Confusion concerning 0dKH and nitrification

[Heelllooo]

Hello everyone,

I'm a bit confused with what I read concerning tank running at 0 dKH.

6 month ago I changed the way I prepare my WC water. I was mixing my really hard tapwater with RO and I started just remineralizing the RO with Calcium and Magnesium compound. Effectively pushing my tanks water to 0 alcalinity (in the the 60L one, the aquasoil doesn't buffer anything since a long time). I've done it after having read and view on this forum that you can perfectly run a tank like this. I didn't encountered any issue and and it definitively improved my plants health.

My tanks are really low stocking level, heavily planted, weak or no filtration, no or low level CO2 injection (DYI jello method). I never mesure pH knowing it doesn't really make sense without buffering capacity.

Now what confuse me :

On one hand, It seems pretty accepted on this forum that the way my tanks run are totally acceptable and it doesn't represent a risk of crash. On the other hand, I see lots of discussion around dosing carbonate and bicarbonate ( like in this recent thread : Sodium and plants ). It seems also pretty accepted that bacteria consume kH when processing ammonia and nitrite.

What do we know about nitrification in water at 0dKH ?
Is nitrification limited by the lack of carbonate and bicarbonate and I would have issue with a higher bioload and less plantmass ?
If nitrification is not limited, what are the advantage to dosing kH (I don't see any apart for the obvious one like having fish needing hardwater) ?
What does nitrifying bacteria consume when they don't find kH or is it different species of bacteria (that don't need kH) who take the job of nitrification in those conditions ?
Would I have issue with pH crash if I was dosing high amount of CO2 ? The CO2/pH/dKH relation is still a bit mysterious for me.
I think I remember reading about other ions buffering the pH ? I dose a good amount of fertilizer, does it play a role ?

 

[aquanoobie]

What does nitrifying bacteria consume when they don't find kH or is it different species of bacteria (that don't need kH) who take the job of nitrification in those conditions ?

It may be different species but it does continue in low pH.

Would I have issue with pH crash if I was dosing high amount of CO2 ?

It's not the pH problem, it is the CO2 concentration causing the fish distress.

I think I remember reading about other ions buffering the pH ? I dose a good amount of fertilizer, does it play a role ?

Most fertilizers don't change pH or KH.

What do we know about nitrification in water at 0dKH ?

It still works.

 

[dw1305]

Hi all,

it may be different species but it does continue in low pH.

I think that it is it. Where ever TAN (NH3 / NH4+), or nitrite (NO2-), are present, in aerobic situations, <"there will be organisms"> that can, and will, utilise them.

We have a thread that may be of interest: - <"Correspondence with Dr Ryan Newton - School of Freshwater Sciences, University of Wisconsin—Milwaukee">.

it seems also pretty accepted that bacteria consume kH when processing ammonia and nitrite.

They do, but the requirement for carbonate hardness is much less than was originally thought. Have a look at <"When to add fish!"> and <"Nitrification in acidic and alkaline environments">.

This is from Gaofeng Ni, Pok Man Leung, Anne Daebeler, Jianhua Guo, Shihu Hu, Perran Cook, Graeme W. Nicol, Holger Daims, Chris Greening; (2023) "Nitrification in acidic and alkaline environments". Essays Biochem 11 August 2023; 67 (4): pp. 753–768.

........ Some of the highest rates of soil nitrification are found in acidic soils (pH < 5.5) that constitute approximately 30% of the world’s ice-free land [75]. The majority of these soils are naturally acidic, and approximately 5% are used for arable crops [75]. AOA are generally the predominant nitrifiers in such soils, which makes them a crucial subject for investigation [19,75,34]. The best-studied acidophilic AOA belong to the genus Ca. Nitrosotalea [42], which is currently represented by three cultivated strains (Ca. Nitrosotalea devaniterrae, Ca. Nitrosotalea sinensis and Ca. Nitrosotalea okcheonensis) with growth pH restricted to 4 and 5.5 (optima between 4 and 5) [42–45]. This is supported by a regional and global phylogenetic analysis of archaeal amoA gene distribution [34,76],....... To date, two acidophilic Nitrobacter strains have been isolated: Nitrobacter NHB1 from acidic heath soil [30] with an activity range of pH 5.0–7.5 and Nitrobacter strain IOacid from an acidic forest soil with maximal nitrite oxidation activity at pH 5.5 (range 4.1–7.2) [53]. Furthermore, canonical (only nitrite-oxidizing) and a comammox clade A Nitrospira strain have been enriched from acidic agricultural soils with activity from pH 4 to 6 [81].

