Potassium carbonate or bicarbonate for mash pH adjustment

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Kaiser

Kaiser

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A recent comment on my blog asks why we brewers are not using potassium carbonate or bicarbonate to raise mash pH? Wine and mead makers seem to use it.

I didn't have a good answer for this. I guess it is because potassium is not a major water mineral and by adding it to the water we wouldn't be building realistic waters. However, malt brings in about 500 mg/l potassium (according to Narziss) and a bit more from potassium salts shouldn't matter. Thus mash pH adjustment with potassium carbonate or the not so caustic potassium bicarbonate should be a viable alternative to calcium carbonate since it does dissolve more readily in water. This assumes that you already getting enough calcium from the water or other salts.

Any thoughts?
 
I've seen potassium carbonate in a recipe from Papazian. I think it was Toad Spit Stout. High ph leads to either a stuck sparge or tannin extraction. I can't remember which. Since you heard this from a wine guy, I would guess it's about the tannins. Tannins can be desireable in wine.
 
Kaiser said:
A recent comment on my blog asks why we brewers are not using potassium carbonate or bicarbonate to raise mash pH? Wine and mead makers seem to use it.

I didn't have a good answer for this. I guess it is because potassium is not a major water mineral and by adding it to the water we wouldn't be building realistic waters. However, malt brings in about 500 mg/l potassium (according to Narziss) and a bit more from potassium salts shouldn't matter. Thus mash pH adjustment with potassium carbonate or the not so caustic potassium bicarbonate should be a viable alternative to calcium carbonate since it does dissolve more readily in water. This assumes that you already getting enough calcium from the water or other salts.

Any thoughts?
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I have used potassium hydroxide for mash adjustment and it worked as expected. I did this because I keep a working solution to use in mead so it is handy (why handle straight calcium hydroxide or make a separate working solution?).

The reasoning in mead making is that potassium is good for yeast and unlike malt honey is not rich in it.

If I were not making mead I would use calcium hydroxide as additional calcium is more likely to benefit beer than potassium, and it is easier to purchase.
 
Good question. Might it be because brewers rarely need to *raise* pH, so there's not been a lot of thought put into the issue? I guess the notable exception is malt bills with a lot of roasted grain, but there's already a long tradition of calcium carbonate there. Seems like if there's no particular reason to increase K+ over Ca++, the default would be Ca++ for all its varied benefits.
 
dwarven_stout said:
Good question. Might it be because brewers rarely need to *raise* pH, so there's not been a lot of thought put into the issue? I guess the notable exception is malt bills with a lot of roasted grain, but there's already a long tradition of calcium carbonate there. Seems like if there's no particular reason to increase K+ over Ca++, the default would be Ca++ for all its varied benefits.
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I think it is also the fascination with replicating naturally occurring water from particular cities (I can see the aesthetic appeal of using naturally occurring water but synthesizing it rather than simply synthesizing the optimal water seems like folly).

Potassium carbonate is much more soluble in water than the calcium counterpart which is reason enough to use it if you want carbonate (I don't and am confident handling hydroxide).
 
remilard said:
I think it is also the fascination with replicating naturally occurring water from particular cities (I can see the aesthetic appeal of using naturally occurring water but synthesizing it rather than simply synthesizing the optimal water seems like folly).

Potassium carbonate is much more soluble in water than the calcium counterpart which is reason enough to use it if you want carbonate (I don't and am confident handling hydroxide).
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Yeah. Witness people dropping in 700ppm sulfate because that's what Burton-on-Trent has.
 
Kaiser said:
A recent comment on my blog asks why we brewers are not using potassium carbonate or bicarbonate to raise mash pH? Wine and mead makers seem to use it.
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There are two reasons, I think, that wine and mead makers use them vs. CaCO3. Most importantly they are more soluable, so they can be added to the must directly. This is less of an issue for beer where the salts are added to the mash. Second, for commercial producers, TTB has limits on CaCO3 quantities you can add, whereas potassium (bi)carbonate are generally recognized as safe and have no limits. My water is slightly calcium deficient, so I need to add some calcium salts to get above the 50ppm for mead, and I do pH adjustment with K-bicarb.

I've never thought about putting anything but CaCO3 in the mash, because of the benefits of calcium for the yeast...
 
