can anyone help me with this ideal gas law question?
Helium is collected over water at 19.0°C and 1.00 atm total pressure. What total volume of gas must be collected to obtain 0.601 g of helium? (At 25°C the vapor pressure of water is 23.8 torr.)
i wa thinking of just making a variation of the ideal gas law, specifically V= (nRT)/P. with the given units.
i find the number ofm oles through dimensional analysis of .601 grams and us 292 K, and 1 atm for the pressure. but if i do that then what with the stuff in the parenthases at the end?
please help, ive been stuck on this for a while
i dont think thats how you would have to do it, at least not with just one PV= nRT.
i think that you need to do partial pressures for both helium, and the water vapor that is collected and see how much of that gas is needed.
please i still need help
if it helps the answer should be around 4ish, not exactly 4, can be around there though
- Dr WLv 71 decade agoFavorite Answer
via daltons law of partial pressures...
Pt = PH2O + PHe
PHe = Pt - PH2O = 1.00 atm - 23.8 torr x (1 atm / 760 torr) = 0.969 atm
from ideal gas law..
PHeV = nRT
V = nRT/PHe
n = 0.601 g x (1 mole / 4.003g) = 0.150 moles
R = 0.08206 Latm/moleK
T = 19.0C = 292.2 K
PHe = 0.969 atm
V = (0.150 moles) x (0.08206 Latm/moleK) x (292.2 K) / (0.969 atm)
V = 3.72 L... 3 sig figs.
now that is the volume of the helium.. but since this is a gas, the helium and the water vapor are fully miscible and fill the same volume. ie the total volume collected needs to be 3.72 L
- michelettiLv 44 years ago
suitable gas regulation purely works on gasses with particular features. the suitable gas regulation assumes that the molecules do not attrat one yet another, have not have been given any mass, etc. this does not carry real below particular situations, like at temperatures close to to absolute 0 (the gas molecules show charm at such low temperatures)
- Anonymous1 decade ago
use pv = nrt
change your temperature into kelvin
change your pressure in to atmosphere
then it is easy and you should get it...