This is an Aquaculture paper: <"Development of an aquaponics microbial inoculum for efficient nitrification at acidic pH - Applied Microbiology and Biotechnology"> Derikvand, P., Sauter, B. & Stein, L.Y. (2021) "Development of an aquaponics microbial inoculum for efficient nitrification at acidic pH". Appl Microbiol Biotechnol 105, pp. 7009–7021

They found (via DNA analysis) that at the lower pH values <"COMAMMOX Nitrospira"> were the principal nitrifying organisms.

........ Aquaponics biofilters enriched with comammox bacteria and adapted to function at pH 5.6 can be a desirable inoculum for freshwater recirculating aquaponics systems to retain nitrification activity and improve crop yields .........

• Microbial communities adapted from pH 7.6 to pH 5.6 retained 81% nitrification activity. • Microbial communities adapted from pH 7.6 to pH 5.6 were enriched in comammox bacteria.

cheers Darrel

 

[LMuhlen]

Most fertilizers don't change pH or KH.

While most of the discussion is centered around KH, and most tests claim to measure KH, our commercial aquarium tests are actually measuring alkalinity, as in the resistance against a pH drop. We assume that in our tanks most of the alkalinity is due to KH (carbonates and bicarbonates), but there are other less impactful sources that could become relevant in a 0 KH setting. There are phosphate based buffers, although I don't know if they can form from simple phosphate ions. There are tannins that supposedly buffer water as well. I believe there must be a lot of different chemical, and maybe even biological processes happening that have some degree of influence on pH stability.

 

[dw1305]

Hi all,

While most of the discussion is centered around KH, and most tests claim to measure KH, our commercial aquarium tests are actually measuring alkalinity, as in the resistance against a pH drop. We assume that in our tanks most of the alkalinity is due to KH (carbonates and bicarbonates),

We do. This thread may be of interest: <"Potassium Carbonate (K2CO3) vs pH ?">

There are phosphate based buffers, although I don't know if they can form from simple phosphate ions.

You could make your own <"phosphate buffers"> using the weak base dipotassium phosphate (K2HPO4.nH2O) and weak acid monopotassium phosphate (KH2PO4). There isn't any point, but it is possible.

cheers Darrel

 

[Andy Pierce]

"Carbonate alkalinity (CA = HCO3 − + 2CO3 2−) typically accounts for >95% of the total alkalinity in the ocean. Many studies (e.g., Broecker & Peng, 1982) use a simple form of alkalinity including only water and carbonate alkalinity terms. In seawater, a slightly more accurate expression is obtained when borate alkalinity is included as well." - Ocean Alkalinity, Buffering and Biogeochemical Processes

 

[dw1305]

Hi all,

Carbonate alkalinity (CA = HCO3 − + 2CO3 2−) typically accounts for >95% of the total alkalinity in the ocean

That is why I've tended to regard hardness (dGH) and alkalinity (dKH) as the two halves of the same whole. Unless you are in the Okavango Delta, or similar, all rivers lead the to the sea.

In the UK freshwater both parameters are nearly always derived from the <"dissolution of limestone"> (CaCO3) to give Ca++ (dGH) and 2HCO3- (dKH) in a 1 : 1 ratio.

