CO2 questions

[Dave S]

I have a planted 400 litre tank; the spray bar is at the back covering
approx half the length of the aquarium, the intake is at the other end in
the corner - What is the best position to locate the Reactor, hidden in the
corner behind plants, or in the centre?

TIA - Dave

[Margolis]

"Dave S" > wrote in message
...
> I have a planted 400 litre tank; the spray bar is at the back covering
> approx half the length of the aquarium, the intake is at the other end in
> the corner - What is the best position to locate the Reactor, hidden in
the
> corner behind plants, or in the centre?
>
> TIA - Dave
>
>


wherever you want to put it, so long as it has good water flow around the
output

--

Margolis
http://web.archive.org/web/20030215212142/http://www.agqx.org/faqs/AGQ2FAQ.htm
http://www.unrealtower.org/faq

[Brian S.]

Hey Dave,

Being you have a large tank, I would attempt to put the CO2 output unit
(reactor, etc) near the middle of the tank possibly in the back wall. This
is where I have mine setup.

Reason for doing this is that the Co2 will travel in both directions in the
tank equally (or there abouts) so that you don't see one side of your tank
flourishing and the other side just slowly growing.

Just my two cents.

Brian S.
"Dave S" > wrote in message
...
> I have a planted 400 litre tank; the spray bar is at the back covering
> approx half the length of the aquarium, the intake is at the other end in
> the corner - What is the best position to locate the Reactor, hidden in
the
> corner behind plants, or in the centre?
>
> TIA - Dave
>
>

[Margolis]

"Dave S" > wrote in message
...
>
> Question: At what point is the CO2 dissapated into the water? (a) As the
> bubble rises to the surface, (b) When the bubble breaks the surface?
>
> TIA - Dave
> >
>
>


using a reactor, the co2 mixes with the water as the bubbles travel up thru
a column of water that is going down thru some bio balls to help mix it up.

If the bubbles make it out of the reactor and make it to the surface of the
tank then it is wasted, it just goes into the air instead of the water. co2
does not like to mix with water, that is why a good reactor is needed.


--

Margolis
http://web.archive.org/web/20030215212142/http://www.agqx.org/faqs/AGQ2FAQ.htm
http://www.unrealtower.org/faq

[Michi Henning]

"Margolis" > wrote in message
...
>
> If the bubbles make it out of the reactor and make it to the surface of the
> tank then it is wasted, it just goes into the air instead of the water. co2
> does not like to mix with water, that is why a good reactor is needed.

No, that's not correct. In fact, CO2 dissolves very readily in water. It's
just that you have to keep the gas in contact with the water for a little
while, to give it time to dissolve. That's why all CO2 reactors work
on the principle of slowing down the ascent of the bubbles, so the CO2
can dissolve into the water instead of having the bubbles rise to the surface
and burst, thereby releasing the CO2 into the air instead of the water.

Cheers,

Michi.

--
Michi Henning Ph: +61 4 1118-2700
ZeroC, Inc. http://www.zeroc.com

[Margolis]

"Michi Henning" > wrote in message
...
>
> No, that's not correct. In fact, CO2 dissolves very readily in water.


I have to disagree very strongly. co2 does not like to mix with water at
all. That is why it dissapates out of the water so rapidly if there is any
surface turbulance at all. That is also why co2 in water in nature is
around 2-3ppm while atmospheric concentrations are around 375ppm.

--

Margolis
http://web.archive.org/web/20030215212142/http://www.agqx.org/faqs/AGQ2FAQ.htm
http://www.unrealtower.org/faq

[Michi Henning]

Margolis" > wrote in message
...
> "Michi Henning" > wrote in message
> ...
> >
> > No, that's not correct. In fact, CO2 dissolves very readily in water.
>
>
> I have to disagree very strongly. co2 does not like to mix with water at
> all.

At room temperature, solubility of CO2 is 0.9 liters in one liter of water.
So, given equal volumes of CO2 and water, almost all of the CO2 can
be dissolved in the water. Venous blood contains 50-60% CO2 by
volume. Most people would call that "readily dissolves" (and many
text books use precisely that phrase).

> That is why it dissapates out of the water so rapidly if there is any
> surface turbulance at all.

No, that is not why this is happening. The reason is that there is an
equilibrium at which the amount of CO2 that dissolves into the water
equals the amount of gas that outgasses from the water. The
concentration at which that equilibrium occurs is determined by the
partial pressure of CO2 in air. If water contains more CO2 than
the natural equilibrium, CO2 will gradually escape from the water
until the natural balance is restored. If you agitate the water surface,
you increase the surface area at which water is in contact with air
and, as a result, CO2 escapes quicker than it would if the surface
were still. However, eventually, the exact same balance will be
established, whether you agitate the water or not.

> That is also why co2 in water in nature is
> around 2-3ppm while atmospheric concentrations are around 375ppm.

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.

One effect of the greenhouse gas problem is that, over the past decades,
we have increased the CO2 levels in the atmosphere, meaning that more
CO2 dissolves into the oceans than was previously the case. Because CO2
forms carbonic acid in water, that causes a pH drop. That drop in pH
(among other things) has been linked to the increasing death of corals
in the Great Barrier Reef. Basically, all the extra CO2 we are producing
readily dissolves into the oceans, raising the pH in the bargain.

