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DIC and pH

LMuhlen

Member
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23 Mar 2022
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Brazil
I've tried to read a bit about how pH determines which form carbon takes when dissolved, whether CO2/acid, bicarbonate or carbonate.

It came close to my knowledge limit for chemistry, but in the end I feel that I got the grasp of it, except that I still couldn't figure if this is relevant at all for aquariums.

That is because I couldn't find information on the dynamics of this equilibrium. Does any one know the order of magnitude for the time it takes for CO2 to turn into bicarbonate or carbonate at higher pH levels? When we inject gas at the rate that we do, and it is either nearly immediately consumed or degassed, does it really form bicarbonates in a relevant rate?

I have seen a substrate manufacturer announcing that its ability to reduce pH would mean that more CO2 is made available.
 
Hi all,
....... Does any one know the order of magnitude for the time it takes for CO2 to turn into bicarbonate or carbonate at higher pH levels? When we inject gas at the rate that we do, and it is either nearly immediately consumed or degassed, does it really form bicarbonates in a relevant rate?

I have seen a substrate manufacturer announcing that its ability to reduce pH would mean that more CO2 is made available..
Yes, it is a dynamic equilibrium, so it happens fairly quickly. The actual rate of reaction will depend upon the solubility of the carbonate source, so a really hard marble (metamorphic Ca CO3) will take longer to go into solution than an aragonite CaCO3 source, like a cuttle "bone'.

The other way around is almost instant, so if you add sodium bicarbonate (NaHCO3) to water that is already saturated with Ca++ and HCO3- ions then Ca CO3 will precipitate out, as a white plume, as soon as you pour the sodium bicarbonate solution in.

Cheers Darrel
 
Hi @LMuhlen

I assume you've seen the three states of this equilibrium pictorially shown as below:

1659562485440.png


JPC
 
Hi @LMuhlen

I assume you've seen the three states of this equilibrium pictorially shown as below:

View attachment 191965

JPC
Yes, this is what I'm wondering about. If this equilibrium is reached almost instantly, it seems very relevant if what we are interested in is CO2/H2CO3. At pH 6.5 we would have less than half of the carbon we inject available to plants and changing the pH would be very effective to increase CO2 availability. But if this equilibrium is relatively slow, maybe it wouldn't matter while we are injecting gas at high rates.

Hi all,
Yes, it is a dynamic equilibrium, so it happens fairly quickly. The actual rate of reaction will depend upon the solubility of the carbonate source, so a really hard marble (metamorphic Ca CO3) will take longer to go into solution than an aragonite CaCO3 source, like a cuttle "bone'.

The other way around is almost instant, so if you add sodium bicarbonate (NaHCO3) to water that is already saturated with Ca++ and HCO3- ions then Ca CO3 will precipitate out, as a white plume, as soon as you pour the sodium bicarbonate solution in.

Cheers Darrel

Either I didn't understand your answer or I wasn't clear with my question, but in our case we are adding CO2, not the carbonates. My concern would be losing CO2 to carbonates.

Assuming that the plants in question aren't proficient with using bicarbonates instead of CO2.
 
Hi all,
........Either I didn't understand your answer or I wasn't clear with my question, but in our case we are adding CO2, not the carbonates. My concern would be losing CO2 to carbonates.....
When you add CO2, you change the equilibrium point of the CO2 - HCO3- - pH relationship, so the very useful curves that @jaypeecee showed aren't relevant, because you have more DIC in total. How much of that Dissolved Inorganic Carbon is in the form of H2CO3- is given by the pH value.

I hope that makes sense?

I only have access to a phone at the moment, but if you search the forum for "Bouncy Castle" & "bromothymol blue" that will give you a more complete explanation.

Cheers Darrel
 
I think this is on point here.
Post in thread 'Lean dosing pros and cons'
Lean dosing pros and cons.

TLDR, high PH has nothing to do with the free CO2 in the water.

