alkalinity as CaCO3

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Kaiser

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Why does the addition of 1 g of CaCO3 to 1 liter of distilled water not cause an alkalinity of 1000 mg/l as CaCO3? If I plug this into a water spread sheet (like Palmer’s) I get an alkalinity of 490.

I always assumed that 1 mg/l alkaliniy as CaCO3 is the same level of alkalinity that 1 mg/l CaCO3 contributes to the water.

Kai
 
Why does the addition of 1 g of CaCO3 to 1 liter of distilled water not cause an alkalinity of 1000 mg/l as CaCO3? If I plug this into a water spread sheet (like Palmer’s) I get an alkalinity of 490.

I always assumed that 1 mg/l alkaliniy as CaCO3 is the same level of alkalinity that 1 mg/l CaCO3 contributes to the water.

Kai

Seems to me like 1 g of CaCO3 to 1 liter of water would produce 1000 mg/L as CaCO3. Could Palmer be taking into account the fact that calcium carbonate isn't very soluble in pure water? I believe the solubility is only 40-50 mg/L.
 
Palmer and others, don't take solubility into account as the salts will eventually dissolve in the mash.

1 g to 1 l is a little much. But the same problem exists with 10 mg to 1l which should give an alkalinity of 10 ppm CaCO3.

Kai
 
Why does the addition of 1 g of CaCO3 to 1 liter of distilled water not cause an alkalinity of 1000 mg/l as CaCO3? If I plug this into a water spread sheet (like Palmer’s) I get an alkalinity of 490.

I always assumed that 1 mg/l alkaliniy as CaCO3 is the same level of alkalinity that 1 mg/l CaCO3 contributes to the water.

Kai

Dear Kai,

It's been a little while since I was a TA for gen chem, but basically this is because of the various equilibria involving carbonate. For example, the 490 figure that you got from the spreadsheet will actually depend upon the pH of the water, and the pressure of CO2 in the surrounding atmosphere. This is because of the following equilibria:

CaCO3(s) <---> Ca(2+)(aq) + CO3(2-)(aq)

H2O + CO3(2-) <----> HCO3(-) + OH(-)

H2O + HCO3(-) <----> H2CO3 + OH(-)

H2CO3 <---> CO2 + H2O

So the second equilbrium, assuming a moderate pH, will lean almost completely to the right, so you'll really have bicarbonate in solution, and not as much carbonate. From here, the bicarbonate can be protonated and turned into carbonic acid, which will quickly dissociate into water and CO2. There is then an equilibrium between CO2 in solution and CO2 in the atmosphere. If the pressure of CO2 in the surroundings is high, it won't escape from the solution (though this is not a factor in practical brewing practice). Thus, as the pH is dropped, more CaCO3 can dissolve as the equilibrium is pushed towards bicarbonate and carbonic acid. At the same time, you're going to be losing some of the carbonate to CO2. What this means is that even if you fully dissolve 1g of CaCO3, you will not get the expected 1000 alkalinity, though you WILL achieve the expected 400 ppm of calcium if added 1g is added to a liter of water. If you really want to convince yourself, you can take the pKa's of the three acid base equilbrium, and calculate the residual alklanity based on neutral water.
 
mkade, I have been thinking about this as well, but none of these calculators seem to use the pH for determining the equilibrium of carbonate and bicarbonate. They all seem to work on a simple calculation of the amount of ions that are added. So I did some more calculating and found this:

CaCO3 + H2O -> Ca++ + HCO3- + HO+

Now if you determine the bicarbonate content that has been contributed you find that 1 ppm CaCO3 will yield 0.61 ppm HCO3-. CaCO3's molecular weight is 100 and HCO3's is 61.

Now there is a formula that converts alkalinity as HCO3- to alkalinity as CaCO3. This assumes that HCO3- has one alkalinity equivalent and that CaCO3 has 2 alkalinity equivalents. As a result you mulyiply [HCO3-] with 50 and divide by 61. The result is 0.5ppm alkalinity as CaCO3.

But there is a problem. The actual alkalinity contributed by CaCO3 was not one equivalent but 2. The formula that calculates the CHO3 concentration forgot about the HO- that was also created for each dissolved CaCO3 molecule. According to Wikipedia (and I tend to believe the science stuff there), alkalinity is defined as the sum of all the possible H+ receptors balanced by the H+ ions:

[A] = 2[CO3--] + [HCO3-] + [HO-] - [H+]

So the alkalinity as HCO3 for the 1ppm CaCO3 addition should have been 2*0.6 ppm HCO3. One for the HCO3- and one for the HO-. Then the alkalinity as CaCO3 is also 1 ppm.

