Chemistry... I freakin hate it

**Imisnew2** · 02-09-09, 04:00 AM

At 37 ºC (normal body temperature), the solubility of N2 in water is 0.0150 g/L when its pressure over the solution is 1.00 atm. Air is approximately 78.0 mol% N2. When a diver inhales air at a pressure of 1.00 atm, ___ mol of N2 are dissolved per liter of blood (essentially an aqueous solution) while at 100 feet below the surface, where the pressure is 4.00 atm, there are ___ mol of N2 dissolved per liter of blood.

I've been trying to get this right for HOURS.

N2 Molar mass = 14.0067

I'm so confused!!

**bust331** · 02-09-09, 07:20 AM

Trying to get your TTP buddies to answer it for you

Cheater

**Kraker Jak** · 02-09-09, 07:32 AM

7

**jmw_man** · 02-09-09, 07:45 AM

Chemistry is actually easy stuff...

when you understand it!

**sacredsarcasm** · 02-09-09, 07:50 AM

is this HS or college chem?

anyway, let me try it. its good practice.

PITFALL: you forgot the MM of nitrogen gas is the sum of both nitrogens. you got the formula right but you forgot the MM is also 2X the mass of N. MM N2=25.0134

get rid of extra info. and it seems the problem simplifies to:

N2 solubility = 0.0150 g/L @ 1.00 atm

so we are at 1.00 atm and N2 is 0.0150 g/L soluble at this pressure. my first thought was we need more info, like the volume of the divers lungs, but since we are doing this per L it again makes it a little easier.

if the blood(water) can absorb 0.0150g/L then using the MM it can absorb 5.35 x 10^-4 mol/L.

to get the solubility at 4 atm use the gas laws. P1/n1=P2/n2. in this case we just need to multiply the 1.00 atm answer by 4, to get the 4.00 atm answer. 2.14 x 10^-3 mol/L.

looks like your problem might have been the molar mass, but also dont forget to try and see through the extra info and get rid of it so you dont get confused.