A scuba diver 60 ft below the ocean surface inhales 95.0 mL of compressed air from a scuba tank at an initial pressure of 2.70 atm and temperature of degrees Celsius. What is the final pressure of air, in the atmosphere in the lungs when the gas expands to 195.0 mL at a body temperature of 37 degrees Celsius and the amount of gas constant?

Question

A scuba diver 60 ft below the ocean surface inhales 95.0 mL of compressed air from a scuba tank at an initial pressure of 2.70 atm and temperature of  degrees Celsius. What is the final pressure of air, in the atmosphere in the lungs when the gas expands to 195.0 mL at a body temperature of 37 degrees Celsius and the amount of gas constant?

aaronq · Answer

you use a variation of the ideal gas law, PV=nRT

Set it equal to zero, $\dfrac{PV}{nRT}=0$

Were gonna compare 2 systems (underwater and above water), so:

                 $\underbrace{\dfrac{P_1V_1}{n_1RT_1}}_{system~1}=\underbrace{\dfrac{P_2V_2}{n_2RT_2}}_{system~2}$

since n is staying constant (one breath) and R is itself a constant, we can ignore them.

                   $\dfrac{P_1V_1}{T_1}=\dfrac{P_2V_2}{T_2}$

Plug your values in and solve for the new pressure. 

(by the way, the initial temperature of the gas is missing from the question, you need that).