But nobody else seems to see it like that. Either everybody copied from the same source or I'm on the wrong track here. The latter is quite possible.

mkade, the equations that you listed change the type of the ions in the water but they don't change the alkalinity of the water. I.e. its ability to resist a pH change by absorbing newly added H+ ions as these H+ ions can be absorbed by CO3--, HCO3- or HO-.

Kai
 
sorry to go off topic. This is why I love brewing. Otherwise dull and boring science stuff becomes of the utmost importance when faced with the greatest quality in a brew. I was never that interested in chemistry class. Although I can't really contribute, I will be following this thread closely!
 
I will not explain this too much because I have to finish working up a patient and frankly I am tired of using my brain today... :D

You are trying to relate alkalinity as a linear function of concentration, when in fact it is logarithmic.

Henderson-Hasselbalch equation goes something like pH= pKa + log ([A-]/[HA]) By adding CaCO3 not only are you losing some to CO2 as mentioned above (leChataliers principal if I am not mistaken...) but the hydrogen ion concentration is only being decreased (alkalizing) according to the disassociation constants and the balance achieved in your water.

I have the feeling I just rambled a lot without making any sense whatsoever, so I am going to go back to reading about protein losing enteropathies before I confuse myself or anyone else....

Edit: I don't know if it was a typo, but it looked like you you were using a monoprotic representation for carbonic acid (HCO3) when it is realy polyprotic (H2CO3)...
 
Is it necessary to add calcium carbonate to distilled water when one is using other minerals. For example, I use distilled and/or RO, and plan to use gypsum, epsom salts, and calcium chloride to arrive at a base mineral profile, then supplement with buffer 5.2 to make sure my pH is on target. In this instance, is CaCO3 a desirable, or even necessary addition?
 
Oh wow....and so many years ago when I was taking chemistry survey classes......balancing equations wondering when am I ever going to have to figure moles per solution.....does an example come up years later.

Since it has been so many years that I've studied molecules, I can't be of any help with the original question. Will be interested to see what the concensus is and also wondering why this question has been asked. Is it more just for reference at how the calculating tools do it, or is it that they're under question?
 
At this point I wonder if we wrongly assume that CaCO3 adds only one equivalent of alkalinity which is the case if we express alkalinity as HCO3- and only count the HCO3- ions but not the CO3-- and HO-. So while the alkalinity as CaCO3 that we get from the water analysis is correct and accounts for all of the alkalinity sources (CO3--, HCO3- and OH-) the assumption that 1 g of CaCO3 in 1 gal will result in 158 ppm carbonate or 158 ppm bicarbonate is incorrect if the bicarbonate content is interpreted as the alkalinity of the water.

I'm trying to find sources (other than home brewing) that show the alkalinity added by CaCO3. But so far I have come up short.

Kai
 
I'm trying to find sources (other than home brewing) that show the alkalinity added by CaCO3. But so far I have come up short.

Yep, I think you'll find the best answer with chemists....I've got a biochemistry friend who'd probably be really impressed if I start asking about alkalinity of water when adding CaCO3. Next time he's online, I'll try flagging him. Just tertiary googling makes me think it can get pretty complicated with the types of titration going on...since water supplies are different and different buffers going on, I don't think it's as cut and dry. Actually, that link that SpanishCastle listed says that it's levels of HCO3, H2CO3, and CO3 that add to alkalinity..getting back to deathweed's chime about carbonic acid (H2CO3)...are you sure it's just equating ions? Might have more luck with water quality discussions:

Acceptance Criteria for Carbonate Speciation
Alkalinity and Water Quality

As for the application with brewing....I guess that's why I'm more just hands on (knowing variables change and I don't want to think that much about it). I just measure Ph and keep plugging until it reaches my target Ph
 
some more food for thought:

I found a reference that "the alkalinity of the water is mostly represented as bicarbonate (HCO3-)". If that is true then there has to be a negligible amount of [CO3--] and [HO-] compared to [HCO3-].

Going back to the addition of CaCO3 there must be a strong preference for the following reaction:

CaCO3 + H20 + CO2 -> Ca++ + HCO3- + HCO3-

The CO2 is present in the water just b/c it is present in the atmosphere. But what is interesting is that the addition of 1 CaCO3 molecule results in the creation of 2 HCO3 molecules. So each mg of CaCO3 actually adds 1.2 mg of HCO3 to the brewing water and not the 0.6 that are assumed in Palmer's spreadsheet and BeerSmith. If the CO2 is not there, the 2nd bicarbonate is not formed and a OH- ion takes its place or that part of the CaCO3 is not dissolved in the first place.

