Re: Palmer Spreadsheet Error

Wed Dec 15, 2010 7:14 am

100 mg of calcium carbonate added to 1 L of water results in alkalinity, by the official definition, of close to 100 mg/L as calcium carbonate. What a surprise! But the Palmer spreadsheet and those based on it calculate only half that amount. So if you decide to use one of those spreadsheets for a dark beer it will tell you that you need RA about 7 times what might be considered reasonable (though still questionable) and then call for twice the amount of chalk required to get the 7 times. Result: RA will be 14 times what one wants.

As has been noted the situation isn't always quite that disastrous and the huge amount of chalk called for will not dissolve and thus not contribute its full weight to the alkalinity. This may be caused by reaction speed as much as anything else but there certainly isn't enough acid in grist to dissolve these huge chalk mounts. Calculation of the equilibrium pH would be very difficult but pH certainly isn't going to be in the 5.2 - 5.5 range in such cases.

As these spreadsheets do not ask about pH there is no way they can model chalk and bicarbonate properly. Also, there is really only one reason to add chalk to brewing water and that is to emulate a natural carbonaceous water. Natural carbonaceous waters arise when limestone (chalk) is dissolved by carbonic acid. These spreadsheets do not model that process either. Fortunately, if pH is less than about 8.3 and no chalk additions are specified these shortcomings really don't matter. And there is seldom, except when trying to model particular water, need for chalk addition.

The EZ spreadsheet doesn't fix all these problems but does, at least, caveat them.
ajdelange
 
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Re: Palmer Spreadsheet Error

Wed Dec 15, 2010 9:52 am

AJ brings up a good point that chalk is difficult to dissolve and contribute its alkalinity. But correspondingly, the chalk's calcium that is countering the alkalinity in the RA equation is also still bound up and not ionized. The overall effect would be that chalk does nothing until there are enough H+ protons to ionize the chalk.

We know that bubbling CO2 through water provides carbonic acid which is sufficient to ionize the chalk. Carbonic acid has a pKa of 6.3. Bubbling air through chalk water is known to dissolve significant concentrations of chalk. The question that I can't answer is what the pKa of Phytic acid is (phytic acid is the 'active ingredient' in phytin). I looked in Malting and Brewing Science and it doesn't get into this reaction too heavily and doesn't provide a pKa. Off hand, since chalk doesn't seem to dissolve very readily in the mash, it appears that phytic acid probably has a less active pKa and can't provide the protons the chalk dissolution needs.
Martin B
Carmel, IN
BJCP National
Foam Blowers of Indiana (FBI)

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mabrungard
 
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Re: Palmer Spreadsheet Error

Wed Dec 15, 2010 11:08 am

Phytic acid has 12 protons and thus 12 pK's and here they are: 0.573, 1.111, 1.163, 1.879, 1.918, 1.934, 2.978, 5.980, 7.491, 9.620, 11.589, 12.100, 12.679. Now I wouldn't take those values to the bank if I were you. They were obtained by titrating a solution of phytic acid with sodium hydroxide, plotting the pH vs mEq base curve and then building a model with 12 pK's whose pH vs base addition curve was calculated based on 12 initially guessed pKs which were then adjusted quasi randomly (simulated annealing) until there was no further improvement in the rms error between the model and the observed data. Simulated annealing is powerful but not elegant and so I have no idea how "observable" the 12 pK's are so I can't calculate the covariance matrix for the estimates i.e. I don't know if those numbers should be, e.g. 1.111 ±0.001 or 1.111 ± 10. Somewhere in between I expect. Anyway, even if I had great confidence in those numbers I couldn't calculate the probable pH of a system in which there is undissolved chalk because the only such calculations I can do are (thermodynamic) equilibrium calculations and in my gut I very much doubt that a seriously overdosed mash ever comes to equilibrium. In a carbonic/bicarb/carbonate only system (i.e. water over limestone in equilibrium with the air) it is the carbon dioxide content of the air which sets the equilibrium conditions (pH, hardness, alkalinity) of the resulting solution and, as we know, it can take days or longer for equilibrium to be reached. That simple system requires satisfaction of 6 simultaneous equations. Simplifying the phosphate question appreciably by assuming all phytin is converted to inorganic phophate (and myo-inositol) adds 4 more, trying to model the phytin itself would obviously take 13 more (1 for each pK and one for apatite solubility). There are, of course, other sources of acid in a mash besides the protons released by reaction of calcium with phosphate. IOW this is a problem that is well beyond me. Even if I could come up with models I'd have little confidence in the pK's I was feeding into it.
ajdelange
 
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Re: Palmer Spreadsheet Error

Thu Dec 30, 2010 9:41 am

The problem with chalk additions in some water programs is that they are calculating the amount of carbonate added to water with the chalk addition. Since chalk is calcium carbonate, it seems perfectly logical that the carbonate number is the correct one. Unfortunately, carbonate chemistry is a little more complicated than that.

First of all, in the typical drinking water pH range, carbonate (CO3) will mostly not exist in the water. It naturally converts to the bicarbonate (HCO3) form in the typical drinking water pH range. Many of you will recognize that the carbonate species exist in different ionic forms depending upon the pH of the system. It exists as carbonic acid (H2CO3) at low pH, as bicarbonate (HCO3) at middle pH, and as carbonate (CO3) at high pH. For the most part, our brewing is in the middle pH range and bicarbonate is prevalent.

The most important reason that we need to convert the carbonate concentration to its equivalent bicarbonate concentration is that the formula that we're using to calculate alkalinity assumes that everything is in the bicarbonate form. So we have to convert our alkalinity producers to their equivalent bicarbonate concentration.

Since the milliequivalent concentration of the carbonate species will not change when they transform to its various forms, we can calculate what the equivalent amount of each species (in mg/L) is as it transforms and find that numerical conversion value.