However, bicarbonates can in some way interfere with the uptake of nutrients like CO2 and cause certain chelates to break. Furthermore, higher KH water is typically associated with higher GH too, and high amounts of calcium can inhibit/interfere with the uptake of various nutrients as well.
 
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What is TLDR, please?

Free CO2 depends on pH.
TLDR, too long didn't read, is usually used to indicate that you will give the summarized version for people who don't want to read a long text.

Ok, so I first read that the leftmost curve in the graphic was for carbonic acid, and I was fine understanding that this wasn't relevant. Then I read elsewhere that the curve could be used for CO2 and carbonic acid interchangeably, and that made me question if I had understood it...

So if I understand correctly, the reason why this is all unimportant to us is that only the thousandth fraction of CO2 which turns to carbonic acid enters this equilibrium with other DIC forms, so most of our precious CO2 is free to exist diluted in water. Is this right?

It felt clear that there should be a reason why this whole thing wasn't right for us, since pH rarely matters in our tanks, at least not at the rate the curve would indicate, but I thought maybe the reason was that this equilibrium would be too slow and our constant supply of gas would make the other kinds of DIC forms irrelevant.
 
the reason why this is all unimportant to us is that only the thousandth fraction of CO2 which turns to carbonic acid enters this equilibrium with other DIC forms, so most of our precious CO2 is free to exist diluted in water. Is this right?
Unfortunately, no.
People are often misguided - like you - by the fact that only tiny fraction of so called free CO2 reacts to form true carbonic acid. It's better to forget this fact and focus your attention solely at CO2 and HCO3-. Imagine that their sum is constant. Then, if pH is 6.35, half of it will be in the form of CO2 and half HCO3-. But if pH is 7.0, then only 4.8 per cent remains in the form of CO2, while HCO3- forms the remaining 95.2 per cent.
Because of that, in acidic water CO2 is more available than in alkaline water. All plants are more or less able & willing to use HCO3- instead of CO2 - which all plants prefer. However, some plants are very unhappy in alkaline water, grow slowly, and easily succumb to parasites or various events which may cause plant's death. Yet other plants readily use HCO3- (bicarbonate) instead of CO2 (carbon dioxide), and as such they enjoy competitive advantage in alkaline waters.
 
Unfortunately, no.
People are often misguided - like you - by the fact that only tiny fraction of so called free CO2 reacts to form true carbonic acid. It's better to forget this fact and focus your attention solely at CO2 and HCO3-. Imagine that their sum is constant. Then, if pH is 6.35, half of it will be in the form of CO2 and half HCO3-. But if pH is 7.0, then only 4.8 per cent remains in the form of CO2, while HCO3- forms the remaining 95.2 per cent.
Because of that, in acidic water CO2 is more available than in alkaline water. All plants are more or less able & willing to use HCO3- instead of CO2 - which all plants prefer. However, some plants are very unhappy in alkaline water, grow slowly, and easily succumb to parasites or various events which may cause plant's death. Yet other plants readily use HCO3- (bicarbonate) instead of CO2 (carbon dioxide), and as such they enjoy competitive advantage in alkaline waters.
Ok, so I see that this is not open and shut. If indeed the entire population of carbon atoms participate in this equilibrium of CO2, carbonic acid, bicarbonates and carbonates, then my original question remains: how fast is this equilibrium reached? We are injecting CO2 at a high rate and the dissolved gas is being absorbed by plants and being degassed to the atmosphere at an equally high rate, forming an equilibrium that keeps the concentration at 30ppm, as an example.

At the same time, we have a daily cycle where the gas concentration is reduced to close to the equilibrium with the atmosphere.

With this constant injection of gas, is there time for this equilibrium of acid and bicarbonates to be reached? If it a fast reaction?
 
@LMuhlen , now you understand it correctly. But to calculate the kinetics of this equilibrium exceeds my abilities, I'm not a chemist by profession. Actually, I'd like to know, too. Perhaps someone will explain?
 
