighter molecules like CH4 will move faster on average compared to heavier molecules like SF6 under the same conditions of temperature.