Considering using potassium carbonate myself. I saw another post on water profile for a RIS, and he said he was having trouble getting a high RA without overdoing the calcium or sodium. Therefore, I figured either using potassium carbonate or bicarb would be a good alternative.
 
RA is not a target. Only pH is a target and you adjust RA only to the degree necessary to achieve an acceptable mash pH. I note that that thread is from 2009 and the understanding of brewing water chemistry has changed markedly.

As to using potassium compounds for brewing water adjustment. I am not a fan. Sure, malt adds a lot of potassium to wort. However, that potassium is associated with organic molecules and ligands that apparently keep it from contributing potassium's somewhat salty taste to beer. Adding potassium carbonate would not have that luxury and there is a potential for creating a salty flavor in beer.

With all that said, we don't typically have to add much alkalinity to mashing water. So its possible that using a small amount of potassium carbonate would not incur much flavor penalty. Give it a go. Report back. Remember, don't aim for a particular RA. Aim for a mash pH.
 
Potassium ions are going to target the same salt receptors as any other potassium ions. They are very ionic - they don't covalently bond to other atoms very well at all. Therefore, your system will detect it the same way.

I understand on the RA thing. I am used to talking chemistry in organic chemist terms, not in typical beer chemistry lingo, so final mash pH adjustment is what I mean, excuse me.

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Potassium carbonate is going to have over twice the effect on pH as sodium bicarbonate. First of all, it pulls hydrogen ions much more than bicarbonate does. It doesn't really sit in equilibrium between carbonate and bicarb like bicarb does with carbonic acid. Also, when it becomes bicarb, it partially absorbs more hydrogen ions.

Then there is the fact that potassium isn't acidic, as calcium is. Therefore, it won't negate some of the alkalinity of the carbonate like calcium does.

Those things being said, to an extent, calcium ions taste good in water. It takes much less potassium to reach an undesirable level. And each potassium carbonate gives two potassium ions. But it should take less to adjust pH with it than with calcium carbonate and way less than sodium bicarb. Honestly, I would think potassium carbonate would always be uses over sodium bicarb.

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OhSoHumuLonely said:
Potassium carbonate is going to have over twice the effect on pH as sodium bicarbonate. First of all, it pulls hydrogen ions much more than bicarbonate does.
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On a molar basis. That's because it is 1) dibasic (and bicarbonate is monobasic) and the first pKb of bicarb is much lower than the first (and only) pKb for bicarbonate. At any reasonable mash pH each mmol of dibasic carbonate absorbs an extra mmol of proton in converting to monobasic bicarbonate which then behaves just like the bicarbonate from sodium bicarbonate (or potassium bicarbonate). A millmole of bicarbonate will absorb 0.9 mmol of hydrogen ions to pH 5.4 (typical mash pH) and 0.7 to pH 6. A millimole of carbonate respectively 1.9 and 1.7 i.e. 1 extra in each case. But all this is on a molar basis. On an equivalence basis there isn't that much difference at all. See the example at the end.



OhSoHumuLonely said:
It doesn't really sit in equilibrium between carbonate and bicarb like bicarb does with carbonic acid. Also, when it becomes bicarb, it partially absorbs more hydrogen ions.
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The equilibrium equations are the same for any carbo (carbonic, bicarbonate and carbonate) system i.e. whether the carbo comes from dissolved CO2, dissolved limestone, added bicarbonate or added carbonate (of any metal). The equilibrium is
H2CO3 <---> HCO3- <---> CO3--
with pK1 over the first arrow and pK2 over the second. If I know the equilibrium carbonic acid concentration and the pH I know the equilibrium carbonate concentration (and the equilibrium bicarbonate concentration).

OhSoHumuLonely said:
Then there is the fact that potassium isn't acidic, as calcium is. Therefore, it won't negate some of the alkalinity of the carbonate like calcium does.
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Neither is acidic. The difference is that CaCO3 is much less soluble than K2CO3 so that if calcium is introduced into a solution containing carbonate a third equilibrium must be considered.

Ca++ + H2CO3 <---> Ca++ + HCO3- <---> Ca++ + CO3-- <--> CaCO3

If there is enough calcium to exceed the solubility product CaCO3 will precipitate upsetting the equilubrium and carbo will flow to the right releasing protons as carbonic converts to bicarbonate and bicarbonate to carbonic. This continues until equilibrium is restored. This is true whether the carbonate came from limestone or potassium carbonate. Add calcium to a solution of potassium carbonate and precipitation will occur with a decrease in pH. It looks as if acid has been added but it hasn't.

