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DIC and pH

LMuhlen

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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.
 

dw1305

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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
 

jaypeecee

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Hi @LMuhlen

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

1659562485440.png


JPC
 

LMuhlen

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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.
 

dw1305

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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
 

Freshflora

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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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LMuhlen

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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.
 

_Maq_

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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.
 

LMuhlen

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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?
 

_Maq_

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@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?
 

Freshflora

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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.
 
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LMuhlen

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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?
 

dw1305

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Hi all,
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.
That is it.

Cheers Darrel
 

X3NiTH

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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.

:)
 

_Maq_

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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.
 

dw1305

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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
 
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