# Aqueous Carbon Chemistry: or Organic Chemistry Finally Paid Off!



## Phil Edwards (Jan 22, 2004)

I read an interesting little bit of information the other day while working in the lab. According to William Schlesinger, one of the best biogeochemists in the world (ranking up there with Tom's old prof, K.R. Reddy) at the pH levels we commonly keep in our tanks Carbon would be available as Carbonate in nature. This got me thinking on our CO2 supplimentation habits and the chemistry going on inside the glass boxes we keep.

After scratching down some chemical reactions I came to the conclusion that Dr. Schlesinger is right. Since chemistry is thirsty work I opened a bottle of soda and the evidence sprayed me in the face. Unless kept under artifically high pressures CO2 gas is not present in aqueous solution. Instead, it is absorbed and converted into Carbonic Acid, which then undergoes some pretty spiffy chemistry.

In our tanks with CO2 gas supplimentation the CO2 is quickly converted to Carbonic Acid by interaction with water which then reacts with other free water molecules to form Bicarbonate and Hydronium ions. These steps are what happens in our reaction chambers and is how CO2 is absorbed and held in aqueous solution.

1. CO2 + H20 --> H2CO3 (Carbonic Acid)
2. H2CO3 + H20 --> H3O+ and HCO3- (Hydronium and Bicarbonate)

3. HCO3- + OH- --> CO3(2-) (Carbonate)

Due to water's property of auto-ionization [where water molecules interact to form OH- (Hydroxide) and H+ (Hydride), OH- being basic and H+ being acidic] the Bicarbonate can react with the strongly basic Hydroxide to form Carbonate and H20. This is pretty unstable stuff, but is more easily accessible by plants than Bicarbonate since it's chemically similar to CO2 and doesn't have a lot of side groups to remove in order to be useful in photosynthesis.

Further evidence is the mechanism we use to test for CO2 levels in water, pH. The definition of an acid is a molecule that has at least one Proton (Hydrogen atom without the electron) available to be donated to or removed by another molecule/atom/ion. Since CO2 has no Hydrogens to provide the protons it can't be an acid. Therefore, CO2 is not available in our aquariums.

If there were nothing adding or removing one or more of the elements/atoms/ions involved in these steps the reactions would go on and on in equilibrium forever. In the cases of our tanks though, we have plants which do a smashing job of removing what carbon they can from the water, throwing a wrench into the lovely equilibrium we've established. That's why we add CO2 or another source of carbon to our tanks.

If you're going by Diana Walstad's tests then it stands to reason that the plants would do better in a tank with harder water because the elements are there to provide/produce higher amounts of carbonate/bicarbonate in the absence of CO2 supplimentation. Decomposition of organic matter produces CO2 but also the harder water is higher in Calcium Bicarbonate, which the plants can use as a source of carbon through biogenic decalcification. The plant absorbs the Calcium Carbonate (CaCO3) and breaks down the molecule to get at the Carbon. This is why we often see Ca deposits on plants that are getting their Carbon by this mechanism.

Also worth mentioning is the high solubility of CaCO3 (Calcite/Limestone) in water. Again, due to the auto-ionization of water the acidic Hydronium (H3O+) interacts with the CaCO3 to form calcium, water, and CO2.

1. CaCO3 + H3O(+) --> Ca + H2O + CO2
2. CO2 + H20 --> H2CO3 etc.

This is why plants can do so well in areas where limestone is available in close contact with surface water or where groundwater flows through limestone, ie: Florida.


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## timdz (Mar 31, 2005)

This again brings up my question of carbonate hardness (KH) test kits. When testing the KH (with Aquarium Pharm, Tropic Marin kits), I assume you are basically adding an acid to a base and see how many drops it takes to cause a color change.

For example, when water goes through a water softener, cations such as Ca+ and Mg+ are are removed and replaced with another cation: Na+.

Baking Soda (Sodium bicarbonate) will also raise the carbonate hardness of the water, thus creating NaCO3?

"Decomposition of organic matter produces CO2 but also the harder water is higher in Calcium Bicarbonate, which the plants can use as a source of carbon through biogenic decalcification. The plant absorbs the Calcium Carbonate (CaCO3) and breaks down the molecule to get at the Carbon. This is why we often see Ca deposits on plants that are getting their Carbon by this mechanism"

So mainly when referring to 'hard' water, CaCO3 isn't the only product increasing the alkalinity. If someone wanted to use RO/DI water mixed proportionately with water through a water softner it is possible to obtain an appropriate dKH', and adding the calcium and magnesium with appropriate buffers to raise the general hardness. But the KH reading is coming from sodium bicarbonates. 
Therefore, when using a KH test kit how can we be sure that we have the 'right' kind of carbonates/bicarbonates in the water? And what would happen to the Na+?