If you live around the <"Mediterranean, or in the mid-west of the USA">, evaporite minerals come into play.

cheers Darrel

 

[hax47]

I am unsure if there is such a thing as 0 dKH alkalinity in aquariums. Theoretically, there will be always bicarbonate once CO2 is dissolved in the water. If no other buffers are present, then the amount is negligible, but once there are other buffers, the bicarbonate levels can go significantly up even in a supposed-to-be 0 dKH water. Significant here means compared to 0 dKH.
If I am not mistaken, nitrification does not consume bicarbonate directly, but bicarbonate buffers the produced H+ ions during the nitrification and is converted to water and CO2, which eventually dissipates. It is the bicarbonate that buffers the H+ ions because it is the most significant buffer in the aquarium, but theoretically, it might be something else as well.

What does nitrifying bacteria consume when they don't find kH or is it different species of bacteria (that don't need kH) who take the job of nitrification in those conditions ?

One candidate as buffer is the phosphate, but it is an effective buffer only in the pH range between about 6 and 8. However, the buffering capacity is very low under "normal" aquarium phosphate concentrations. 3 ppm contributes to about a maximum of 0.1 measured dKH, but that is probably only true around a pH of 8; at 7 it might be only half of it. The other candidates are some organic acids released from decaying plants and soil.

Would I have issue with pH crash if I was dosing high amount of CO2 ? The CO2/pH/dKH relation is still a bit mysterious for me.

pH is a function of the ratio of the bicarbonate and CO2. It is a logarithmic correlation, so the higher the CO2 level or the lower the KH, the lower will be the pH. The risk of having low KH is that KH can be easier "consumed" on the logarithmic scale than higher KH. For example, a consumption of 0.9 dKH from 1, would decrease the bicarbonate to 10%, which results in a 1 pH drop, consumption of 0.99 results in a 2 pH drop. At the starting KH of the 2, the pH drop is much less, about 0.2 in both cases.

I believe that in heavily planted tanks with low bioload, most of the produced ammonia could be uptaken by plants and not much nitrification would be required. I guess that most of the stable tanks with near-zero KH are these types of aquariums.

 

[dw1305]

Hi all,

Theoretically, there will be always bicarbonate once CO2 is dissolved in the water. If no other buffers are present, then the amount is negligible, but once there are other buffers, the bicarbonate levels can go significantly up even in a supposed-to-be 0 dKH water. Significant here means compared to 0 dKH.

I think you are right, you would need @Andy Pierce for confirmation, unfortunately inorganic chemistry <"isn't my strong point">.

If I am not mistaken, nitrification does not consume bicarbonate directly, but bicarbonate buffers the produced H+ ions during the nitrification and is converted to water and CO2, which eventually dissipates. It is the bicarbonate that buffers the H+ ions because it is the most significant buffer in the aquarium, but theoretically, it might be something else as well.

I'm pretty sure that is right as well <"https://www.sciencedirect.com/science/article/pii/S0964830517312933">.
Leon Steuernagel, Erika Lizette de Leon Gallegos, Asma Azizan, Ann-Kathrin Dampmann, Mohammad Azari, Martin Denecke, (2018)
"Availability of carbon sources on the ratio of nitrifying microbial biomass in an industrial activated sludge", International Biodeterioration & Biodegradation, 129, pp 133-140.

.....The major fraction of nitrifying bacteria in activated sludge plants are autotrophs that require an inorganic carbon (IC) source for biosynthesis (Bitton, 2005). Nitrifiers assimilate one-carbon compounds such as CO2 via the Calvin-Benson-Bassham-Cycle, catalyzed by the enzyme Ribulose-1,5-bisphosphate carboxylase/oxygenase or RubisCO (Schramm et al., 1998, Utåker et al., 2002). Hommes et al. (2003) showed that certain organic compounds are available as carbon source for N.europaea but inorganic carbon (IC) is the preferred substrate for biosynthesis. Jiang et al. (2015) stated that gaseous CO2 is a less preferable source of inorganic carbon compared to bicarbonate. ........

I'm guessing that in more oligotrophic and acidic conditions there are a different suite of micro-organisms that preferentially utilise CO2 or have a much higher affinity for HCO3- ions etc. It is back to the argument in <"Test reading not what i expected ?">.