Kai
 
I need to refresh my memory on this business, but it's amazing...I see the *exact* same discussions at work in boiler feedwater. And yes, I am a gigantic chemistry nerd.

For an industrial-type discussion about hardness, look at this url:
Chapter 07 Precipitation Softening

The application may be different but there is material there applicable to brewing (like the use of slaked lime (Ca(OH)2) to reduce hardness by precipitating calcium bicarbonate as calcium carbonate.

-

Hardness is typically expressed as equivalent to ppm (mg/L) of CaCO3. The solubility of CaCO3 in water at room temperature is (I believe) 40-50 ppm. If you have anything higher than that, there are other materials adding hardness.

Calcium sulfate or calcium chloride will add permanent (non-carbonate) hardness from the calcium. Those salts are more soluble than CaCO3, which is how you could get a hardness over 47 ppm.

Hardness is converted from ppm calcium (MW=40) to ppm CaCO3 (MW=100) by multiplying by 100/40 = 2.5. Normally the titration to determine hardness directly measures Ca and Mg by reaction with a chelant; this ignores the anions so CaSO4, CaCl2, CaCO3 would all behave the same.

-

Alkalinity can be measured on different standards...in the document I linked to they mention "P" and "M" alkalinity...or how much acid you need to lower the pH below ~9 (strong bases neutralized) and ~4 (strong and weak bases neutralized). For brewing purposes, the M alkalinity is important. You total up the amount of acid used to neutralize all this, see how much CaCO3 would be required to neutralize that much acid, and that's the alkalinity as ppm CaCO3. (or you could total the normality in norm/L, then multiply by 50 g/norm to get ppm CaCO3).

Adding carbonate, bicarbonate, phosphate, or diphosphate will increase alkalinity. Adding calcium salts will increase alkalinity. You could add sodium carbonate to add alkalinity *and* remove hardness (soda softening).
 
Let's not forget the original question:

How much alkalinity is added by the addition of 1 mg CaCO3 per liter?

I say it is 1.2 ppm HCO3- and 1 ppm as CaCO3, but all brewing water spreadsheets I have checked so far disagree with me.

I guess I have to shoot John Palmer an e-mail regarding this.

Kai
 
Assuming the water is not saturated in CaCO3, adding 1 mg/L will add 1 mg/L of both alkalinity and hardness. If the water *is* saturated, it will drop harmlessly to the bottom.

Adding 1 mg/L of CaCl2 (MW = 40 + 35.5 x 2 = 110 g/mol) will add 1 x 100/110 = .9 mg/L of hardness as CaCO3. It will add zero alkalinity because Cl- isn't included in the alkalinity calculations.

If the pH of your water is lower (like in a mash), the equilibrium should shift slightly. Depending how low it is, a little CO2 may be evolved and more of the calcium may dissolve, adding additional hardness. Palmer's How To Brew does mention that malt will affect alkalinity (and pH) but not hardness.
 
Let's not forget the original question:

How much alkalinity is added by the addition of 1 mg CaCO3 per liter?

I say it is 1.2 ppm HCO3- and 1 ppm as CaCO3, but all brewing water spreadsheets I have checked so far disagree with me.

I guess I have to shoot John Palmer an e-mail regarding this.

Kai

Dear Kai,

I thought about this some more, and had to look up the actual definition of alkalinity (it's really a silly term, as a chemist we usually use more logical terms such as molarity, pH, etc). Here is a simple definition:

"Alkalinity can be measured by titrating a sample with a strong acid until all the buffering capacity of the aforementioned ions above the pH of bicarbonate or carbonate is consumed. This point is functionally set to pH 4.5. At this point, all the bases of interest have been protonated to the zero level species, hence they no longer cause alkalinity."

I think this is important because it changes the way we should think about things. Now I've made the assumption that we're dealing with a buffered system in my following calculations, because at the end of the day, your mash will be a buffered system. So this may not be the answer you're looking for, or what Palmer and co. use when calculating water chemistry. Basically, your total alkalinity will be the total numbers of moles of 2*CO3(2-) + HCO3- + OH- multiplied by the molecular weight of CaCO3 (because alkalinity is expressed in units of ppm CaCO3) * 10^3). So I right off the bat ignored the equilibrium between carbonate and bicarbonate because the pKa of that reaction is about 10.5, meaning that at pH 6.5, over 99.99% is in the bicarbonate form. The interesting thing is that pKa of carbonic acid is 6.37, meaning that in a buffered system, the residual alkalinity will change significantly with pH. Then, I considered a buffered system since it's much easier to do calculations because we can deal with a constant pH. Then with simple algebra, you can calculate the concentration of both CO3(2-) (assumed to be zero because it is much less than HCO3-, and negligible in these calculations), HCO3-, and OH-. This sum is then converted from moles to the units of alkalinity. Unfortunately, I don't think I can attach a spreadsheet, though I'd be willing to email it if anyone is interested. I have copied and pasted the cells below, and they are ugly. I also have calculated alkalinity at several pH's for a system when adding 1 g CaCO3 per liter of water below, including contribution from OH- and from CO3-.

