On a summer day, you take a road trip through Death Valley, California, in an antique car. You start out at a temperature of 21掳C, but the temperature in Death Valley will reach a peak of 51掳C. The tires 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) reaches 269 kPa. Answer the following questions and show your work.
?How many moles of nitrogen gas are in each tire?
?What will the tire pressure be at peak temperature in Death Valley?
?Will the tires burst in Death Valley? Explain.
?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.)Pressure, Temperature, Chemistry?I have no calculator on me. But you can plug in the values.
PV = nRT
P = pressure = 249kPa --%26gt; convert to atm --%26gt; 2.457438934atm
V = 15.6L
n = ?
R = 0.0821L*atm*mol^-1*K^-1
T = 21掳C --%26gt; convert to K --%26gt; 294K
n = PV / RT
P1V1 / T1 = P2V2 / T2
249kPa*15.6L / 294K = ?kPa*15.6L / (273 + 51)K
Isolate for P2 and you'll get your answer. If it is equal or greater than 269kPa, then the tires will burst.
For the last one, I believe we can assume that the maximum pressure the tires can be at peak temperature is 268kPa.
P1V1 / T1 = P2V2 / T2
?*15.6L / 294K = 268kPa*15.6L / (273 + 51)K. The answer will give you the maximum initial tire pressure.
I hope this helps.