The other candidates are some organic acids released from decaying plants and soil.

I think that is likely.

I believe that in heavily planted tanks with low bioload, most of the produced ammonia could be uptaken by plants and not much nitrification would be required. I guess that most of the stable tanks with near-zero KH are these types of aquariums.

I think that is likely as well.

cheers Darrel

 

[fredi]

With the test kits available to most aquarists, we are measuring DKh
From memory 1DKh = 17.3ppm (please don’t shoot me if I am out a bit 😂😂)
Lod (limit of detection) also applies, just because we cannot “see” it, doesn’t mean it’s not there

 

[dw1305]

Hi all,

With the test kits available to most aquarists, we are measuring DKh

They are actually measuring alkalinity (via titrimetric methods), but normally it is a fair supposition that the alkalinity is from dissolved bicarbonates.

From memory 1DKh = 17.3ppm (please don’t shoot me if I am out a bit 😂😂)

It is a really strange derivation. The details are in <"CO2 gaseous equilibrium with atmosphere"> and refer back to Larry Frank's article at the Krib <"Water Hardness">.

1dKH = 17.86 ppm CaCO3
From above; 1dKH = 17.8575 mg/liter CaCO3. 7.143 mg/liter of this is Ca, the rest ;(17.8575-7.143)= 10.7145mg/liter CO3
1dKH = 10.7145 ppm CO3

Cheers Darrel

 

[fredi]

It is a really strange derivation. The details are in "" and refer back to Larry Frank's article at the Krib.

I learn something new every day

 

[aquanoobie]

I believe that in heavily planted tanks with low bioload, most of the produced ammonia could be uptaken by plants and not much nitrification would be required. I guess that most of the stable tanks with near-zero KH are these types of aquariums.

+1

That is why I've tended to regard hardness (dGH) and alkalinity (dKH) as the two halves of the same whole.

SO4 compounds will change that.

pH is a function of the ratio of the bicarbonate and CO2. It is a logarithmic correlation, so the higher the CO2 level or the lower the KH, the lower will be the pH.

Only up to a point since carbonic acid H2CO3 created by CO2 injection is a mild acid.

 

[bazz]

Hi @dw1305 Darrel,
Does this work the other way round for GH?
1dKH = 17.86 ppm CaCO3
From above; 1dKH = 17.8575 mg/liter CaCO3. 7.143 mg/liter of this is Ca, the rest ;(17.8575-7.143)= 10.7145mg/liter CO3
1dKH = 10.7145 ppm CO3
Cheers!

 

[aquanoobie]

17.86 ppm CaCO3 = 1 dKH
And if 1 dKH is made of other compounds like baking soda NaHCO3 for example, then we can say the 1 dKH is an equivalent of 17.86 ppm of CaCO3.

 

[Andy Pierce]

I am unsure if there is such a thing as 0 dKH alkalinity in aquariums. Theoretically, there will be always bicarbonate once CO2 is dissolved in the water. If no other buffers are present, then the amount is negligible, but once there are other buffers, the bicarbonate levels can go significantly up even in a supposed-to-be 0 dKH water. Significant here means compared to 0 dKH.

For practical purposes in a real world aquarium at any reasonable pH there will always be some amount of bicarbonate in the water. Strictly however, the amount of bicarbonate is pH dependent so at pH 5 there isn't very much bicarbonate and at pH 4 there is almost none. Since the pH of pure distilled water exposed to air is around 5.8, there will generally be some bicarbonate present. To achieve a lower pH you'd need to add a stronger acid than composite carbonic acid. I don't think the presence of "other buffers" is relevant to the amount of bicarbonate, other than indirectly in as much as those other buffers may affect the pH which in turn does affect the equilibrium levels of bicarbonate.