pH 7:
Total: 810 ppm CaCO3
CO3-: 810 ppm CaCO3
OH-: 0.01 ppm CaCO3

pH 8:
Total: 977.1 ppm CaCO3
CO3-: 977.2 ppm CaCO3
OH-: 0.1 ppm CaCO3

pH 5.2:
Total: 63 ppm CaCO3
CO3-: 63 ppm CaCO3
OH- : 0.0001 ppm CaCO3

I think that these numbers demonstrate two very interesting points. Firstly, for those using pH 5.2 stabilizer, the residual alkalinity in your water is highly neutralized. Another point of note is that even at pH 8, the total alkalinity is almost completely arising from bicarbonate, and not OH-. Anyhow, it seems to me that calculations of this sort are actually quite complex, and calculating the "alkalinity" contribution of CaCO3 is silly. Instead, we should focus on the concentrations of the individual ions. I apologize for the length of this post. Hope at least one of you enjoys.

-Matt



pKa H2CO3 6.37

Ka = [hco3-][h+]/[h2co3]
pka 6.37
Ka 4.2658E-07

pH 7
10^-pH 0.0000001
1 g / L
moles CaCO3 0.009991308
molarity 0.009991308
moles H2CO3 9.9913e-3 - X
moles H+ 10^pH
moles HCO3- X


product top X*10^pH
cross mult Ka*(9.991e-3)-Ka*X
cross mult 2 10^pH*X
solve x 10^pH*X + Ka*X = Ka*(9.991e-3)
simplify X(10^pH+Ka) = Ka * (9.991e-3)
X X = [Ka * 9.991e-3]/[10^-pH+Ka)

X 0.00809391
alkalinity CO3 moles 0.00809391
alkalinity HCO3- 810.0951586
moles OH 0.0000001
alkalinity OH 0.0100087
total alkalinity 810.1051673
 
As mkade and Scotty_g are ultimately saying - it is a pH issue. In distilled water CaC03 is practically insoluble. As the pH changes, so does the solubility. I must assume then that the disagreement in the calculations is probably the result of adjusting for differences in solubility with pH.

I use the precipitation softening procedure explained in the link Scotty_g provided to soften my water. I add the lime and then let the precipitate (CaCO3) that forms settle out overnight. I get a nice layer of white sludge at the bottom.
 
Matt, I think we are getting somewhere here.

The influence of the Ph makes this certainly more complicated but given that Palmer and other author&#8217;s spreadsheets don&#8217;t take water pH into account I wonder if it is really necessary to go into that detail.

So the definition of Alkalinity is based on the amount of acid it takes to &#8220;consume&#8221; all the buffering capacity. This is the idea of titration and the amount of acid needed is then expressed in equivalents of CaCO3 where each CaCO3 molecule is equivalent to needing 2 H+ molecules from the acid. This makes sense to me and I trust the alkalinity that is given in a water report. This is also the alkalinity that affects the mash pH.

Now what I need to figure out is how much alkalinity (i.e. buffering capacity) am I adding when I add 1ppm CaCO3 to the water/mash. What are the actual reactions that lie behind the alkalinity contributions that you listed. If the addition of CaCO3 does not contribute 2 HCO3- to the water, as it seems to be the case for all the pH values that you mentioned. What happens to the missing HCO3- ions? Do they gobble up H+:

HCO3- + H+ -> H20 + CO2

If that is the case the pH needs to increase as the concentration of H+ ions is lowered. Conversely if the pH doesn&#8217;t change though the addition of CaCO3, it needs to disassociate into Ca++ and CO3&#8212;or two HCO3- b/c if it was to create HO- the pH would change due to the need of the product between [HO-][H+] to remain constant.

But what matters is not how much CO3&#8212;or HCO3- are added by the CaCO3. What matters is how much acid can be neutralized by the CaCO3 addition. Basically you could add 1ppm CaCO3 to water and titrate this to determine its alkalinity as CaCO3

Why I ran into this problem is b/c I tried to check if a particular water report made sense. I.e. that the anions and cations are balanced and I found that the water I build from distilled water with CaCO3 doesn&#8217;t make sense as there were too few HCO3- counted.

