Question

Please help, will fan and medal.

On a spring day, you take a road trip through Yakima, WA, in a sports car. You start out at a temperature of 21°C, but the temperature in Yakima will reach a peak of 51°C. Each tire on your car hold 15.6 L of nitrogen gas at a starting pressure of 249 kPa. The tires will burst when the internal pressure (Pb) exceeds 269 kPa for 10 minutes. Answer the following questions and show your work.

Answer 1

1. How many moles of nitrogen gas are in each tire?

2. What would the tire pressure be at peak temperature in Yakima?

3. Will the tires burst in Yakima? Explain.

4. If you must let nitrogen gas out of the tire before you go, to what pressure must you reduce the tires before you start your trip? (Assume no significant change in tire volume.)

5. Create a drawing of the tire and show a molecular view of the air molecules in the tire at 249 kpa vs the molecular view of the air molecules after the tires have been heated. Use a caption to describe the average kinetic energy of the molecules in both scenarios.

Answer 2

a) use the ideal gas law, PV=nRT

Answer 3

or rather, 1) not a)

Answer 4

I only need 5. now.

Answer 5

would the same type of molecules move faster or slower at higher temperatures?

If they move faster, would they not have more collisions with one another, keeping a greater average distance between one another?