Does this work the other way round for GH?
1dKH = 17.86 ppm CaCO3
From above; 1dKH = 17.8575 mg/liter CaCO3. 7.143 mg/liter of this is Ca, the rest ;(17.8575-7.143)= 10.7145mg/liter CO3
1dKH = 10.7145 ppm CO3

I think it's best not to confuse dGH and dKH and it's really unfortunate that these are both quoted in CaCO3 equivalents. dGH is the amount of divalent cations (Mg++, Ca++) and dKH is (usually) titratable alkalinity of which most will be carbonate-species based. They aren't inherently linked at all.

 

[hax47]

I don't think the presence of "other buffers" is relevant to the amount of bicarbonate, other than indirectly in as much as those other buffers may affect the pH which in turn does affect the equilibrium levels of bicarbonate.

This is the reaction:

CO2 + H2O <- ... -> H+ + HCO3-

The way I like to imagine this is that without any other buffers in RO water, this reaction stops at very low bicarbonate concentrations, equal to that of the produced H+. If there is another buffer, it binds the produced protons and shifts the reaction to the right side, so more bicarbonate is created. With reasonable buffer capacity, a reasonable amount of bicarbonate will be produced until the equilibrium in the reaction is reached. In terms of alkalinity (measured KH) this would not make too much shift since one buffer binds the protons while an equal amount of bicarbonate is created. The consequence is that the pH does not change as much as expected from CO2 concentration changes, since the bicarbonate concentration changes together with CO2 (so 10x increase in CO2 would not lead to 1 pH change).
But I guess this is the same as what you say, other buffers affect the pH and therefore the equilibrium. With the same CO2 concentration, if the pH is different when other buffers are present, in equilibrium, the bicarbonate levels must be also different. We can't have the same bicarbonate with the same CO2 and different pH.
I am not sure if this has any relevance in aquariums, but if it has, it must be at very low, near-zero KH values where the amount of other buffers may be comparable to that of bicarbonate.
In aquariums, the one place where this is probably more important is the blood in the fish. When we dose CO2, not only the water pH is affected, but also the pH inside the fish because of higher CO2 levels (fish don't have too many options to limit the uptake of CO2). However, in fish, the blood contains a reasonable amount of proteins that act as buffers, therefore limiting the effect of the CO2 on the pH inside the fish. So 10x CO2 increase inside the fish would not lower the pH by 1 as is the case in the water.

 

[dw1305]

Hi all,

SO4 compounds will change that.

I'm not sure how much effect they will have. The sulphate ion (SO4--) is a base, <"but a weak one">.

Because it is such a weak base, sulfate ion undergoes negligible hydrolysis in aqueous solution.

I think that is saying that basically we can regard it as neutral? but we would need @Andy Pierce or @hax47 .

I think the same would apply to (ortho)phosphate (PO4---), it is a <"stronger base">, but I don't know how much effect they have in practice - <"Conjugate (acid-base theory) - Wikipedia">.

I think this is the base bit, <"PO4---"> can accept protons (H+ ions) and bases are "proton acceptors".

The phosphate or orthophosphate ion [PO4]−-- is derived from phosphoric acid by the removal of three protons H+ . Removal of one proton gives the dihydrogen phosphate ion [H2PO4]− while removal of two ions gives the hydrogen phosphate ion [HPO4]2−.

cheers Darrel

 

[dw1305]

Hi all,

Hi @dw1305 Darrel,
Does this work the other way round for GH?
1dKH = 17.86 ppm CaCO3
From above; 1dKH = 17.8575 mg/liter CaCO3. 7.143 mg/liter of this is Ca, the rest ;(17.8575-7.143)= 10.7145mg/liter CO3
1dKH = 10.7145 ppm CO3

It does, it really is the <"most horrible mess">. @_Maq_ discusses units etc. in <"Some handy facts about water">.

17.86 ppm CaCO3 = 1 dKH
And if 1 dKH is made of other compounds like baking soda NaHCO3 for example, then we can say the 1 dKH is an equivalent of 17.86 ppm of CaCO3.

I think it's best not to confuse dGH and dKH and it's really unfortunate that these are both quoted in CaCO3 equivalents.