In the context of an aquarium all that matters is the following: the more CO2 you inject and dissolve into the water, the lower your PH will be. CO2 does not depend on PH, PH depends on CO2. You had it right in response to my original post—only a small fraction of DIC forms carbonic acid. I believe it is only 1%. New Insights on Carbonic Acid in Water. The nature of DIC follows Henry’s Law, and thus the proportion of carbonate and bicarbonate to carbonic acid does not dictate the amount of free CO2 in the water.
 
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In the context of an aquarium all that matters is the following: the more CO2 you inject and dissolve into the water, the lower your PH will be. CO2 does not depend on PH, PH depends on CO2. You had it right in response to my original post—only a small fraction of DIC forms carbonic acid. I believe it is only 1%. New Insights on Carbonic Acid in Water. The nature of DIC follows Henry’s Law, and thus the proportion of carbonate and bicarbonate to carbonic acid does not dictate the amount of free CO2 in the water.
If the DIC chart refers to total carbon, regardless of the portion that turns into carbonic acid, like Maq said and is said to be in some other sources as well;
If the pH of the tank is high for reasons other than those related to DIC, high enough that even after the CO2 injection it is still high;
If the rate at which CO2 turns into bicarbonate due to the high pH is fast;

If all that were true, it would seem logical that a significant part of the injected gas would quickly turn into bicarbonates, and we would need to inject even more CO2 to reach the same result, all the while generating lots of bicarbonates.

And this is the part where my head starts to hurt, because then pH is dropping because part of the OH- is being consumed to turn carbonic acid into bicarbonate, but the CO2 concentration isn't growing at full proportion, so the 1 pH drop would mean less CO2 concentration than expected, and not that it is harder to reach a 1 pH drop.

Where is the logic flaw? If it is the first point and only the carbonic acid matters to the DIC chart, why do some sources mention that CO2 and carbonic acid can be used interchangeably, and even calling the left curve CO2 instead of carbonic acid?
 
Rate of change is very rapid as addition of soluble compounds into water alters the electrical neutrality of water (always neutral), if there is an upset to the equilibrium of anions to cations then charge balance has to be made to bring the water back to neutrality. The consequence of this is precipitation of compounds and in this particular case it’s charge balance via precipitation between Hydrogen CO2 and Oxygen.

:)
 
CO2 does not depend on PH, PH depends on CO2.
In fact, it works both ways, depending of which factor is a constant and which is a variable. You take for a variable CO2 content. Then yes, pH goes down with higher content of CO2.
However, if you add strong acid to the water and do not change the volume of DIC, then CO2 content increases and HCO3- decreases. In that way, CO2 depends on pH.
Yet there's one more way to see it. In an ideally open system the CO2 content is constant (reflecting partial pressure of CO2 in air) up to about pH=8.3, which is the point where all CO2 turns into bicarbonates.
 
Hi all,
You take for a variable CO2 content. Then yes, pH goes down with higher content of CO2.
However, if you add strong acid to the water and do not change the volume of DIC, then CO2 content increases and HCO3- decreases. In that way, CO2 depends on pH.
That is true generally, because it is back to the <"Bronsted Lowry definition of acids">, that they are proton (H+ ion) donors.
Yet there's one more way to see it. In an ideally open system the CO2 content is constant (reflecting partial pressure of CO2 in air) up to about pH=8.3, which is the point where all CO2 turns into bicarbonates.
<"It is that one">, @Manuel Arias, gives the <"chemistry in this post">. In <"this specific case">, when we add CO2, we are simulating an <"atmosphere richer in carbon dioxide (CO2)">, when we turn the CO2 off we return to the present CO2 ~ carbonate ~ pH equilibrium. I still like the <"Bouncy Castle"> analogy

Have a look at <"Circulation? Low flow? Planted tank with a betta..."> for more about pH.

cheers Darrel
 
After the read on @_Maq_ 's article on water parameters, I was brought back to this question of mine, so I went digging again through google to see if I could find some answers to the parts that, in my head, remains unresolved.

