"Part of the carbonate hardness is not available. You add 2kh mathematically, but no 2kh is created. Microcrystalline precipitates are formed, which then form small fractions that are mainly deposited on the surfaces. This happens within several hours and then they go back into re-dissolution over several hours due to the carbonic acid and CO2 activity that the material brings with it.
Very roughly: you then only measure 1kh and over time the rest goes into the re-dissolution phase and provides the alkanity for the next few hours."
Thanks for the translations.
That is not any sort of viable explanation for why it is not present in a beaker test of the alk in the product. An alkalinity titration takes the pH far lower (pH in the 4's) than is needed to dissolve any tiny calcium carbonate particles, so if they formed they would be counted and is also far lower than is present in the tank at any time of day.
It is certainly possible that such solids form, and they are more likely to form with higher pH additives. They seem to want to claim it is a benefit when using bicarbonate and a detriment that breaks the buffer system when using high pH additives such as carbonate or hydroxide. Which is it?
FWIW, there's no evidence that the pH in tank gets low enough to dissolve calcium carbonate particles. There are local situations where it may happen (deep in sand beds, inside sea cucumber GI tracts, under bacterial biofilms, etc, but calcium carbonate particles simply coating surfaces such as live rock are exposed to water near pH 8, where they are not going to readily dissolve.
About the precipitations of carbonates and hydroxides:
"The problem with precipitation is, and this is often argued, that only magnesium hydroxide is formed. That is one of the typical myths we have.
And it has now been proven by Armin Glaser (German chemist) and countless studies (University of Hawaii etc.) that this is not the case.
We have a production of various carbonates, but these hydroxides, which dissolve again afterwards, also pull organic substances, amino acids and trace elements close to them. So the adsorption capacity is even greater than with carbonates and they often remain in the deposits, even if some of them dissolve again. What is forgotten is the adsorption capacity of these ultra-fine materials."
I'm not going to defend what anyone else has said, but I have said hundreds of times that calcium carbonate may also be formed when dosing high pH additives, which is why you want to mix in high pH additives quickly.
I have also said hundreds of times that calcium carbonate surfaces rapidly bind many things, including organics. That statement is literally part of every repost of my directions of how to stop excessive precipitation of calcium carbonate. I don't know if their assertion about hydroxides remaining intact and not dissolving over time is true or not, but the assertion of their surfaces remaining hydroxide is likely incorrect and obviously is without any supporting data. Exposed hydroxides are likely to rapidly react to form carbonate. Regardless, there is no doubt that precipitated minerals bind things. That is why many people dose calcium carbonate particles (coral snow) to begin with.
"The buffer system is related to the fact that we have various precipitation mechanisms, not only for the precipitation of chemicals during dosing, which is one thing. Other precipitations also occur on organic surfaces or on ceramic/calcareous surfaces, which practically cling to the surfaces like micro crystals.
I certainly agree that is possible in many scenarios. I discuss that with reefers all the time (excessive precipitation scenarios).
They then dissolve again due to the storage of CO2 or carbonic acid.
They may redissolve for various reasons where the local pH may be lower than in the tank water (mentioned above), but I have no idea what he is talking about with stored CO2 or carbonic acid. CO2 and carbonic acid are ONLY a function of tank water ph. There's no hidden CO2 waiting to spring out and do something, unless one is talking about metabolism or organics that produce Co2, such as when organic carbon dosing. Bicarbonate bolus dosing certainly does not store up hidden CO2 or carbonic acid that is not exactly reflected by the pH.
It is a combination of the addition and the saturation point of the lime. Seawater is not particularly stable. If we increase the pH value or change the carbonate content or the calcium content compared to the balanced average, which is maintained by magnesium, we automatically produce precipitates. No matter what we dose, this has nothing to do with bicarbonate.
Seawater at any pH above 7.7 and normal alk (or higher) is supersaturated with calcium carbonate, and it can precipitate. The higher the alkalinity and the higher the pH, the more prone the seawater is to have precipitation. A number of things, not just magnesium, but also including organics and phosphate, serve to slow or stop that precipitation. Certainly in the local area where alkalinity additives (such as bicarbonate) are added, this can happen from high alk and with dosing of high alk and pH supplements such as carbonate and hydroxide, it is even more likely. That, again, is why it is important to mix in rapidly, and I'd never claim it does not happen even then. It happens, and is no big deal. I said this decades ago. Abiotic precipitation happens all the time.
cs.lbl.gov