My favorite way of having a successful planted tank is to start from scratch (RO/DI) and adding my personal favorite buffers from SeaChem to adjust GH and alkalinity. But when using well water diluted with RO/DI water I am not sure I feel comfortable with what my KH test kit is actually testing............


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## cS (Jan 27, 2004)

Phil Edwards said:


> Unless kept under artifically high pressures CO2 gas is not present in aqueous solution. Instead, it is absorbed and converted into Carbonic Acid, which then undergoes some pretty spiffy chemistry.
> 
> In our tanks with CO2 gas supplimentation the CO2 is quickly converted to Carbonic Acid by interaction with water which then reacts with other free water molecules to form Bicarbonate and Hydronium ions. These steps are what happens in our reaction chambers and is how CO2 is absorbed and held in aqueous solution.


Phil. Not all of the carbon dioxide (CO2) we inject into our tanks become carbonic acid (H2CO3). When gaseous CO2 is added to water, it becomes CO2(aq), a hydrated species.

*CO2(g) => CO2(aq)*

This aqueous carbon dioxide is what plants uptake. Only a _*very*_ small amount of CO2(aq) goes on to react with water (H2O) to become carbonic acid (H2CO3).

*CO2(aq) + H2O(l) => H2CO3(aq)*

H2CO3 is a weak acid so it will disassociate as follow:

*H2CO3(aq) <=> H+(aq) + HCO3-(aq) <=> H+(aq) + CO3(2-)(aq)*

But at the pH range relevant to the typical planted tank, HCO3- is the dominant species. So, it's basically:

*H2CO3(aq) <=> H+(aq) + HCO3-(aq)*

The liberation of the hydrogen proton (H+) lowers pH, which is why we see a decrease in pH as a result of CO2 injection. Furthermore, the equilibrium between H2CO3 and HCO3- (bicarbonate) is then employed by hobbyists to calculate the concentration of CO2(aq) in the water (pH/KH/CO2 chart).

---

In nature, three main sources of carbon are derived from (1) the atmosphere, (2) organic decomposition by bacteria, and (3) KH. (1) and (2) account for only a few ppm CO2 so plants in nature are usually limited by CO2. Mother nature, in her infinite wisdom, gifts many plants with the ability to utilize the water's KH (the measure of the concentration of HCO3- [ok, CO3(2-) (carbonate) as well :mrgreen:]). This extraction is energetically expensive so plants do not resort to this unless CO2 concentration is low. *THEREFORE, for plants possessing this ability, high KH water is beneficial.*

---

What I do not understand is that why some plants (soft water plants) are adversely affected by the presence of HCO3-. It doesn't make much sense in my head; but at least we are no longer blaming high Ca/Mg as the cause for these plants' poor health.


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## Phil Edwards (Jan 22, 2004)

cS,

Seeing as you're actually a chemist would you please tell me how CO2 would dissolve into H2O without becoming Carbonic Acid? Seriously, I always thought that to be soluble there had to be an exchange of e- between compounds or a "taking" of e- (ie: van der Waal's) that created some sort of bonding between the two. The only way I can think of is a hydrogen bond between the two unless somehow there are London Dispersion forces holding them together at an H-C interaction. However, the H bonding seems the much more plausible and stable interaction to me. 

The only practical stuff I have on this type of chemistry is out of my environmental science books and they all only mention the CO2 + H2O --> HCO3 reaction. I just happened to try writing out the reactions step-wise to see how all the stuff moved around, which I blame my Org. prof for getting me into the habit of doing.  

By the way, I LOVED the EluciDate forward you posted in "Faces..." way back when.


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## SCMurphy (Jan 28, 2004)

cS said:


> This aqueous carbon dioxide is what plants uptake. Only a _*very*_ small amount of CO2(aq) goes on to react with water (H2O) to become carbonic acid (H2CO3).
> 
> *CO2(aq) + H2O(l) => H2CO3(aq)*


 That would be the normal situation, but we are pumping CO2 into the tank. We are overloading the CO2(aq) side of the equation and pushing the CO2 to dissolve into carbonic acid and we see that in the pH shift.


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## Edward (May 25, 2004)

cS said:


> What I do not understand is that why some plants (soft water plants) are adversely affected by the presence of HCO3-. It doesn't make much sense in my head; but at least we are no longer blaming high Ca/Mg as the cause for these plants' poor health.


I don't think it's the HCO3 presence as much as the absence of Ca and Mg in other forms, as sulfates and chlorides. Soft water plants seem to prefer those more.

