DIC and pH

[LMuhlen]

Some very small part of H2CO3 will dissociate into bicarbonate in a reversible reaction. 30 ppm CO2 will increase HCO3 concentration by about 0.00002 ppm, lowering back the CO2 concentration will reverse this reaction, unless you have a base like CO3- or OH- available.

This is the part that I've been receiving conflicting feedback on.

If I understand you correctly, you are saying that a tiny fraction of the CO2 will become bicarbonate in an equilibrium. If this is the case, this entire discussion is pointless, since the effect of pH into the CO2/bicarbonate equilibrium would affect a tiny fraction of the total CO2 population, which would make this a non-issue.

On the other hand, @_Maq_ is saying that the entire CO2 population is participating in this equilibrium where pH dictates the fraction of CO2 that becomes bicarbonate. This would mean that at pH7 we would have roughly 80% of the CO2 assuming the form of bicarbonate, which is very significant.

I confess that the second understanding seems unaligned with our observation of planted tanks, it would be very noticeable if we had to inject 3~5 times more CO2 in order to reach the same result in tank with a pH of 7, compared to a lower pH. That is why I'm trying to bend my thinking over itself to find a solution, which could be the difference between the dynamic situation vs the theory on a static equilibrium. However, in my brief search, I haven't found an explicit explanation in an article of how this really works. Maybe because it is such a simple thing for the experts that it is not worthy of elucidating.

 

[Ria95]

This is the part that I've been receiving conflicting feedback on.

If I understand you correctly, you are saying that a tiny fraction of the CO2 will become bicarbonate in an equilibrium. If this is the case, this entire discussion is pointless, since the effect of pH into the CO2/bicarbonate equilibrium would affect a tiny fraction of the total CO2 population, which would make this a non-issue.

On the other hand, @_Maq_ is saying that the entire CO2 population is participating in this equilibrium where pH dictates the fraction of CO2 that becomes bicarbonate. This would mean that at pH7 we would have roughly 80% of the CO2 assuming the form of bicarbonate, which is very significant.

I confess that the second understanding seems unaligned with our observation of planted tanks, it would be very noticeable if we had to inject 3~5 times more CO2 in order to reach the same result in tank with a pH of 7, compared to a lower pH. That is why I'm trying to bend my thinking over itself to find a solution, which could be the difference between the dynamic situation vs the theory on a static equilibrium. However, in my brief search, I haven't found an explicit explanation in an article of how this really works. Maybe because it is such a simple thing for the experts that it is not worthy of elucidating.

Pretty much a non-issue in most planted aquariums as there is no significant additional source of a strong base. Some people are just good story tellers, no need to bend reality to match legend.

At ph 7 in a water-carbonate system, bicarbonates would be aprox. 80% of the total inorganic carbon content (Ct) not 80% of CO2. That s what the figure highlights, but most hobbyist become overly focused on CO2 and ignore the details. For example, water at a pH of 7 would have 82% of its Ct both at 10 KH , 35 ppm CO2 and at 1 KH, 3.5 ppm CO2.
If you want to inject additional CO2 in the water with 1 KH to reach 35 ppm, the fraction of HCO3 will change to 31% of Ct (same total amount, different fraction) and the pH will got to 6 . You can follow this on the graph as well. This 1 pH drop for 10x CO2 is no coincidence as the functions used to make the graph are essentially the same (and have the same limitations) as those used to determine the CO2 concentration in a water-carbonate system.

 

[LMuhlen]

Just as a curiosity, our future overlord ChatGPT seems convinced that the influence of pH on the concentration of dissolved gaseous CO2 in fresh water is much smaller than the influence of the partial pressure of the atmosphere. It says that pH influences the proportions of DIC - except gaseous CO2.

If you want to inject additional CO2 in the water with 1 KH to reach 35 ppm, the fraction of HCO3 will change to 31% of Ct (same total amount, different fraction) and the pH will got to 6 . You can follow this on the graph as well. This 1 pH drop for 10x CO2 is no coincidence as the functions used to make the graph are essentially the same (and have the same limitations) as those used to determine the CO2 concentration in a water-carbonate system.

I understand that when you get the 1 pH drop, you get the 10x CO2 concentration. My concern is if you need more or less or the same amount of CO2 to get this 1 pH drop starting from a higher pH water compared to starting from a lower pH water. If the pH influences significatly the CO2 concentration, you maybe would need more CO2 to get the 1 pH drop in a higher pH tank, since most of the CO2 injected would become bicarbonates instead of CO2. Or maybe less, because the pH drop itself would "release" CO2 trapped as bicabonates? Or both effects cancel each other?

Chat GPT seems convinced that the pH of the water is of much lesser significance than the rate of CO2 injected. It did take some insisting for it to converge to a final opinion, though.

 

[dw1305]

Hi all,

that the influence of pH on the concentration of dissolved gaseous CO2 in fresh water is much smaller than the influence of the partial pressure of the atmosphere. It says that pH influences the proportions of DIC - except gaseous CO2.

It's sort of halfway there. The atmospheric CO2 level controls the amount of (D)TIC (Dissolved(Total) Inorganic Carbon) in solution and the pH value is dependent upon the <"ratio of HCO3- and CO2"> (from the 0.15% <"dissolved as H2CO3"> which disassociated into a proton (H+) and bicarbonate ion (HCO3-).