OhSoHumuLonely said:
Those things being said, to an extent, calcium ions taste good in water.
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Well, it's flavor neutral for the most part.

OhSoHumuLonely said:
But it should take less to adjust pH with it than with calcium carbonate and way less than sodium bicarb. Honestly, I would think potassium carbonate would always be uses over sodium bicarb.
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A milliequivalent of sodium bicarbonate is as effective as a milliequivalent of potassium carbonate and each is associated with a milliequivalent of a metal ion.

For example, suppose I propose to do a beer with a proton surfeit of 100 mEq at pH 5.4 so that I need 100 mEq of base. I can get that from 111 mmol of sodium bicarbonate (9.3g) or 100 mmol of NaOH (4 g), 100 mmol of KOH (5.6 g), 52.6 mmol of Na2CO3 (6.5g) or 52.6 mmol K2CO3 (6.5 g) or 50 mmol of Ca(OH)2 (3.7g). If I choose a hydroxide base I wind up with 100 mEq of cation. If I choose a bicarbonate base I get 111. If I choose a carbonate base I get 105.2. That's an 11% difference between hydroxide and bicarbonate and 6% difference between bicarbonate and carbonate which isn't enough to get excited about IMO. Also note that of the carbo added, from whichever source is chosen, 90% leaves the solution (at pH 5.4) as CO2 and if a hydroxide base is chosen all the hydroxyl ions turn into water. In the last analysis all that's left is the 100 - 111 mEq of the metal ion.


The problem with carbonate is that the alkali earth salts are hard to dissolve and it reacts slowly so that one doesn't get the total expected effect until well into or even after the completion of the mash.
 
ajdelange said:
On a molar basis. That's because it is 1) dibasic (and bicarbonate is monobasic) and the first pKb of bicarb is much lower than the first (and only) pKb for bicarbonate. At any reasonable mash pH each mmol of dibasic carbonate absorbs an extra mmol of proton in converting to monobasic bicarbonate which then behaves just like the bicarbonate from sodium bicarbonate (or potassium bicarbonate). A millmole of bicarbonate will absorb 0.9 mmol of hydrogen ions to pH 5.4 (typical mash pH) and 0.7 to pH 6. A millimole of carbonate respectively 1.9 and 1.7 i.e. 1 extra in each case. But all this is on a molar basis. On an equivalence basis there isn't that much difference at all. See the example at the end.



The equilibrium equations are the same for any carbo (carbonic, bicarbonate and carbonate) system i.e. whether the carbo comes from dissolved CO2, dissolved limestone, added bicarbonate or added carbonate (of any metal). The equilibrium is
H2CO3 <---> HCO3- <---> CO3--
with pK1 over the first arrow and pK2 over the second. If I know the equilibrium carbonic acid concentration and the pH I know the equilibrium carbonate concentration (and the equilibrium bicarbonate concentration).


Neither is acidic. The difference is that CaCO3 is much less soluble than K2CO3 so that if calcium is introduced into a solution containing carbonate a third equilibrium must be considered.

Ca++ + H2CO3 <---> Ca++ + HCO3- <---> Ca++ + CO3-- <--> CaCO3

If there is enough calcium to exceed the solubility product CaCO3 will precipitate upsetting the equilubrium and carbo will flow to the right releasing protons as carbonic converts to bicarbonate and bicarbonate to carbonic. This continues until equilibrium is restored. This is true whether the carbonate came from limestone or potassium carbonate. Add calcium to a solution of potassium carbonate and precipitation will occur with a decrease in pH. It looks as if acid has been added but it hasn't.

Well, it's flavor neutral for the most part.



A milliequivalent of sodium bicarbonate is as effective as a milliequivalent of potassium carbonate and each is associated with a milliequivalent of a metal ion.