Edward


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## cousinkenni (Jan 24, 2005)

Phil Edwards said:


> cS,
> 
> Seeing as you're actually a chemist would you please tell me how CO2 would dissolve into H2O without becoming Carbonic Acid? Seriously, I always thought that to be soluble there had to be an exchange of e- between compounds or a "taking" of e- (ie: van der Waal's) that created some sort of bonding between the two. The only way I can think of is a hydrogen bond between the two unless somehow there are London Dispersion forces holding them together at an H-C interaction. However, the H bonding seems the much more plausible and stable interaction to me.
> 
> ...


Phil,

I don't know if this is relevent to your question, but think about other gases than C02. Take Nitrogen for example. Atmospheric nitrogen is relatively inert. but it disssolves relatively easily in our blood. At 1Atm our blood holds a certain amount. If you increase the pressure (i.e. through diving) more N2 will dissolve.

Don't always think of dissolving as a dissociation (NaCl -> Na+ Cl-). That only happens mainly to salts. Think of dissolving as a single molecule of whatever being surrounded by water molecules on all sides. Take sugar for instance. In solid state, sugar is next to sugar which forms a crystal structure. If you place this crystal sugar in water the individual sugar molecules break apart because water surrounds it. The crystal (which is more structured) turns into moving chaos (dissolved). It is basically entropy at work. Entropy is also why there is a loss of heat (energy) through evaporation to get back to the crystal structure (you have to keep the laws of physics happy).

Sorry to go off on a tangent. I though it might help you understand it better

Ken T.


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## cousinkenni (Jan 24, 2005)

SCMurphy said:


> That would be the normal situation, but we are pumping CO2 into the tank. We are overloading the CO2(aq) side of the equation and pushing the CO2 to dissolve into carbonic acid and we see that in the pH shift.


SCmurphy,

Cs actually makes note of this in her statment.

Ken T.


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## cS (Jan 27, 2004)

timdz said:


> This again brings up my question of carbonate hardness (KH) test kits. When testing the KH (with Aquarium Pharm, Tropic Marin kits), I assume you are basically adding an acid to a base and see how many drops it takes to cause a color change.


Timdz. Alkilinity test kits measure the concentration of the anion of weak acids. In our carbonate buffered system at the operative pH, H2CO3 is the weak acid whose anion is HCO3- (see above post). These test kits use some kind of strong acid to do so -- represented below as HX.

*HCO3- + HX => H2CO3*

Then they use some fancy mathematics to convert how much HX you used into ppm CaCO3, which is then converted further into degrees KH.

_It's important to take note that "ppm CaCO3" is used here as an "unit of measurement". We are NOT actually testing for CaCO3.  Think of it as a conversion factor: how much CaCO3 would we need to add to produce the same KH reading as what you have just measured._



> For example, when water goes through a water softener, cations such as Ca+ and Mg+ are are removed and replaced with another cation: Na+.


Water softeners are not recommended for the planted tank since it replaces Ca/Mg with Na. Effectively, it reduces the GH and raises the salinity to toxic levels. If you want to use "pure" water, then use RO or RO/DI water.



> Baking Soda (Sodium bicarbonate) will also raise the carbonate hardness of the water, thus creating NaCO3?


No. Sodium bicarbonate is NaHCO3. It dissolves in water into sodium ions (Na+) and bicarbonate ions (HCO3-). The HCO3- is what actually raises carbonate hardness (KH). The Na+ just tags along. For most people, the Na+ coming from baking soda is neglible. Such is not the case with the use of water softeners, which have the potential to raise Na+ quite high.



> So mainly when referring to 'hard' water, CaCO3 isn't the only product increasing the alkalinity. If someone wanted to use RO/DI water mixed proportionately with water through a water softner it is possible to obtain an appropriate dKH', and adding the calcium and magnesium with appropriate buffers to raise the general hardness. But the KH reading is coming from sodium bicarbonates. Therefore, when using a KH test kit how can we be sure that we have the 'right' kind of carbonates/bicarbonates in the water?


I am not sure that I understand what you're asking but I'll try to clarify. Think of GH and KH as separate entities. For all intents and purposes, they do not concern one another, unless you adding a substance that raises both the GH and KH, such as CaCO3 & CaMg(CO3)2.

*GH is the measure of the concentration of primarily Ca2+ and Mg2+ ions.*
To raise GH, one can use CaCO3, CaCl2.xH2O, CaSO4 (gypsum), MgSO4.7H2O (Epsom salt), CaMg(CO3)2 (dolomite), or any other products containing "Ca" and "Mg" in the formula.