Yes, both dGH and dKH are actually defined by their (entirely theoretical) calcium oxide (CaO) content, so 17.86 ppm CaCO3 is both 1 dGH and 1 dKH <"Water Hardness">.

dGH is the amount of divalent cations (Mg++, Ca++) and dKH is (usually) titratable alkalinity of which most will be carbonate-species based. They aren't inherently linked at all.

They aren't, but because both divalent cations and carbonate anion <"normally derive from CaCO3"> (a 1 : 1 ratio of dGH : dKH) practically they are often linked.

If you don't get your calcium (Ca++) ions from CaCO3? You can make designer water. @Roland for example <"has decoupled"> dGH and dKH in his <"Soft water tank">.

But I guess this is the same as what you say, other buffers affect the pH and therefore the equilibrium. With the same CO2 concentration, if the pH is different when other buffers are present, in equilibrium, the bicarbonate levels must be also different. We can't have the same bicarbonate with the same CO2 and different pH.

That one.

In aquariums, the one place where this is probably more important is the blood in the fish. When we dose CO2, not only the water pH is affected, but also the pH inside the fish because of higher CO2 levels (fish don't have too many options to limit the uptake of CO2). However, in fish, the blood contains a reasonable amount of proteins that act as buffers, therefore limiting the effect of the CO2 on the pH inside the fish. So 10x CO2 increase inside the fish would not lower the pH by 1 as is the case in the water.

We have some discussion of the Bohr (and Root) effects, and how they <"may differentially effect different fish"> when you add CO2. Personally I'm not a CO2 user, so I don't have any practical experience of the effect on different fish of elevated CO2 levels.

cheers Darrel

 

[Andy Pierce]

I'm not sure how much effect they will have. The sulphate ion (SO4--) is a base, <"but a weak one">.
I think that is saying that basically we can regard it as neutral? but we would need @Andy Pierce or @hax47 .
I think the same would apply to (ortho)phosphate (PO4---), it is a <"stronger base">, but I don't know how much effect they have in practice - <"Conjugate (acid-base theory) - Wikipedia">.
I think this is the base bit, <"PO4---"> can accept protons (H+ ions) and bases are "proton acceptors".

You can read these off from a table of pKa values.

The acid version of the sulphate ion is sulfuric acid H2SO4. There are two protons that can be lost. The first is lost essentially all the time (that's why sulfuric acid is such a strong acid). The second is half-on and half-off at a pH of 1.99 (pKa2). You can math it up, but essentially that means at pH 3 or higher there is no capacity for the sulphate ion to accept a proton, which means at the type of aquarium pH we are considering the sulphate ion does not affect pH at all.

The acid version of the phosphate ion is phosphoric acid H3PO4. There are three protons that can be lost. The pH values at which each of these protons are half-on/half-off are 2.16, 7.21 and 12.32 (pKa1, pKa2 and pKa3 respectively). That means at our usual pH range, the first proton is always lost (pKa 2.16) , the third proton is always still on the phosphate (pKa 12.32), but the second proton can come on and off, so you'll have a mix of HPO3-- and H2PO3-. You get the most "buffering capacity" at the pKa, so phosphate will be an excellent buffer around pH 7 (best at pH 7.21). If you were to add a purely phosphate salt like K3PO4 to aquarium water, the K3 will all immediately come off, and the loose PO3--- will immediately grab (on average) 1.5 protons which means the pH will go up because the concentration of free H+ will go down (some got stuck on the phosphate). People usually don't want this to happen, which is why they use "h̶y̶d̶r̶a̶t̶e̶d̶" "mono- and/or dibasic" versions of potassium phosphate, typically KH2PO4.

For funzies we can look at (composite) carbonic acid H2CO3, with pKa1 of 6.35 and pKa2 of 10.33. That means the first proton can come on and off around pH 6-7, whilst the second proton is essentially always on below pH of 9.5, so there will be a mix of H2CO3 and HCO3- at neutral(ish) pH.