For example, suppose I propose to do a beer with a proton surfeit of 100 mEq at pH 5.4 so that I need 100 mEq of base. I can get that from 111 mmol of sodium bicarbonate (9.3g) or 100 mmol of NaOH (4 g), 100 mmol of KOH (5.6 g), 52.6 mmol of Na2CO3 (6.5g) or 52.6 mmol K2CO3 (6.5 g) or 50 mmol of Ca(OH)2 (3.7g). If I choose a hydroxide base I wind up with 100 mEq of cation. If I choose a bicarbonate base I get 111. If I choose a carbonate base I get 105.2. That's an 11% difference between hydroxide and bicarbonate and 6% difference between bicarbonate and carbonate which isn't enough to get excited about IMO. Also note that of the carbo added, from whichever source is chosen, 90% leaves the solution (at pH 5.4) as CO2 and if a hydroxide base is chosen all the hydroxyl ions turn into water. In the last analysis all that's left is the 100 - 111 mEq of the metal ion.


The problem with carbonate is that the alkali earth salts are hard to dissolve and it reacts slowly so that one doesn't get the total expected effect until well into or even after the completion of the mash.

We could sumarize everything here by saying that if you need to knock out a mEq of protons you are going to pick up approximately a mEq of cation assuming that you use the usual bases (di and tribasic phosphates, hydroxides, mono and di basic carbonates) no matter which one you choose.
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Calcium is an acidic cation. It reacts with water to make calcium hydoxide (which is fairly covalent, so not very basic like free hydroxide) and 2H+. That is why adding calcium chloride or sulfate lowers pH.

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pKb of Ca(OH)2 is 2.37, while that of NaOH or KOH (free hydroxide) are 0.2 and 0.5 respectively.

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OhSoHumuLonely said:
Calcium is an acidic cation. It reacts with water to make calcium hydoxide (which is fairly covalent, so not very basic like free hydroxide) and 2H+.
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I think you are confusing the ion and the metal. If you put the metal in the water you get

Ca + 2H2O --> H2 + Ca++ + 2OH-

Looks to me as if the water has been reduced by the calcium. The pH of the water goes way up if you do this. That's why the family to which calcium belongs are called the 'alkali earths'. And how does this differ from what happens if you put an alkali metal (K, Na...) into water? If Ca is an acid by this reasoning so must K and Na be.

Ca++ + H2O ---> Ca++ + H2O AFAIK. At least that's what it does when I put a calcium salt into solution. For Ca++ to be an acid it would have to be able to protonate something or accept an electron pair from something. There are hydroxyl ions for sure in water but calcium cannot protonate them or accept electrons from them unless the concentrations of calcium and hydroxide go so high that the solubility of Ca(OH)2 is exceeded. One does not prepare calcium hydroxide by adding a calcium salt to water. He does it by adding a calcium salt and then enough hydroxide (from lye, for example) that the solubility product of Ca(OH)2 is exceeded and that compound precipitates. I'm guessing, given the electronegativities of calcium and oxygen, that the bonds would be for the most part ionic but given even ionic bonds have some covalent character I suppose you could argue that calcium is a Lewis acid. And that potassium, being even less electronegative would form a more ionic bond and thus be less acidic. Is this what you are talking about? This would be a stretch to my way of thinking, especially in the current context where the Lowry-Bronsted definition reigns as we are interested in solutions well below saturation levels except when talking about the mechanism by which calcium acidifies mash where it is phosphate that comes into play (and to a lesser extent carbonate).


OhSoHumuLonely said:
That is why adding calcium chloride or sulfate lowers pH.
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Except that it doesn't unless phosphate (or carbonate) is present in which case it is the precipitation of apatite (this is the reaction by which calcium acidifies mash) that releases protons by the mechanism I sketched for carbonate in the previous post. Adding either of these salts to pure water does not change the pH (unless the calcium chloride contains come calcium hydroxide which it often does). I guess I have never thought of the precipitation of apatite as the neutralization of a Lewis acid by a Lewis base but if the bonds have any covalent character I guess it is. For that matter, then, any reaction is an acid/base reaction (and I believe that what Lewis's central hypothesis was). To solve problems such as the current one the Lowry-Brønsted model with solubility products considered clearly has its advantages.
 
What I am saying is calcium carbonate will not have the same effect on pH as potassium carbonate because calcium will complex hydroxide. Therefore, though the carbonate will absorb the same amount of H+, the calcium will replace some of it by binding water to form calcium hydroxide and 2H+.

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I am not confusing calcium ion with calcium metal. Calcium metal is extremely basic (donate electrons to form Ca2+, the definition of a lewis base). Calcium ion is somewhat acidic. Here:

Ca + 2H2O --> Ca(OH)2 + 2H2

Ca2+ + 2H2O --> Ca(OH)2 + 2H+

Very different reactions. First is Ca as lewis base. Second is Ca2+ as lewis acid.