*KH is the measure of the concentration of HCO3- and CO3(2-) ions.*
To raise KH, one can use CaCO3, CaMg(CO3)2 (dolomite), NaHCO3 (baking soda), KHCO3, or any other products containing "CO3" or "HCO3" in the formula.

Now, if you look at the above listing of chemicals, you'll notice that CaCO3 and CaMg(CO3)2 appear in both places. That's because if you look carefully at their chemical formulas, you'll notice that they contain elements that affect both the KH and GH.

CaCO3: the Ca raises the GH while the CO3 raises the KH.
CaMg(CO3)2: the Ca and Mg raise the GH while the CO3 raises the KH.



> And what would happen to the Na+?


The Na+ goes on to increase the salinity of the water. For most folks, this increase is neglible, so we don't talk about it.



> My favorite way of having a successful planted tank is to start from scratch (RO/DI) and adding my personal favorite buffers from SeaChem to adjust GH and alkalinity. But when using well water diluted with RO/DI water I am not sure I feel comfortable with what my KH test kit is actually testing.


I am assuming that the Seachem buffer you are using is Seachem Equilibrium? If so, then you can trust your KH test kit. Just don't start adding any "pH Up" and "pH Down" products and you'll be fine. Use your CO2 injection to adjust that.

I hope that things are a bit clearer?


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## SCMurphy (Jan 28, 2004)

cousinkenni said:


> SCmurphy,
> 
> Cs actually makes note of this in her statment.
> 
> Ken T.


 Yes, I know, that's what I'm getting at. Her statement that not much CO2 dissolves is a common one based on what happens in nature where most dissolved CO2 is in the *CO2(aq) *form. When we overload the equation the way we do, a lot more CO2 dissolves that what would be the norm. You know, the old 'overload one side of a chemical equation and watch the shift to the other side...'


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## cousinkenni (Jan 24, 2005)

SCMurphy said:


> Yes, I know, that's what I'm getting at. Her statement that not much CO2 dissolves is a common one based on what happens in nature where most dissolved CO2 is in the *CO2(aq) *form. When we overload the equation the way we do, a lot more CO2 dissolves that what would be the norm. You know, the old 'overload one side of a chemical equation and watch the shift to the other side...'


Sean,

I thought that CS's post and the first post in the section were saying that most plants IN NATURE use HCO3- not CO2. The reason for this is that CO2 is not readily available. CS mentions that the use of HCO3- requires more energy than using CO2 but that energy has to be expended for the life of the plant.

If you see a big pH shift, then yes, the CO2 from the left side of the equation is driving the release of H+ protons. You will have to measure the pH before CO2 addition (no plants present) and then after CO2 addition (no plants) to see how much the CO2 is affecting the equation.

The plants may keep the pH at a stable point because they are constantly taking from the left side of the equation as we add to the left side....correct? This doesn't necessarily mean that the plants are using the HCO3- (IN OUR TANKS).

When we add sodium bicarbonate to the system it is acting as a buffer so we can obtain the pH we desire correct? It just so happens that in our case as we add more buffer we increase the potential that the water has to hold carbon in a non CO2 form. As the plants diminish the CO2(aq) it can be replaced by the balancing of the equation.

So whether the plants IN OUR AQUARIUMS use CO2(aq) or HCO3- I am not sure, but from what CS says it sounds like they use CO2(aq).

Any info on these points would be greatly appreciated.

Thanks,

Ken


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## plantbrain (Jan 23, 2004)

Phil, about a ratio of H2CO3:CO2 is roughly 1:400 in seawater, I forgot what this ratio is in FW and it has other factors involved at arriving at that ratio.

As far as CO2, well, that is what the plant wants, if it's limiting, then , if they have the ability, the weeds will go after the bicarb.

Plants and algae have several methods they can use to do this.

Some algae and plants are much better at concentrating carbon sources than others. So CO2 competition often occurs when lower amounts exist in our tanks, in nature etc. 

How to solve this issue?
Add excess CO2.

I think the main things is the practical matter and keep your eye on the aquatic weed's needs.

If you have the abilty to maintain stable amounts of CO2, you can limit it and see how certain species respond vs others. You can see some plants grow well, while others rot. Adding CO2 to a non limiting level, then both grow equally as well.

I'm not so sure that "most" plants use HCO3 as the lion's share in nature, a few certainly do, but many use the CO2 that's there in early part of the day and then they stop when the concentration falls below a critical level, flowing waters also are quite different and don't have pronounced diel variations in CO2 levels.

If the body of water is a lake, the HCO3 will quickly be removed if the lake is heavily infested with weeds.

Then what?

The total carbon issue will show that plants will grow better in harder water since the harder water has more total carbon, but the plants will not use it if there is ample CO2 available for the light intensity.

