"Michi Henning" wrote in message
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"NetMax" wrote in message
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"Michi Henning" wrote in message
No, that's not the reason. The reason is that air only contains
around 0.0314% CO2. The amount of CO2 that dissolves in water
is limited by the partial pressure of CO2 in the air, not by the
solubility
of CO2 in water. If you put a 100% CO2 atmosphere above water,
you end up with the above mentioned 0.9 liters of CO2 per liter
of water.
Another way to look at this is, the rate at which a gas leaves (CO2
quickly outgassing) is indicative of how easily or slowly it enters.
I don't think that's correct. Basically, if you put a gas mixture (such
as
air) over water, the equilibrium that establishes itself depends on a
number of factors:
- temperature
- the partial pressure of each gas in the mixture
- the solubility in water of each gas in the mixture
heh heh, I guess I oversimplified it too much. I just like to put things
in terms that even I understand.
For air at atmospheric pressure and water at 20 C, you end up
with 3 ppm of CO2 and 9.1 ppm of O2.
in water
If we disturb the equilibrium, so the water contains a gas in excess,
the rate at which the gas leaves the water depends on how much
exceess of the gas you have. For example, if you saturate a liter of
water with 0.9 liters of CO2 and then leave the water standing in the
open air, the rate of outgassing of CO2 would initially be high and
drop as
the liquid gets closer to equilibrium. If you were to graph the CO2
concentration of the liquid against time, you'd see a curve that is
initially steep and then gradually flattens toward the equilibrium
line.
Non-linear curve makes sense.
Michi, can the same thing be said about O2, or are there additional
variables at work?
What is the equilibrium %s of O2 in water vs air?
At 20 C, equilibrium of O2 in freshwater is 9.1 ppm, whereas,
for CO2, it's 3 ppm. But that doesn't indicate solubility. Instead
it simply indicates the equilibrium values. To get solubility, you
have to put water into a pure CO2 or O2 atmosphere and see
how much of the gas you end up with in the water.
If you do this with O2, you get 40 ppm, that is, you can dissolve
around 0.04 g of O2 in a liter of water. I you do this with CO2,
you get 1767 ppm, that is, you can dissolve around 1.76 g
of CO2 in a liter of water.
It's probably more meaningful to talk about volume instead
of weight, because that way, we don't distort the figures
by the different molecular weights of O2 and CO2. Expressed
as volume, this means that you can dissolve 900ml of CO2
in a liter of water, but only 28ml of O2. In other words, solubility
of CO2 in water is about 32 times better than that of O2 in water.
Using the same 0.9 litre example above, how would O2 figure?
See above.
My understanding is that under normal circumstances, CO2 (injection)
will
not crowd out O2. Using the litre saturation points, where would CO2
and
O2 level off if applied simultaneously?
At the 20-30 ppm of CO2 that we aim for in a planted tank, we are
nowhere near the saturation value of CO2, so there is indeed no issue
with "crowding out" oxygen. If you expose water to a mixture of CO2
and O2, then the gases dissolve in the water in proportion to their
partial pressures (Henry's law).
Cheers,
Michi.
--
Michi Henning Ph: +61 4 1118-2700
ZeroC, Inc. http://www.zeroc.com
Thanks, I think I have it now. O2 is 32 times less soluble in water than
CO2, but since the atmospheric O2 concentrations are higher, the
resultant equilibrium has O2 at 3 times higher than CO2. Can I
extrapolate that atmospheric O2 is in concentrations of 96 times higher
than CO2?
Your description explains how they can put a large amount of pure O2 into
fishbags for trans-shipping. The water simply will not absorb more than
about 4 times what it already has. I imagine fishbags don't have the
best O2 barrier anyways.
--
www.NetMax.tk