How would CaCl2 lower pH if it wasn't acidic?!
 
Granted, Ca2+ is not as acidic as Mg2+ or Al3+, but I just taught my gen chem students this somewhere. There are two factors that cause it to be acidic. First of all, it will precipitate hydroxide ions, raising the levels of H+ ions. But even before it precipitates, it will complex hydroxide in solution. Again, not as much as some more acidic cations, but it definitely does it.
 
OhSoHumuLonely said:
Granted, Ca2+ is not as acidic as Mg2+ or Al3+, but I just taught my gen chem students this somewhere. There are two factors that cause it to be acidic. First of all, it will precipitate hydroxide ions, raising the levels of H+ ions. But even before it precipitates, it will complex hydroxide in solution. Again, not as much as some more acidic cations, but it definitely does it.
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If that were true the pH would drop when you add a calcium salt to pure water. It doesn't (unless, as I have noted, the calcium salt is contaminated with Ca(OH)2 which, as I also noted, it sometimes is in which case the pH goes up). Go into the lab and try it. If I added calcium chloride to water and Ca(OH)2 precipitated I would expect to see the precipitate. I don't. It doesn't complex hydroxide in solution. Well, OK, ion pairs form but they are short lived and come apart pretty quick. Mg++ can be removed from water by raising the pH to the level where [OH]- becomes high enough to exceed the solubility of Mg(OH)2. This happens at pH 11 or 12. The log equilibrium constant for Mg++ + 2H2O <--> Mg(OH)2 + 2H+ is -16.8 for
Ca++ + 2H2O <--> Ca(OH2) + 2H+ it is -22.8. So while it does happen it just does not happen to any appreciable extent except at perhaps pH 14 - 15 and above. That's a loooong way from mash pH.

OhSoHumuLonely said:
How would CaCl2 lower pH if it wasn't acidic?!
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I already explained that but perhaps I wasn't clear. In the first place it doesn't lower pH except when phosphate is present. I encourage you, for the sake of your students, to verify this in the lab. When phosphate is present in any amount even though the amount of tri-basic phosphate is very, very small some hydroxyl apatite will precipitate because it is very, very insoluble. At a fixed pH this upsets the equilibrium between tribasic phosphate and dibasic. So some dibasic yields up a proton and converts to tribasic. That in turn upsets the equilibrium between dibasic and monobasic phosphate and so some monobasic phosphate yields up a proton to convert to dibasic. If there is still enough calcium present (and, of course, initially there will be) to saturate WRT apatite more will precipitate and the process continues. Calcium is being removed from the solution and protons added to it (two of which, per molecule, actually do come from water) so the pH goes down lowering the tribasic phosphate concentration and the calcium concentration. The process persists until enough calcium has been removed and the pH is lowered enough that the water is no longer saturated. The overall equation, assuming that we start with 100% monobasic phosphate (which is reasonable at mash pH) is:

10Ca++ +6(HPO4)- + 2H2O ---> Ca10(PO4)6(OH)2 + 14H+

Thus 20 mEq of precipitated calcium produces 14 mEq of protons 2 of which come from water and the remainder of which come from the monobasic phosphate ion. Ca++ + H20 ---> Ca(OH)2 + 2H+ just isn't a factor here unless you want to consider the hydroxyl part of the apatite as being from that mechanism and I guess that's reasonable.
 
OhSoHumuLonely said:
I will do it and video tape it for you.
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This if for your benefit (and your students) not mine. I know what will happen. I just went and did a sanity check as I had to check the pH of my fermenting beer anyway. Having done that I can perhaps save you some time by pointing out that if you do the test with DI water then you have to know the pH of the DI water before you add the CaCl2 and that's tough to get because DI water is very high resistance. At the equivalent of 50 grams CaCl3/L I got 5.2 pH which compares to a calculated pH (based on PaCO2 of 0.0004 bar) for DI water of about 5.56. Thus, a more relevant test as the real question is 'Does CaCl2 acidify brewing water?' is to take a sample of your brewing water and see if a CaCl2 addition can overcome any of its alkalinity. My water runs alkalinity of about 1.5 mEq/L. I added CaCl2 at the equivalent rate of 4 grams per liter (modified by the fact that it is supposed to be the dihydrate but is probably more hydrated than that) so say 2 - 3. It is ACS grade stuff and so the Ca(OH)2 content should be very low. My observation: the pH went up slightly. So then I thought lets add some phosphate and see what happens. Not difficult since the 7.00 buffer used to calibrate the meter was right there. Shooting in a bit of that got the pH down under 7 pretty quickly. Thus, and I didn't really need convincing, I have verified both aspects of what we have been discussing i.e. that adding calcium in the absence of phosphate does not reduce the alkalinity of brewing water but that doing it in the presence of phosphate (malt) does.