Regards, 
Tom Barr


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## HeyPK (Jan 23, 2004)

To add to what Tom said, many of our newer and fussier aquarium plants can't use bicarbonate. The really old aquarium plants, such as Elodea, Eigeria, Ceratophyllum, Vallisneria, the earliest Amazon swords (I think) can utilize bicarbonate. Crypts, Anubias, etc can't utilize bicarbonate and really do need CO2 fertilization. I suspect that these kinds of plants may have high levels of CO2 in their natural environments. They live in streams where the water may come largely from ground seeps and may be quite high in CO2. Water that seeps out of the ground or comes out of swamps usually has high CO2, iron and manganese levels.


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## cS (Jan 27, 2004)

> Seeing as you're actually a chemist


Phil, you're too kind; but my job has NOTHING to do with chemistry whatsoever. If you want your nails done and gossip, then please have a seat. :mrgreen:
It's people like Roger Miller, Paul Sears, Paul Krombholz (HeyPK), cousinkenni, etc. who are the true chemists.







I simply regurgitate information that they have presented elsewhere because there is just no way that I would be able to intuit any of this on my own.



> would you please tell me how CO2 would dissolve into H2O without becoming Carbonic Acid? Seriously, I always thought that to be soluble there had to be an exchange of e- between compounds or a "taking" of e- (ie: van der Waal's) that created some sort of bonding between the two. The only way I can think of is a hydrogen bond between the two unless somehow there are London Dispersion forces holding them together at an H-C interaction. However, the H bonding seems the much more plausible and stable interaction to me.


Cousinkenni summarized it even better than I can Phil. Thanks Kenni. :mrgreen: But I suppose you want the actual mechanism by which it happened? I'll make an attempt to describe the chemistry involved. This is all from memory so I'm bound to make mistakes. So to all the real chemists out there, _please_ double check.

Carbon dioxide (CO2) is electrically neutral (non-polar) due to its symmetrical molecular structure.










Water (H2O), on the other hand, is electrically charged (polar) due to its assymetrical molecular structure. There's a slight negative charge at the oxygen end and a slight positive charge at the hydrogen end.










If nothing else, then I think anyone who walks away from a chemistry class would remember the "like dissolves like" rule of thumb. CO2, being non-polar, is generally not soluble in H2O, a polar substance. However, there is a _slight_ solubility between CO2 and H2O at standard conditions, mostly due to the intermolecular forces induced by the H2O molecules. Which one(s)? I have no idea what the proper term(s) is; but this is how I think it works, which could very well be quite wrong. No one looks while my proverbial pants are down ok? 8-[

Ok, imagine the polar H2O molecule approaching the non-polar CO2 molecule. Now normally, the electron cloud of CO2 is symmetrically distributed between the 2 oxygen atoms. But when the negative end of the H2O molecule approaches the oxygen atom on the CO2 molecule, the eletron cloud on the CO2 oxygen moves away in order to reduce the negative-on-negative interaction (repulsion of like charges). As a result, the non-polar CO2 molecule temporarily experiences a dipole moment where there is a separation of charges. This separation causes the CO2 molecule to become weakly attracted to the H2O molecule (opposite charges attract). CO2(aq) denotes this state of hydration. If the attraction is strong enough then you'll get the formation of H2CO3 which very quickly degrades into H+ and HCO3- ions.

Because this attraction is so weak, pressure and temperature play a more important role in how much non-polar gases can dissolve in a given volume of water. Let's ignore temperature and focus on pressure for the sake of brevity. As you increase pressure, you force more CO2 to dissolve. This accounts for the carbonation process in soda drinks. When you open a soda can, the pressure drops dramatically so all that dissolved CO2 leaves the sugar solution and escapes into the atmosphere. But who cares about soda drinks, let's look at how all this applies to our tank:

We dissolve CO2(g) into water by turbulence in our reactors. In effect, we are supersaturating the water with CO2. *CO2(g) => CO2(aq)* This equation is not exactly true because the relationship is in equilibrium, NOT to completion as I have depicted previously. *CO2(g) <=> CO2(aq)*

Now, as mentioned above, we are operating under standard condition of 1 atm of pressure over the surface of the water. And at this pressure, the equilibrium favors CO2(g) (recall how weak the intermolecular attraction between CO2 and H2O is). So this means that all that 30 ppm of CO2(aq) is constantly trying to escape as CO2(g). You can test this by taking a water sample from your tank and let it stand for a couple of hours. Test its CO2 concentration now. You'll find that there's only a couple ppm CO2 left. This process is accellerated by surface turbulence; ergo the often heard advice to reduce surface turbulence in a planted tank to conserve CO2.