Kai

 
As mkade and Scotty_g are ultimately saying - it is a pH issue. In distilled water CaC03 is practically insoluble. As the pH changes, so does the solubility. I must assume then that the disagreement in the calculations is probably the result of adjusting for differences in solubility with pH.

I do acknowledge that and I also see that when I add chalk to my brewing water. It just doesn’t dissolve. But I don’t care about that as I know that it will eventually dissolve in the mash as the mash will have enough acidity.

But what I’m interested in is determining how much acid neutralizing power the CaCO3 has that I add to the mash or brewing water. Many spreadsheets/calculators say ~1 equivalent (i.e. one CaCO3 can neutralize 1 H+) and I say it should have 2 equivalents (i.e. one CaCO3 can neutralize 2 H+).

Kai

 
Here's a clue I think that I found on wikipedia

Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.
CaCO3 + CO2 + H2O &#8594; Ca(HCO3)2

This tells me that it will only form (HCO3)2 in an amount equal to the dissolved CO2 in the water. It will be (HCO3 X ppm added) + (HCO3 X ppm of CO2 in sol'n), not HCO3 X ppm added X 2.

Many spreadsheets/calculators say ~1 equivalent (i.e. one CaCO3 can neutralize 1 H+) and I say it should have 2 equivalents (i.e. one CaCO3 can neutralize 2 H+).

So it looks like it is actually somewhere in the middle, depending on how much CO2 is dissolved in the water.


Therefore I suspect that once the C02 is used up, it is other ions that are reacting with the chalk allowing it to dissolve, but these ions won't give you more HCO3 as part of the reaction
 
I actually don't care if the chalk I'm adding dissolves or no. If it doesn't dissolve in the water it will dissolve in the mash. B/c the mash adds acids and with them H+ ions. These ions may or may not "gobble up" the HCO3- and HO- created by the CaCO3. All I'm interested in is how many H+ can the CaCO3 neutralize that I add. And that is 2 thus the alkalinity contribution of the addional chalk (no matter if it dissolves before or after doughing in) should be 1 ppm as CaCO3 for each ppm CaCO3 added.

Kai
 
Sorry for sounding like a broken record, but is a CaCO3 addition needed, or desirable, for someone building up a base mineral profile with RO/distilled water, while also using a buffering agent (a la 5.2)? It would seem a moot point, given the neutralizing power of the buffer...
 
Why would you go through the effort of designing a water and then add 5.2 which changes that water profile in an unknown way?

I just don't like the 5.2 stuff.

If you have 5.2 and really hard water and don't want to mess with water chemistry, I sugest blending the tap water with RO water and use the 5.2. Just 5.2. and RO water is not desirable as I don't know if the 5.2 adds any calcium. You need at least some Ca in the wort.

Kai
 
Why would you go through the effort of designing a water and then add 5.2 which changes that water profile in an unknown way?

I just don't like the 5.2 stuff.

If you have 5.2 and really hard water and don't want to mess with water chemistry, I sugest blending the tap water with RO water and use the 5.2. Just 5.2. and RO water is not desirable as I don't know if the 5.2 adds any calcium. You need at least some Ca in the wort.

Kai

The 5.2 stuff is a phosphate buffer, and should not affect flavor itself. It is not contributing calcium, magnesium, sulfates, chloride, etc that are usually being mentioned as important for beer flavor. Of course, some may experience a change in flavor because of what is no longer happenening (i.e. the change in pH will alter residual hardness). I myself have other curiosities about 5.2. Clearly, even if you use 5.2, you need to have some minerals (i.e. calcium for enzymes, yeast metabolism). I think more about the final product. While the mash pH may be 5.2, often the final pH of bottled beer is closer to 4 (I think some of this has to do with dissolved CO2, which lowers pH due to its acidity). I have wondered before if the phosphate buffer changes the final pH of bottled (or kegged) beer, and has some flavor effect in that way.
 
The drop of the pH during fermentation is not so much the CO2 but more the yeast. They create an acidic environment by excreting acidic substances so they have pH gradient between the outside and their inside. This is needed for the transport of nutrients into the cell. How much the pH falls is actually dependent on the buffering capacity of the wort. Interestingly enough a low pH wort pH can lead to a higher beer pH than a high wort pH. This has to do with additional buffers that are created by enzymes (phytase) that are favored by low mash pH. While this all is very interesting I have not looked into that yet nor do I know how important it is for practical brewing.