As you desperately seem to want Ca++ to be appreciably acid I doubt this will have much influence but I have posted a picture of M(OH) solubilities (sorry about the lousy quality) from Stumm and Morgan. As you can see, even though the Ca curve is in the worst part of the picture, Ca(OH)2 is plenty soluble until pH's gets over 12. As the equilibrium constant data I posted suggest, the Ca++ curve lies 3 pH units higher than the Mg curve.

MOH Solu.jpg
 
"A molecular entity (and the corresponding chemical species ) that is an electron-pair acceptor and therefore able to react with a Lewis base to form a Lewis adduct, by sharing the electron pair furnished by the Lewis base." Surely you could have looked that up yourself. Just search 'IUPAC Gold Book'

I wonder if this might be at the heart of your confusion. Clearly a hydrogen ion (naked proton) is a Lewis acid and a Hydroxyl ion a Lewis base as the hydroxyl ion can donate any of its three unshared pairs to the hydrogen forming water. But if two hydroxyl ions are attracted to a calcium ion that's not really the same thing as donating a pair. It is partially sharing two pair assuming that the bond has some covalent character.
-H:O: Ca++ :O:H-

Don't know whether the definition covers this situation or not. I do remember reading somewhere that Lewis wanted to cast all reactions, including redox, as acid base reactions so well it might and if the key word is 'shared' then I think it probably does. There is no such thing, AFAIK, as a pure ionic bond.

Whether calcium ion is a Lewis acid or not is actually irrelevant here as it is the Lowry-Brønsted definition that governs it's behaviour here. Under this definition it is not an acid and it doesn't shift water pH unless (and I am getting a little weary of repeating this over and over again) phosphate, or some similar anion, is present.

Do you want me to say Ca++ is a Lewis Acid? OK - it is. Is that relevant at mash pH? No it isn't as you will verify when you do your experiment.
 
You do seem to be mostly right, though I think we are actually talking about an Arrhenius acid. A Bronsted-Lowry acid is an H+ donor, where an Arrhenius base increases the amount of H+ in solution either by donating it or by absorbing OH-. In water, calcium doesn't seem to do this until it reaches very high concentrations. If you add ENOUGH calcium chloride, it will eventually precipitate as calcium hydroxide, but in small amounts, that doesn't seem to happen.

That being said, in a mash, it is important to consider Lewis acidity. Calcium does complex phosphates and other organic anions (and possibly neutral ligands, as well, but I don't know that for sure). Therefore, it acts as a Lewis acid and lowers pH. When adding calcium carbonate, this has to be taken into account.

Back to the matter at hand, however, potassium carbonate would NOT do this. Potassium is not a Lewis acid. Therefore, it will raise pH because carbonate will bind H+, but it will not be counteracted by the counter ion binding basic phosphates and such.

I apologize for my confusion with the ability of calcium ion to create H+ in water, and I am sorry for getting snippy. I am pretty sure that you're right about that now, as I can read everything you said clearly since I am now on my computer rather than on my phone.
 
OhSoHumuLonely said:
You do seem to be mostly right, though I think we are actually talking about an Arrhenius acid. A Bronsted-Lowry acid is an H+ donor, where an Arrhenius base increases the amount of H+ in solution either by donating it or by absorbing OH-.
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An Arrhenius acid is something that increases H+ in aqueous solutions and an Arrhenius base is something that increases OH- in aqueous solution. A Lowry-Brønsted acid is a proton donor and a Lowry-Brønsted base is a proton acceptor regardless of solvent but as the solvent is water here that's not something that needs to be of concern. For what we do in brewing the Lowry-Brønsted definition is very convenient. It is actually quite easy to predict mash or water pH by tracking the 'proton condition' which is the balance between proton deficit and surfeit. If, for example, we add some sodium bicarbonate to pure water its pH will change. Some of the bicarbonate will act like a base and absorb protons, some will act like an acid and yield protons and the water will act like an acid and give up protons. The pH is the pH at which the sum of the protons given up by the basic bicarb and the water equals the number absorbed by the basic bicarbonate and is easily found by varying the pH. This would be difficult to do using the Arrhenius definition as the bicarbonate is simultaneously acid and base and doesn't directly add any OH-.