Therefore, to compensate for the loss of CO2(aq) to the atmosphere and those consumed by plants/algae, we have to constantly pump CO2(g) in order to drive the CO2(g) <=> CO2(aq) equilibrium (_Le Chatelier's Principle_) to the right so that adequate CO2(aq) can be maintained. In effect, we are making the rate at which CO2(g) enters the solution much higher than the rate at which CO2(aq) leaves.

I hope that things are a bit clearer now Phil?











> By the way, I LOVED the EluciDate forward you posted in "Faces..." way back when.


::laughs:: It's worth it to suffer through your family/friends/co-worker's daily "send this to five people or you'll burst into flames" emails for the rare gems that pop in occasionally. ::laughs::


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## cS (Jan 27, 2004)

cousinkenni said:


> I thought that CS's post and the first post in the section were saying that most plants IN NATURE use HCO3- not CO2. The reason for this is that CO2 is not readily available. CS mentions that the use of HCO3- requires more energy than using CO2 but that energy has to be expended for the life of the plant.


You've taken me out of context Ken. My apologies for not being clearer. I was replying to Phil's post, which focuses on CO2 and KH. I used that to clarify some things as well as addressing WHY certain plants possess the ability to extract CO2 from the KH and why high KH is therefore beneficial to them.  Allow me to summarize as well as addressing what Sean is trying to say.

*Most of the CO2(g) injected into our tanks take on the CO2(aq) form. Of these CO2(aq), X% becomes H2CO3 (which degrade into blah blah blah). So as we add more and more CO2(g), then we get more and more CO2(aq). And as more and more CO2(aq) is formed, we get a greater QUANTITY of H2CO3. The percentage stays roughly the same: X%. The kinetics is explained in the above post. I hope it makes things clearer.*

But as Tom has already pointed out, CO2 is the primary and preferred source of carbon for aquatic macrophytes. But relative to our home aquaria, the concentration of dissolved CO2 in nature is much less (thus the comment on CO2 being the limiting agent in nature, givent he high level of light). Mother Nature can't just purchase a CO2 cyclinder. She's limited to the air, which contains relatively little CO2. Furthermore, the distribution and primary source of CO2 in nature is, as I understand it, quite varied depending on the habitat under discusion. Let's compare and contrast three such habitats: streams, Florida lakes, and the Amazonian black water river systems. I should note here that the following are regurgitation of information I've read elsewhere as well as my own speculation, so take it with a grain of salt and correct me where necessary please. 

*STREAMS*
In streams, your primary source of CO2(aq) comes from the air. The sheer turbulence of the water is synonymous with the turbulence created in our CO2 reactors. The net result is the equilibrium concentration of only a couple ppm CO2(aq). These streams are populated by various algae and mosses. CO2 coming from KH and organic decomposition play little role here I would imagine. Streams aren't particularly high in KH since the water has yet to dissolve much minerals. Well, I suppose that depends how far upstream we're discussing.  CO2 coming from organic decomposition would be virtually nil I am guessing since streams are very low in organics and any CO2 created would be overpowered by the turbulence. So here, macrophytes feed mainly on dissolved CO2 obtained from the air.

*FLORIDA LAKES*
Florida lakes are influenced by 4 sources of CO2 acquisition: from the air, from organic decomposition (substrate), spring water, and KH. CO2 from the air dissolves in water through wave actions so it's confined to the upper water layer. The deeper layers in deeper lakes rarely mix. So it's not surprising that many of Florida's shallow lakes have lots of plants. Secondly, lakes are stationary and accumulate a lot of muck. The substrate decomposes, releasing CO2 in the vicinity of plants for immediate uptake. Third, Tom often notes that many of Florida's lakes are spring fed. And as Paul noted earlier, spring water is loaded with CO2. Last but not least, Florida is riddled with limestone so the KH (carbonate content) is relatively high.

*AMAZONIAN BLACK WATER RIVER SYSTEMS*
These rivers are influenced primarily by CO2 from the air and organic decomposition in the substrate. Carbonate hardness is virtually absent in these systems due to all the humic acid from nearby forests. Springs in the Amazons? I have never heard of them. Maybe they exist. I don't really know. Nonetheless, the acidic soft water systems of the Amazon are where some of the "soft water plants" hail from.

---

There are just so many variables in natural water bodies. Many plants inhabit the shallow areas of these water bodies. They often exhibit both submersed and emersed leaves. These emersed leaves may be the primary sites of production and food is exported to the rest of the plant. Etc. And of course we are ignoring the algae in marine systems. All I know is that I need a vacation after all this talk. [smilie=l:


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## SCMurphy (Jan 28, 2004)

cousinkenni said:


> Sean,
> 
> I thought that CS's post and the first post in the section were saying that most plants IN NATURE use HCO3- not CO2. The reason for this is that CO2 is not readily available. CS mentions that the use of HCO3- requires more energy than using CO2 but that energy has to be expended for the life of the plant.
> 
> ...