At ~400 ppm atmospheric CO2, and if you have a buffer of carbonate (limestone ~ CaCO3 ~ <"calcite">), then the pH equilibrium point is ~pH8. Have a look at <"Geochemistry: Simple Conditions Of Calcite Precipitation — Canada (Ontario) Beneath Our Feet"> it is a good summary of the reactions in hard, alkaline (carbonate buffered) water.

Additionally if we have high atmospheric pressure, and cooler conditions, more CO2 (and <"all other gases">) will go into solution

1. increasing the pressure of CO2 gas in a system causes calcite to dissolve and go into solution;
2. decreasing the pressure of CO2 gas in the system causes calcite to precipitate from solution;
3. increasing the water temperature decreases the amount of CO2 dissolved in water and causes calcite to precipitate;
4. decreasing the water temperature increases the amount of CO2 dissolved in water and causes calcite to dissolve.

When you add CO2 to water you've just <"mimicked an atmosphere richer in carbon dioxide"> (CO2). <"http://www-naweb.iaea.org/napc/ih/documents/global_cycle/vol i/cht_i_09.pdf">

In terms of the pH it is the other way around, the pH is dependent on atmospheric CO2 levels (in an open system). Have a look at the <"Coke answer"> and @John q 's <"experiment with carbonated water"> - <"The pH of water from various sources: an overview for recommendation for patients with atopic dermatitis"> .

The carbonated drink scenario is a closed system, where CO2 can't escape into the atmosphere (until equilibrium is reached) but it held under pressure in solution and the <"greater amount of protons"> depresses the pH. In carbonated water, as soon as the bottle is opened (pressure is released), CO2 will escape and the pH will rise (apparently this doesn't work for Coke, because it also has phosphoric acid (H3PO4) in the recipe).

cheers Darrel

 

[Ria95]

I understand that when you get the 1 pH drop, you get the 10x CO2 concentration. My concern is if you need more or less or the same amount of CO2 to get this 1 pH drop starting from a higher pH water compared to starting from a lower pH water. If the pH influences significatly the CO2 concentration, you maybe would need more CO2 to get the 1 pH drop in a higher pH tank, since most of the CO2 injected would become bicarbonates instead of CO2. Or maybe less, because the pH drop itself would "release" CO2 trapped as bicabonates? Or both effects cancel each other?

1 pH drop will get you about the same 10x increase in concentration over a wide range of KH values. In this scenario KH is at least one of the major influencers of your pH value. You inject CO2, you lower the pH, a smaller % of the Ct is HCO3. See previous post about how little CO2 is converted to HCO3 and unless the aquarium has a source of readily available bases added it won’t keep that carbon as HCO3, the reaction will reverse. In this situation, carbonic acid is also not strong enough to break bicarbonates and release the CO2 from them. Try it yourself! Test KH , inject CO2, test KH , degass CO2, test KH, repeat

 

[LMuhlen]

See previous post about how little CO2 is converted to HCO3

This is the key point to me. That's how I thought it worked based on simple observation. Just like you said, we test KH and it doesn't change noticeably with or without CO2 injection. This makes it all simple and not an issue. However, it doesn't seem to be the unanimous understanding of DIC vs pH relationship. Some people say that at a higher pH we have no dissolved CO2, it is all HCO3- and some CO3--. That would sound like an issue.

At ph 7 in a water-carbonate system, bicarbonates would be aprox. 80% of the total inorganic carbon content (Ct) not 80% of CO2.

So you are saying, based on your 2 combined answers, that CO2 is not part of the total inorganic carbon content, right? Only H2CO3, HCO3- and CO3--


From a link in Darrel's post, it seems that the confusion originates from the fact that CO2d turns into H2CO3 very quickly and, therefore, for concerns regarding the availability of H2CO3 we can consider the entire CO2d reserve as H2CO3. However, the equilibrium keeps only a very small amount of CO2d as H2CO3. Some authors say that we can add CO2d and H2CO3 as a single entity for the calculations, but that doesn't mean that we don't have CO2, far from it.

But that post raised another concern for me. It is all about debunking the "3ppm CO2 as the base value" thing. If we don't have 3ppm, but instead close to 0.5ppm from the atmosphere equilibrium, then the 1 pH drop would put us at 5ppm CO2, and not 30ppm... How come we don't see significant discrepancies when comparing the 1pH drop and the drop checker methodologies?

 

[Ria95]

It depends on the authors of the articles you read if they offer separate values for CO2 and H2CO3 or if they handle both of them in one parameter and only mark it as ’ CO2’ for ease. Both are valid if it‘s made clear what they mean.

What I’m saying in the quote is that the 80% figure is calculated from the total inorganic carbon, CO2 is only a part of Ct. Maybe that’s the source of confusion. As modeled in the graph from the first page, in a closed system (no inorganic carbon enters or exists) you have a fixed Ct, playing with the pH (via HCl, NaOH) will move the % of the forms accordingly. If now in that same system, you stop playing with the pH (via HCl, NaOH), you start adding CO2, your Ct will also increase , CO2 will be a bigger part of Ct, HCO3 a smaller part, CO3 even smaller and you will see the pH starting to lower.

People say a lot of things, I’m not inclined to start bending reality to match legends. Up to them to explain what they mean and back it up.