A lot of 5.2 could actually become a problem for a good beer pH if it creates to strong of a buffer that wants to hold the wort/beer at 5.2

Kai
 
The drop of the pH during fermentation is not so much the CO2 but more the yeast. They create an acidic environment by excreting acidic substances so they have pH gradient between the outside and their inside. This is needed for the transport of nutrients into the cell. How much the pH falls is actually dependent on the buffering capacity of the wort. Interestingly enough a low pH wort pH can lead to a higher beer pH than a high wort pH. This has to do with additional buffers that are created by enzymes (phytase) that are favored by low mash pH. While this all is very interesting I have not looked into that yet nor do I know how important it is for practical brewing.

A lot of 5.2 could actually become a problem for a good beer pH if it creates to strong of a buffer that wants to hold the wort/beer at 5.2

Kai

Yeah, it's definitely an interesting point. Five Star Chemicals says that it supercharges your wort, which I don't doubt. I guess the question is what does it do to the rest of the process. The problem is that those phosphates are highly soluble, and even if they don't directly affect flavor, you can't get rid of them, and their buffering capacity might affect the beer in other ways.
 
I actually don't care if the chalk I'm adding dissolves or no. If it doesn't dissolve in the water it will dissolve in the mash. B/c the mash adds acids and with them H+ ions. These ions may or may not "gobble up" the HCO3- and HO- created by the CaCO3. All I'm interested in is how many H+ can the CaCO3 neutralize that I add. And that is 2 thus the alkalinity contribution of the addional chalk (no matter if it dissolves before or after doughing in) should be 1 ppm as CaCO3 for each ppm CaCO3 added.
Kai

I understand what you are saying, the point I was trying to make was that when the chalk finally does dissolve, not all of the the ions produced by the carbonate ion will be as HCO3 so the buffering capacity won't be = to 2 eq. of HCO3 per CaCO3

CaC03 + H20 to Ca+ +(HCO3-)2 will only occur if you pump in enough CO2 into the solution. To balance the equation there must be extra C somewhere on the left side of the equation because as I've written it here the left side has 1Ca, 1C, 2H and 4O, and the right side 1Ca, 2C, 2H and 6O, making the left side short 1C and 20, hence the need for CO2. Once the CO2 in solution is used up, you won't get any more HCO3 to act as a buffer as the acids in the wort dissolve the remaining CaCO3
 
You don't need CO2 to dissolve CaCO3. Acid will do as well:

CaCO3 + 2H+ -> Ca++ + H2O + CO2

This is the reaction that i was talking about that "gobbles up" the H+ ions and neutralizes the acid. I.e. prevents the pH change.

Kai
 
Yikes, lots of stuff in this thread. I gotta stop working and hang out here all day.

Normally your mash will not be saturated in CO2. As you heat it, dissolved gases (namely CO2) will evaporate. When it *does* dissolve, it forms carbonic acid:
CO2 + H2O -> H2CO3. This will protonate into H+ / HCO3- or 2H+ and CO3--. Left to its own devices (nothing else in the wort) it will mainly exist as H+ and HCO3-, I believe (mkade's post describes this--it's related to the pKa).

CaCO3 will provide a little bit of buffering against some acids; again, depending on the material you could generate some CO2 and raise the pH. It's not a *good* buffer because it is so insoluble. Phosphate buffer (pH 5.2) is effective because it will consume leftover H+ to neutralize acid or release H+ to neutralize bases (going from PO4 to HPO4 to H2PO4 to H3PO4).

If there is no acid in the water, CaCO3 dissolves to form Ca++ and CO3--. But again, not very much of either. If there is acid, *then* CO3-- will start to consume it. If the other acid is strong (it sheds H+ easily) then you can convert CO3-- to H2CO3 (which then dissociates to form H2O and CO2). If the other acid is weak, you only end up with HCO3-. Phosphoric acid can lose 3 H+; the first one is lost easily (strong acid), the second with difficulty (weak acid) and the third with more difficulty (even weaker acid).

Mash acids are all weak acids (unless you lower pH with hydrochloric or phosphoric). They will lead to more CaCO3 dissolving if there is solid chalk at the bottom of the MLT, but not a lot of it.

The reason you use the 5.2 buffer is that it takes a lot of mash acid / alkalinity to move your pH away from the desired level. If you don't buffer, your pH will change more easily--so your carefully crafted water may change to an unwanted pH when you throw in the malt.

The phosphate provided by the buffer is relatively small in the grand scheme of things, and it is a yeast nutrient. Most of the phosphate will end up as yeast mass and settle out. The industrial wastewater treatment plant at work is fertilized with phosphoric acid, and 90-95% of it is consumed by the bacteria with only the small fraction remaining dissolved in the effluent.
 