OhSoHumuLonely said:
In water, calcium doesn't seem to do this until it reaches very high concentrations. If you add ENOUGH calcium chloride, it will eventually precipitate as calcium hydroxide, but in small amounts, that doesn't seem to happen.
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The chart from Stumm and Morgan and the equilibrium constant tell slightly different stories but using the number log[Ca++] = 21 - 2pH. Thus at pH 7 the molar concentration of Ca++ would be 1E7 (if we could extend the line back that far which of course we can't). It is clear that at ph 7 and below this is not a factor.


OhSoHumuLonely said:
That being said, in a mash, it is important to consider Lewis acidity. Calcium does complex phosphates
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Yes and in that case the reaction of calcium with water does produce 2 out of every 14 protons released.


OhSoHumuLonely said:
...and other organic anions (and possibly neutral ligands, as well, but I don't know that for sure). Therefore, it acts as a Lewis acid and lowers pH. When adding calcium carbonate, this has to be taken into account.
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It does complex with proteins for sure but I don't know any of the details. But if you consider the complexation of calcium with carbonate while the calcium does act as a Lewis acid there is no hydrogen ion release:

Ca++ + CO3-- --> CaCO3

The hydrogen ions come from bicarbonate ions when this reaction is in play. It is the same as with the precipitation of apatite. Only the portion of the calcium that reacts with water produces protons. The rest come from Lowry-Brønsted dissociations. But yes, the protons released by whatever mechanism do have to be taken into account. I think I have come up with a way of explaining and simplifying the estimation of mash pH based on the proton condition. I have been saying that calcium is, of course, not an acid but that it acts like one and has a proton surfeit of 1/3.5 mEq/L per meq/L calcium hardness based on Kohlbach's finding.

OhSoHumuLonely said:
Back to the matter at hand, however, potassium carbonate would NOT do this. Potassium is not a Lewis acid. Therefore, it will raise pH because carbonate will bind H+, but it will not be counteracted by the counter ion binding basic phosphates and such.
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Why isn't potassium a Lewis acid? Will it not form a partially covalent bond with (OH)- of there is enough of it and enough (OH)-? Is it infinitely soluble? The answer is that it is indeed a Lewis acid but that whereas it is largely irrelevant that calcium is a Lewis acid it is utterly irrelevant that potassium is as the conditions (pH, concentration) under which it reacts with water to produce hydrogen ions are even more bizarre than for calcium.

OhSoHumuLonely said:
I apologize for my confusion with the ability of calcium ion to create H+ in water, and I am sorry for getting snippy.
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No need. The discussion has resulted in me gaining several valuable insights.
 
I revert to my previous point about Ca2+ being a Lewis acid. I understand that it may not be so acidic as to complex water and form two hydronium ions and calcium hydroxide. It is acidic enough, however, to complex phosphates and other anions at much lower concentrations than hydroxide at mash pH. These anions are basic and by removing them, they can no longer bind to H+. So while the carbonate of CaCO3 should absorb as many protons as the carbonate in K2CO3, potassium won't appreciably bind phosphate, and therefore, while the carbonate in calcium carbonate is counteracted by the reduction of other basic anions, the carbonate in potassium carbonate should not be.

I may have missed something in your posts. As I said, I was reading everything on my phone, and I tend to miss things that way. I will go back and double check, but unless I find something that changes my mind here, I still say it should take fewer carbonate ions to raise the pH the same amount using potassium carbonate than it does calcium carbonate.
 
OhSoHumuLonely said:
I revert to my previous point about Ca2+ being a Lewis acid. I understand that it may not be so acidic as to complex water and form two hydronium ions and calcium hydroxide.
Click to expand...