One of them is why it seems that pH plays such a strong role on CO2 availability (determining the proportions in which the different DIC options stay in equilibrium), but in practice we don't seem to observe it so strongly. Higher pH (7~7.5) injected tanks seem to do just fine, even though most of the CO2 should turn into bicarbonates. I thought that maybe it was a matter of the speed of reaction, which doesn't appear explicitly when looking at the equilibrium results. In an injected tank, we have new CO2 added constantly, and we lose most of this to the atmosphere. When we have CO2 turning from gas to diluted to outgassed, does it have time to react to bicarbonates?

I found this article:

I don't have the critical capabilities to look for any problems with it and, while it tackles an issue slightly different than ours (it covers the DIC rate of transformation when paired with its pH fluctuation, while in our tanks, for most of the day, we have a pH equilibrium paired with a constant supply of CO2), it seems to show that, while the DIC forms reach equilibrium between themselves in an order of magnitude of a second or less, the transformation from diluted CO2 to other DIC forms is in the range of 2 to 4 minutes.

It seems possible to me, considering this information, that given a constant CO2 supply, the actual CO2 equilibrium point in our tanks at higher pH is at a much higher concentration than what the DIC phase charts would have us at, making the increase of pH less significant to us from the perspective of making CO2 less available.
 
Hi all,
Higher pH (7~7.5) injected tanks seem to do just fine, even though most of the CO2 should turn into bicarbonates. I thought that maybe it was a matter of the speed of reaction, which doesn't appear explicitly when looking at the equilibrium results. In an injected tank, we have new CO2 added constantly, and we lose most of this to the atmosphere. When we have CO2 turning from gas to diluted to outgassed, does it have time to react to bicarbonates?
I'm not entirely sure. Once you have more than ~2dKH (~ 44 mg / L HCO3-) a one unit in pH fall indicates about 30 ppm CO2, despite the base 10 logarithmic nature of the pH scale. @Manuel Arias gives the full explanation in <"Drop checker color relation to KH question">.
it seems to show that, while the DIC forms reach equilibrium between themselves in an order of magnitude of a second or less, the transformation from diluted CO2 to other DIC forms is in the range of 2 to 4 minutes.

It seems possible to me, considering this information, that given a constant CO2 supply, the actual CO2 equilibrium point in our tanks at higher pH is at a much higher concentration than what the DIC phase charts would have us at, making the increase of pH less significant to us from the perspective of making CO2 less available.
I'm guessing that is right, but it is purely a guess.

cheers Darrel
 
Quite a few variables though, what acid, how much , how much HCO3, is the system open to CO2 escape?
If you were to adjust the pH of 4KH water with a strong acid to quickly break apart all HCO3 you would get 62 ppm CO2... but it would be a one time deal and better make sure to keep all that CO2 in the system if you want to see a change in the drop checker.

And in this specific case you would be right to use the figure posted by jaypeecee. You have a closed system, a fixed total inorganic carbon content (Ct) and play between species by changing the pH. Much like the ph-KH-CO2 chart is correctly read for pure water-bicarbonate systems with CO2 injection. Just plugging in KH and pH for a running tank without injected CO2 will give some outlandish values mainly due to organic acids!

In a system where you are continuously adding and loosing CO2, you are changing the Ct with one carbonate species which would make the distribution go to the left side (H2CO3/CO2). What fraction of the Ct is 30 ppm CO2/H2CO3 will move the point of the H2CO3 to the left or right, and thus the pH if the system pH is mainly driven by carbonates, H+/OH- . The Ct of the system should always be taken into consideration.

Consider: Ct is made mainly from 87 ppm HCO3 and 30 ppm H2CO3/CO2 the fraction -> fraction of H2CO3 is 0.3.
Ct is made mainly from 21 ppm HCO3 and 30 ppm H2CO3/CO2 the fraction -> fraction of H2CO3 is 0.6.

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.
 
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