Scotty, that's more or less what I thought. In response to Kaiser, I intend to build up a mineral profile of about 200 ppm, split between calcium, magnesium, sulfate, and chloride in proper amounts. This would be for flavor only. I would then use the 5.2 to make sure my mash pH was where it is supposed to be. As someone that has dealt with the fuss of pH meters, etc, I'd much rather use the buffer.

So with that in mind, for someone that's using a buffer for pH and minerals for taste, CaCO3 is rendered a moot point, yeah?
 
Kaiser, here is how Palmer calculates Contributed Alkalinity from 1 gram of chalk (CaCO3) in one litre of water to be 490:

- he states in Chapter 15, Table 16 of HowToBrew.com that 1 gram of CaCO3 in one gallon of water will dissociate into 105 ppm of Ca(+2) and 158 ppm CO3(-2)

- therefore, to convert to the CO3-2 concentration of 1 litre instead of 1 gal, multiply by 3.785 l/gal to get 598 ppm of CO3(-2) (expressed as HCO3 since it will readily associate with free H+ assuming pH is low enough)

- convert from ppm to Alkalinty (as CaCO3) as per Chapter 15, Table 13 by dividing by 61 and multiplying by 50: 598/61*50 = 490

Hope that was clear -- unfortunately vB doesn't support superscripts and subscripts, so the chemistry notation doesn't come through well.
 
- therefore, to convert to the CO3-2 concentration of 1 litre instead of 1 gal, multiply by 3.785 l/gal to get 598 ppm of CO3(-2) (expressed as HCO3 since it will readily associate with free H+ assuming pH is low enough)

This is where I think the problem lies. 598 ppm of CO3-- are not equivalent to 598 ppm of HCO3-. CO3 can neutralize 2 H+ while HCO3 can neutralize only 1 H+. Hence to get the same alkalinity equivalent of CO3-- expressed as HCO3- you need to multiply the CO3 ppm by 2 since HCO3 and CO3 roughly weigh the same (61 vs. 60).

Kai
 
BTW, it might be worth mentioning that on page 8 of Principles of Brewing Science, Fix provides the calculations of molecular weights of Ca(2+) and CO3(2-) to concentrations in water (his example is for a 0.1 gram addition of CaCO3 to 1 litre of water, but you can easily scale by a factor of 10 to get Palmer's numbers, accepting a small bit of rounding error).
 
This is where I think the problem lies. 598 ppm of CO3-- are not equivalent to 598 ppm of HCO3-. CO3 can neutralize 2 H+ while HCO3 can neutralize only 1 H+. Hence to get the same alkalinity equivalent of CO3-- expressed as HCO3- you need to multiply the CO3 ppm by 2 since HCO3 and CO3 roughly weigh the same (61 vs. 60).

Kai
Perhaps, but regardless the difference is slight.

This is actually discussed pretty well by Fix on pages 5 and 8 - 9 in POBS.
 
Sorry Kai -- I re-read what you were saying, and now I think I understand where you were going.

Actually, you are right that the CO3 ion has twice the equivalence of the HCO3 ion for determining alkalinity. In fact, Fix states that the equation for Alkalinity = [HCO3-] + 2*[CO3(2-)]. However, the CO3 ion concentration is not considered in brewer's estimation of alkalinity because CO3(2-) is not present in significant concentrations in the pH range that we are working with (i.e. 6 and below). So the latter part of this equation is generalized for ease of use, I imagine.

:mug:
 
However, the CO3 ion concentration is not considered in brewer's estimation of alkalinity because CO3(2-) is not present in significant concentrations in the pH range that we are working with (i.e. 6 and below). So the latter part of this equation is generalized for ease of use, I imagine.

Yes, CO3- is not really present in water at the common water pH levels. But if you add CaCO3 then you add 2 equivalents of HCO3- there is no way around that. CaCO3 can disassociate this way:

CaCO3 -> Ca++ + CO3—

This results in carbonate which we know doesn’t really exist in our water. So the carbonate needs to disassociate further:

CO3-- + H20 -> HCO3- + HO-

This produces one HCO3- but also yields HO- which would raise the pH since the [HO-][H+] product needs to remain constant. So there needs to be something else happening here. One option is the CO2 that is dissolved in the water:

HCO3- + CO2 + HO- -> HCO3- + HCO3-

Now we have 2 bicarbonate molecules and no HO- anymore. The pH would not change but the alkalinity was increased by 2 HCO3- equivalents.

But what if there is not enough CO2 for aforementioned reaction -> The CaCO3 will not dissolve. So it is not dissolved the water but it is still present as a suspension or just lies at the bottom of the kettle. Until we add the malt. At that point we add acids and with them H+ that will allow for this reaction:

CaCO3 + 2H+ -> Ca++ H20 + CO2

No HCO3- or CO3- is left but the CaCO3 neutralized 2 acid equivalents (2H+) hence it contributed 2 HCO3- worth of alkalinity. The latter is defined as the ability to neutralize acid.