Calcium being a Lewis acid really doesn't have anything to do with it. Adding a Lewis acid to a solution does not necessarily result in a lowering of pH as would be the case with an Arrhenius or Brønsted-Lowry acid. In the example of a Lewis acid given in the Gold Book,

Me3B + :NH3 ---> Me3BNH3

Me3B is a Lewis Acid but in reacting with ammonia no protons are released and the pH isn't lowered. What is significant with calcium is that it form adducts with the anions of certain acids and in so doing releases the protons from those acids. It is these protons that are responsible for mash pH reduction. Among these acids are water, phosphoric acid and its monobasic and dibasic anions, carbonic acid and the bicarbonate ion, protonated side chains on proteins and who knows what else in a mash. The release of protons has nothing to do with the acidity or basicity of calcium but rather with the solubility of the adducts.

In

Ca++ + 2H2O = 2H+ + Ca(OH)2

we don't care a fig about what kind of acid Ca++ may be nor whether indeed it is an acid at all. All we care about is that it can pull the anions off the water leaving protons. The source of the protons is the water. Apparently chemists sometimes extend the definition of an acid to include substances that have this property in combination with water so that Ca++ + 2H2O is considered an acid in the Brønsted - Lowry sense. It 'dissociates' to yield protons. And there is an equilibrium constant for this but that equilibrium constant is not a dissociation constant in the usual sense of the word but related to the solubility product. If the solubility product of Ca++ and OH- is exceeded then the reaction above takes place and hydrogen ions are released. This happens when log[Ca++] > 22.8 - 2pH IOW it doesn't happen with beer.

A similar reaction takes place with bicarbonate ion

HCO3- + Ca++ --> H+ + CaCO3

and another similar one with monobasic phosphate

10Ca++ + 6HPO4- + 2H2O --> Ca10(PO4)6(OH)2 + 14 H+

We don't care about the reaction mechanism of these either. We care how soluble the adducts are. The least soluble will determine whether any protons are released or not. Apatite is the least soluble and so determines whether any protons will be available from this source. One can solve a system of 9 simultaneous non linear equations in an attempt to predict how much, if any, acid is released from this source depending of the available calcium and phosphate but that's a pretty tall order.



OhSoHumuLonely said:
It is acidic enough, however, to complex phosphates and other anions at much lower concentrations than hydroxide at mash pH.
Click to expand...
I don't think acidic is the term we want here. What is really at issue is the solubility of the adducts. The other point is that binding to the anion does not release protons. The acid must be present i.e.
Ca++ + CO3-- --> CaCO3 no proton
Ca++ + HCO3- --> H+ + CaCO3 proton.

OhSoHumuLonely said:
These anions are basic and by removing them, they can no longer bind to H+.
Click to expand...
But they always come from an acid. By taking them out you upset the equilibrium between the fully deprotonated anion and the one-less-deprotonated anion causing conversion with freeing of protons.

OhSoHumuLonely said:
So while the carbonate of CaCO3 should absorb as many protons as the carbonate in K2CO3, potassium won't appreciably bind phosphate, and therefore, while the carbonate in calcium carbonate is counteracted by the reduction of other basic anions, the carbonate in potassium carbonate should not be.
Click to expand...

The mechanism when a calcium or magnesium base is added is that the calcium precipitates as apatite releasing protons some of which immediately neutralize the base producing water. The extent to which this takes place depends on the pH, among other things. If carbonates are involved things get ever stranger because of the slow reactions (again depending on pH). I did some experiments for the Palmer/Kaminsky book and think I found that calcium carbonate and hydroxide were about 30% as effective as KOH on a per equivalent basis. There is discussion of all this in Chapter 6.


OhSoHumuLonely said:
I may have missed something in your posts. As I said, I was reading everything on my phone, and I tend to miss things that way. I will go back and double check, but unless I find something that changes my mind here, I still say it should take fewer carbonate ions to raise the pH the same amount using potassium carbonate than it does calcium carbonate.
Click to expand...

Let's say less hydroxyl ions from potassium hydroxide than hydroxyl ions from calcium hydroxide because of the apaptite effect. When carbonates get into the picture things are just too strange. We advise against the use of carbonate in any form but find lime (Ca(OH)2) to be fine as long as one is aware that its 'efficiency' is less than the stoichimetry might suggest.
 
I agree the idea of Lewis acidity is not that important as beer goes. In reality, I think it is the Lewis acidity that causes the complex to form between phosphate and calcium ions. None the less, you've gotten to the core of the discussion. Raising pH with calcium reagents is less efficient than raising them with potassium reagents on a per-ion basis.
 

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