In summary: CO2 *does* readily dissolve in water. If you want proof,
put a hose into your tank and pump CO2 into it. Watch each
CO2 bubble as it rises -- it gets visibly smaller as it makes its way
to the surface. The amount by which the bubble gets smaller is
the amount of CO2 that has dissolved into the water.

Cheers,

Michi.

--
Michi Henning Ph: +61 4 1118-2700
ZeroC, Inc. http://www.zeroc.com

[NetMax]

"Michi Henning" > wrote in message
...
> Margolis" > wrote in message
> ...
>> "Michi Henning" > wrote in message
>> ...
>> >
>> > No, that's not correct. In fact, CO2 dissolves very readily in
>> > water.
>>
>>
>> I have to disagree very strongly. co2 does not like to mix with water
>> at
>> all.
>
> At room temperature, solubility of CO2 is 0.9 liters in one liter of
> water.
> So, given equal volumes of CO2 and water, almost all of the CO2 can
> be dissolved in the water. Venous blood contains 50-60% CO2 by
> volume. Most people would call that "readily dissolves" (and many
> text books use precisely that phrase).
>
>> That is why it dissapates out of the water so rapidly if there is any
>> surface turbulance at all.
>
> No, that is not why this is happening. The reason is that there is an
> equilibrium at which the amount of CO2 that dissolves into the water
> equals the amount of gas that outgasses from the water. The
> concentration at which that equilibrium occurs is determined by the
> partial pressure of CO2 in air. If water contains more CO2 than
> the natural equilibrium, CO2 will gradually escape from the water
> until the natural balance is restored. If you agitate the water
> surface,
> you increase the surface area at which water is in contact with air
> and, as a result, CO2 escapes quicker than it would if the surface
> were still. However, eventually, the exact same balance will be
> established, whether you agitate the water or not.
>
>> That is also why co2 in water in nature is
>> around 2-3ppm while atmospheric concentrations are around 375ppm.
>
> 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.
>
> One effect of the greenhouse gas problem is that, over the past
> decades,
> we have increased the CO2 levels in the atmosphere, meaning that more
> CO2 dissolves into the oceans than was previously the case. Because CO2
> forms carbonic acid in water, that causes a pH drop. That drop in pH
> (among other things) has been linked to the increasing death of corals
> in the Great Barrier Reef. Basically, all the extra CO2 we are
> producing
> readily dissolves into the oceans, raising the pH in the bargain.
>
> In summary: CO2 *does* readily dissolve in water. If you want proof,
> put a hose into your tank and pump CO2 into it. Watch each
> CO2 bubble as it rises -- it gets visibly smaller as it makes its way
> to the surface. The amount by which the bubble gets smaller is
> the amount of CO2 that has dissolved into the water.
>
> Cheers,
>
> Michi.
>
> --
> Michi Henning Ph: +61 4 1118-2700
> ZeroC, Inc. http://www.zeroc.com

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.

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?

Using the same 0.9 litre example above, how would O2 figure?

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?

--
www.NetMax.tk who loves the science of keeping fish happy : )

[Michi Henning]

"NetMax" > wrote in message
...
> "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

For air at atmospheric pressure and water at 20 C, you end up
with 3 ppm of CO2 and 9.1 ppm of O2.

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.

> 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

[NetMax]

"Michi Henning" > wrote in message
...
> "NetMax" > wrote in message
> ...
>> "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

[Michi Henning]

"NetMax" > wrote in message
...
>
> > 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

Yes, right -- it would be different for oil ;-)

> 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?

Not quite. When I said O2 is 32 times more soluble, I was referring to
volume and the saturation point. When you say that O2 is 3 times higher
than CO2, you are using weight (because ppm is a measure of weight,
not volume), and you are referring to equilibrium, not saturation.

We can work out the ratio by volume:

Using partial pressure, air is 21% O2 and 0.033% CO2, so that's 636
times more O2 than CO2.

In water, at 20 C, we get 9.1 ppm O2 and 3 ppm CO2. As volume, that's around
6.4 ml O2/l and 1.5 ml CO2/l. That's a ratio of 4.26.


So, you could say that O2 concentration in air is 636 times that of CO2,
and that it is 4.26 times that of CO2 in water. (Incidentally, the difference
in O2 concentration between water and air is one reason why animals
originally evolved to leave the water and move onto land. In air, you
have 21% oxygen whereas, in water, you only have around 0.6%
oxygen. The maneuverability of an animal is largely determined by
how much energy it can produce, and energy is produced by
oxidizing nutrients. Ergo, an animal on land can move faster and
for longer periods than an animal in water because air contains so
much more oxygen so a land animal can produce more energy
per unit of time.)