 Yup, most of the marsh plants that we coax into living submerged are unable to use bicarbonate.

I never said 'big', I don't know what a big pH shift would be to you. You could measure the pH in the morning and afternoon and see the pH shift, unless you have a pH controller that works. I don't use a controller.

The plants are not keeping the pH at a stable point, the constant injection of CO2, the bicarbonate buffering system, and the degassing of CO2 is what creates the illusion of a stable point. The plants in the aquariums just use what they can as all this is going on around them. Think of it as the CO2 EI approach. Many plants might not use HCO3- directly, but if they do effect the pH via removal of CO2 and supplimentation is inadequate to make up the removal then the equations 'backup' and HCO3- returns to CO2 and H2O and the plants get to use the CO2.

Why to you add a buffer? Is it to prevent a 'big' pH shift? You are adding a buffer to KEEP the CO2 in a CO2 form, NOT to encourage it to dissociate. Why would you add bicarbonate to encourage CO2 to become bicarbonate?

I don't add sodium bicarbonate, I add Ca and Mg and let the bicarbonate form as the buffering system pulls CO2 in. The plants help out by adding some humics to the water and even more CO2 is pulled into reserve. The reason people have to create the buffer is they aren't using the draw method to pull CO2 into solution to form the buffer.

In Nature, CO2 is at equilibrium with the atmosphere, the fact that is has a pathway to dissociate as well as dissolve is why it is more solulable than other gases which only dissolve via partial pressure (cS does a nice job of describing this.) The more ions in the water, the more CO2 dissociates, the more bicarbonate is produced, the more CO2 is held in reserve.

Oh btw, Phil, good job!


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## Phil Edwards (Jan 22, 2004)

cS said:


> Carbon dioxide (CO2) is electrically neutral (non-polar) due to its symmetrical molecular structure.
> 
> 
> 
> ...


Well, for not being a chemist you certainly can translate what they say very well! 

You're pretty close there with the interactions between CO2 and H2O through dipole-dipole interaction (the delta + and delta - stuff) What's going on is the Oxygens are pulling harder on the available electrons than the Carbon is, creating an area with a slightly negative charge in character/behavior. Likewise on H2O the Oxygen is slightly negative while the Hydrogens are slightly positive in character. When the partial + and partial - meet they can interact to form a bond between the two molecules. In this case it could be as weak as a brief dipole-dipole bond or as strong as a Hydrogen bond, depending on the circumstances.

Those two forces could potentially cause CO2 to dissolve in an ideal environment. What I'm curious about is the solubility of the gas from CO2 (g) to CO2 (aq) in a system with an acidic pH. I went and talked with the Biogeochem prof at my school and he confirmed that in pH between 4.6-7.something CO2 is found in majority as HCO3-. When the pH drops below 4.6 then it is found as CO2.

Yes, these are things in a natural environment and not in the artifically induced conditions of are aquariums. Like Sean said, pumping CO2 into our tanks overbalances the equilibrium and it has to go somewhere. LeChatlier's Principle if I'm not mistaken.  On the opposite, if we want CO2 in the system to increase we have to overbalance in the product side of the reaction so we can then compensate with injection to balance it out at a level we want.

WOW! What a discussion!! I hope we're all learning something here and not getting mad at eachother. We've done a good job so far of keeping it on the level, let's keep it up.


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## cousinkenni (Jan 24, 2005)

Phil,

Now you have gone and confused me. You are basically saying that at the pHs we keep our tanks at, most of the CO2 is in the HCO3- form, while Tom and CS are trying to say that it exists more in the CO2(aq) form (Tom mentioned a 1:400 ratio in saltwater). So it seems we are at a stand still..........

What is the correct answer?

My other question for Tom is that you always say to keep the [CO2] at roughly 30ppm. How do you do this? Are you using the pressureized CO2 to cause a pH drop by supersaturaing the H2O? I use passive diffusion to add CO2 and my pH never drops below 6.8 (itusually stays between 6.8 and 7.0). My KH stays at roughtly 60ppm

Sould I use a reactor? I don't add anything to my water besides CaCl2 and MgSO4 (1tsp and 1/2 tsp respectively per water change in a 20 gallon tank).

Ken T.