So while a water analysis of water to which 1ppm CaCO3 has been added may not see a 1ppm increase of the alkalinity as CaCO3 b/c not all of the chalk actually dissolved, the effect that the 1ppm CaCO3 has on the mash when it comes to neutralizing acids is worth 1 ppm alkalinity as CaCO3. And this is what eventually matters for the mash pH.

I have been thinking about this for 2 days now and I am getting more and more convinced that I’m correct here. I also e-mailed one of the more knowledgeable brewers on HBD who seems to know much more about water than I (you should see his water spreadsheet :http://www.wetnewf.org/Brewing_articles/BURP_OCT08) and he agrees that 1ppm CaCO3 adds 1ppm alkalinity as CaCO3.

Kai
 
I ran an experiment intended to confirm my theory but it didn't rule in my favor:

I prepared 3 different waters:

RO: straight from the reverse osmosis tap. About 45 TDS, pH ~5.70 (was jumping around a
bit. I guess it is very weakly buffered), exact mineral composition is unknown

Water A: 1.5 l RO water + 0.12g CaCO3 + 0.44g CaCl2*2H2O. If CaCO3 adds 2 equivalents of
alkalinity the RO water's residual alkalinity was increased by 0 dH. If CaCO3 adds 1
alkalinity equivalent the RO water's residual alkalinity was lowered by 2.2 dH (~ 50 ppm
CaCO3)

Water B: 1.5 RO water + 0.23g CaCO3 + 0.23g CaCl2*2H2O. If CaCO3 adds 2 equivalents of
alkalinity the RO water's residual alkalinity was increased by 4.4 dH (~100 ppm CaCO3).
If CaCO3 adds 1 alkalinity equivalent the RO water's residual alkalinity remains
unchanged.

I then heated 3 200ml samples of this water to 64C in the microwave and added 50g of malt
to each. The malt was added at the same time. After 5 min I took a small sample from each
and cooled it to 22C. Then I measured the pH and here is what I got:

RO water mash : 5.76
Water A mash : 5.69
water B mash : 5.77

This means that the Water B had the same RA as the RO water. But for that to be true
CaCO3 has contributed only 1 alkalinity equivalent. Which does not make sense to me.
Could this be the result of an incomplete reaction by the chalk?

Now I have to figure out what exactly is going on here.

Kai
 
How much water did you use in your experiments (was that 1.5 gal, or 1.5 L)? How long did you let the chalk and CaCl2 sit before measuring hardness? You can only get 47 ppm of chalk to dissolve in water; I doubt the pH 5.7 RO water has enough acidity to help dissolve chalk. I have not done the math so I can't say for sure.

1 ppm chalk dissolved will raise alkalinity (or hardness) 1 ppm equivalent of CaCO3 (by definition).

Depending on your pH meter, it may not be able to distinguish that small a difference in pH.
 
1 ppm chalk dissolved will raise alkalinity (or hardness) 1 ppm equivalent of CaCO3 (by definition).

That was my point until I ran this experiment.

In the previous experiment the chalk didn't dissolve. Which is equivalent to just adding the chalk to the mash.

So I ran another experiment in which I dissolved the chalk w/ CO2:

In this experiment I prepared 3 waters:

* (A) RO as the control
* (B) RO + 80 ppm CaCO3 + 290 ppm CaCl2
* (C) RO + 80 ppm CaCO3 + 290 ppm CaCl2 + CO2

The water for B and C was mixed in the same bottle and some was poured off for B. I then added CO2 to the soda bottle's head space and started shaking. The bottle immediately contracted which was a sign that CO2 was dissolved in the water. After a while the solution cleared up. Once the water was clear again I was sure that most of the CaCO3 dissolved. I then prepared 3 mashes the same way I did in the last experiment and measured their pH in a cooled sample (21-22 C):

A: 5.69
B: 5.58
C: 5.68

I kept measuring between the samples and after a while it seemed as if the pH of B started to rise towards 5.70 (where the other two samples were). I'm not quite sure why that was happening. Looks like a reason to repeat this experiment with twice the amount of salts in oder to make the difference in RA more pronounced.

But it seems as if chalk dissolved in the mash water w/ CO2 (this is important as the use of an acid would defeat the purpose) is twice as potent in raising the alkalinity as chalk that is simply added to the mash or the strike water.

Why that would be the case is a mystery that I haven't solved yet. But I'm open to suggestions.

Kai
 

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