So why is the actual CO2/O2 equlibrium ratio in water not 636?
The ratio at equilibrium isn't just determined by partial pressure, but
also by a constant that is different for each gas and essentially captures
solubility at a given temperature. For CO2, the constant is
3.38 x 10^-2, for O2, it is 1.28 x 10^-3 at a temperature of
25C. Solubility is determined by multiplying
the constant by the partial pressure of the gas.
To work the problem for different temperatures, you also need to
know the enthalpy of dissolution for each gas. For CO2, that's
-5.242 kcal, and for O2, that's -3.5 kcal. An equation called
the Clausius-Clapeyron equation then allows you adjust
solubility for different temperatures. If you do all that, you
end up with the 32 times at 20 C that I mentioned earlier.

> 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.

Right. There is limit to how much O2 can dissolve into the water. However,
as the fish breath and use up O2, more O2 from the gas volume above
the water can dissolve into the water, repleneshing what's used up by the
fish.

Cheers,

Michi.

--
Michi Henning Ph: +61 4 1118-2700
ZeroC, Inc. http://www.zeroc.com

[NetMax]

"Michi Henning" > wrote in message
...
> "NetMax" > wrote in message
> ...
>>
>> > 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
>
> Yes, right -- it would be different for oil ;-)
>
>> 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?
>
> Not quite. When I said O2 is 32 times more soluble, I was referring to
> volume and the saturation point. When you say that O2 is 3 times higher
> than CO2, you are using weight (because ppm is a measure of weight,
> not volume), and you are referring to equilibrium, not saturation.
>
> We can work out the ratio by volume:
>
> Using partial pressure, air is 21% O2 and 0.033% CO2, so that's 636
> times more O2 than CO2.
>
> In water, at 20 C, we get 9.1 ppm O2 and 3 ppm CO2. As volume, that's
> around
> 6.4 ml O2/l and 1.5 ml CO2/l. That's a ratio of 4.26.
>
>
> So, you could say that O2 concentration in air is 636 times that of
> CO2,
> and that it is 4.26 times that of CO2 in water. (Incidentally, the
> difference
> in O2 concentration between water and air is one reason why animals
> originally evolved to leave the water and move onto land. In air, you
> have 21% oxygen whereas, in water, you only have around 0.6%
> oxygen. The maneuverability of an animal is largely determined by
> how much energy it can produce, and energy is produced by
> oxidizing nutrients. Ergo, an animal on land can move faster and
> for longer periods than an animal in water because air contains so
> much more oxygen so a land animal can produce more energy
> per unit of time.)
>
> So why is the actual CO2/O2 equlibrium ratio in water not 636?
> The ratio at equilibrium isn't just determined by partial pressure, but
> also by a constant that is different for each gas and essentially
> captures
> solubility at a given temperature. For CO2, the constant is
> 3.38 x 10^-2, for O2, it is 1.28 x 10^-3 at a temperature of
> 25C. Solubility is determined by multiplying
> the constant by the partial pressure of the gas.
> To work the problem for different temperatures, you also need to
> know the enthalpy of dissolution for each gas. For CO2, that's
> -5.242 kcal, and for O2, that's -3.5 kcal. An equation called
> the Clausius-Clapeyron equation then allows you adjust
> solubility for different temperatures. If you do all that, you
> end up with the 32 times at 20 C that I mentioned earlier.
>
>> 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.
>
> Right. There is limit to how much O2 can dissolve into the water.
> However,
> as the fish breath and use up O2, more O2 from the gas volume above
> the water can dissolve into the water, repleneshing what's used up by
> the
> fish.
>
> Cheers,
>
> Michi.
>
> --
> Michi Henning Ph: +61 4 1118-2700
> ZeroC, Inc. http://www.zeroc.com
>


Whew, you just know that I had to read that more than once for it to sink
in, thanks for the lessons Michi!!

One last question (well maybe), since CO2 and O2 have a different
enthalpy of dissolution as a function of temperature, can I take that to
mean that as water gets warmer, then the point of equilibrium for CO2
will drop more quickly than O2? (making it more difficult to keep
elevated CO2 levels in warmer water tanks (80F+)).
--
www.NetMax.tk

[Michi Henning]

"NetMax" > wrote in message
.. .

> Whew, you just know that I had to read that more than once for it to sink
> in, thanks for the lessons Michi!!
>
> One last question (well maybe), since CO2 and O2 have a different
> enthalpy of dissolution as a function of temperature, can I take that to
> mean that as water gets warmer, then the point of equilibrium for CO2
> will drop more quickly than O2? (making it more difficult to keep
> elevated CO2 levels in warmer water tanks (80F+)).

In general, as temperature goes up, less of a gas can dissolve into
a given volume of water. That's true for all gases. Because CO2
has a larger enthalpy of dissolution than O2, the solubility of CO2
changes more quickly as a function of temperature than the solubility
of O2. So, yes, as temperature goes up, you'll be able to put relatively
less CO2 into the water than at lower temperature. But I think the overall
effect of the higher temperature is more significant than this shift in the
balance between O2 and CO2. Getting CO2 into an 80F+ tank is harder
mainly because the warmer water dissolves less of any gas.

Cheers,

Michi.

--
Michi Henning Ph: +61 4 1118-2700
ZeroC, Inc. http://www.zeroc.com