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## plantbrain (Jan 23, 2004)

cousinkenni said:


> Phil,
> 
> Now you have gone and confused me. You are basically saying that at the pHs we keep our tanks at, most of the CO2 is in the HCO3- form, while Tom and CS are trying to say that it exists more in the CO2(aq) form (Tom mentioned a 1:400 ratio in saltwater). So it seems we are at a stand still..........
> 
> ...


I think you should keep things simple.
If your KH is 60ppm, add enough CO2 to get a pH of 6.5.
Use only CO2 to do this.
Passive diffusers stink.
Make a DIY CO2 reactor, I have one shown in detail on my site for the public to try out.

It's highly effective, easy to use, works with any source of CO2 and cost about 2-3$.

There ya go.

Now as far as what is in solution, it's CO2 gas at super staurated values relative to the air.

What is your GH? I doubt you need to add the CaCl2/MgSO4, but adding it will not hurt, but you don't need that much in a 20 gal tank.

Damn you Phil!
There, now we are no longer getting along.

Regards, 
Tom Barr


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## cousinkenni (Jan 24, 2005)

Tom,

Before addition of the Cacl2 and MgSO4 my GH is 140ppm with about 20ppm Ca++

Thanks for the advice I will build one this weekend  .

Kenni


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## HeyPK (Jan 23, 2004)

Another little interesting factoid that I picked up somewhere is that if you cram aquatic plants into a small volume of hard water (has Ca and Mg bicarbonate), give intense light and monitor until the plants have taken up all the CO2 they can, you will get a pH of about 9 if the plants can only take up CO2, but you will get a pH of 11!! :shock: if you have plants that can utilize bicarbonate.

If you have bicarbonate users sharing a tank with CO2 users, and the CO2 is limited, the bicarbonate users will win and the CO2 users will be loosers. Hygophila polysperma is a CO2 user, and when it can't get any, it gets a characteristic appearance where its leaves are slanted. http://www.aquaticplantcentral.com/gallery/showimage.php?i=221&c=3

And, it isn't just the other plants that can lose. You can lose fish if you have a well-lit tank with bicarbonate users and you are not paying attention to the CO2 levels. I once had a tank with a bunch of Eigeria densa where I noticed one dead zebra danio and the others looking sick. I measured the pH and got 9.5. After an emergency CO2 addition the other fish recovered.


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## MatPat (Mar 22, 2004)

All I can say is WOW. Reading and re-reading this and trying to figure it out is gonna keep me busy for the next few months. 

I am always amazed at the amount of "brain-power" and experience that exists on this site!


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## Phil Edwards (Jan 22, 2004)

Tom, I know you didn't mean that!


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## cousinkenni (Jan 24, 2005)

HeyPK said:


> Another little interesting factoid that I picked up somewhere is that if you cram aquatic plants into a small volume of hard water (has Ca and Mg bicarbonate), give intense light and monitor until the plants have taken up all the CO2 they can, you will get a pH of about 9 if the plants can only take up CO2, but you will get a pH of 11!! :shock: if you have plants that can utilize bicarbonate.
> 
> If you have bicarbonate users sharing a tank with CO2 users, and the CO2 is limited, the bicarbonate users will win and the CO2 users will be loosers. Hygophila polysperma is a CO2 user, and when it can't get any, it gets a characteristic appearance where its leaves are slanted. http://www.aquaticplantcentral.com/gallery/showimage.php?i=221&c=3
> 
> And, it isn't just the other plants that can lose. You can lose fish if you have a well-lit tank with bicarbonate users and you are not paying attention to the CO2 levels. I once had a tank with a bunch of Eigeria densa where I noticed one dead zebra danio and the others looking sick. I measured the pH and got 9.5. After an emergency CO2 addition the other fish recovered.


CS or any one else,

Can you explain HeyPK's post to me. I cannot figure out why the pH is raising so much. The only thing that I can figure is that the plants are somehow stripping the water of hydronium ions by using the bicarbonate. What I can't figure out is where or how the plants are doing this (what is the first half of the equation?). Please enlighten me/us with your omniescience in the form of a balanced equation 

Thanks

KT


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## Rolo (May 12, 2004)

KT - I believe when plants use bicarbonate the following reaction takes place: 

HCO3- -----> CO2 + OH-

OH- is left outside the plant and causes pH to rise. 

So really the plant is still using CO2, just indirectly. I was under the understanding though that few plants use biogenic decalcification (like those from African Rift lakes?).


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## BruceF (Aug 5, 2011)

Stumbled into this thread today. Well worth a read.


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## Zapins (Jul 28, 2004)

Without reading every reply I want to mention that CO2 to carbonic Cid to bicarb is an equilibrium reaction which means that it will never go to completion and there will always be some of each present in the water.


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