Some important chemistry information:

milliequivalents per liter is equal to the ionic concentration divided by the ion's equivalent weight.

The ion's equivalent weight is equal to the ion's molecular weight divided by the ion's charge.

For Carbonate, the eq wt = 60 mg/mole divided by its charge (-2), or 30 mg/mole
For Bicarbonate, the eq wt = 61 mg/mole divided by its charge (-1), or 61 mg/mole

I'm going to add another alkalinity ion for something else I'll present later. The equivalent weight of hydroxide (OH) is 17 mg/mole divided by it charge (-1), or 17 mg/mole.

Since the milliequivalents per liter do not change when we convert from one form of carbonate ion to another, we can calculate what that numerical conversion from carbonate to bicarbonate is. That conversion is simply the ratio of the equivalent weights of the ions. In the case of carbonate and bicarbonate, that ratio is 61/30 or 2.033333333. To convert a calculated concentration of carbonate ion to its actual concentration of bicarbonate ion at the typical drinking water pH range would be to multiply the carbonate concentration by 2.033.

So for a typical 1 gram per gallon chalk addition, the calcium concentration would be 105.7 ppm. But instead of the 158.4 ppm carbonate concentration, the bicarbonate concentration is actually 322.3 ppm (158.4 x 2.033). Note that the mEq/L are equal: 158.4/30 = 5.28 and 322.3/61 = 5.28.

The real problem with chalk is that it just isn't that soluble in water. There are entire book chapters written on the subject of calcium carbonate solubility since it is critical to life and critical to potable water supply engineers like myself.

At standard temperature and pressure (STP), the solubility of chalk is about 47 mg/L, which is not that much. That equates to less than 0.2 grams of chalk in each gallon of water. Those of you that use chalk know that it just doesn't seem to dissolve in water. You can bubble air through the water to get it to dissolve faster, but if you're working with air at atmospheric pressure, then you're only going to get that 47 mg/L into the water. That amount of chalk provides about 55 ppm HCO3 or about 45 ppm alkalinity, which may not be enough for the typical brown or black beer mash.

Work by Troester and DeLange have confirmed that chalk solubility in the mash isn't much higher. Apparently, the acids present in the mash are pretty weak and cannot provide the protons needed to dissolve the chalk. It takes extra effort in the form of adding CO2 to the water to get the chalk to dissolve in water.

I have done tests with water and chalk added at a rate of 2 grams per gallon and have easily dissolved it when I added CO2 to the headspace of the soda bottle and pressurized to over 15 psi with a carbonator cap. This improves the solubility by over 10 times, but that may not really be practical if your dealing with water needed for a 14 barrel mash.

To add alkalinity to mashing water we can also add baking soda (NaHCO3), but then we have to worry about a practical limit for sodium (150ppm, but it should really be kept below 100 ppm).

So, we need another option to add alkalinity to their mashing water.

Pickling Lime (aka Slaked Lime) is calcium hydroxide (Ca(OH)2). It is very soluble in water and does not face the solubility problems that chalk has. But I haven't seen anyone discussing how it should be added.

We need to go through the same milliequivalent/liter game that we went through with the carbonate/bicarbonate transformation. The ratio of equivalent weights between bicarbonate and hydroxide is 61/17 = 3.588.

Therefore, the concentration of calcium added when 1 gram of pickling lime is added to 1 gallon of water is 142.8 ppm and the concentration of hydroxide is 121.1 ppm. Converting that hydroxide concentration to its equivalent bicarbonate concentration is: 121.1 (ppm OH) x 3.588 = 434.7 ppm.

As you might expect with a strong base like pickling lime, it has pretty high alkalinity producing potential. When its added in the small amounts needed to control mash pH, it doesn't really convert into bicarbonate in the mash. It just consumes any acid it comes in contact with, converting those OH ions directly into H2O when an acid (H) is encountered. Since Alkalinity is defined as the measure of the capacity of a water to neutralize strong acid, it doesn't matter that the alkalinity is from carbonate, bicarbonate, or hydroxide. But since our brewing chemistry analyses are based alkalinity calculated from bicarbonate content, it is important to perform the conversion of hydroxide to its equivalent bicarbonate concentration.

Since the issue of errors in some water calculation programs was the genesis of this discussion, I should end with its discussion. Those programs assume the carbonate concentration can be treated as a reduced concentration of bicarbonate. Considering the limited solubility of chalk, its not a bad assumption. Unfortunately, the severely limited solubility of chalk make even that assumption to optimistic unless the brewer is going to dissolve the necessary quantity of chalk under CO2 pressurization. In addition, if the brewer does actually use CO2 to dissolve the chalk in the water, then the alkalinity calculated for the chalk addition would definitely be wrong with those water calculation programs. The real solution to adding alkalinity (without too much sodium) is to get brewers up to speed with using pickling lime for adding mash alkalinity and forget about chalk.

I trust this information will be helpful.
Martin B
Carmel, IN
BJCP National
Foam Blowers of Indiana (FBI)

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mabrungard
 
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Re: Palmer Spreadsheet Error

Fri Dec 31, 2010 6:26 am

The problems with carbonate additions stem from the fact that you can't buy calcium bicarbonate as a powder. If we could do that we could just add as much calcium bicarbonate as was needed to get a desired level of alkalinity paired with calcium and be done with it. OTOH we can make solutions of calcium bicarbonate but to do so in a manner consistent with what nature does we have to imitate her and though we all studied that in grammar school (remember stalactites and stalagmites) none of the popular spreadsheets model what she does which is dissolve limestone with carbonic acid (CO2 gas from subterranean respiring bacteria dissolved in water). The amount of limestone which dissolves per liter depends on the pressue of CO2 (and temperature) as does the pH. From the pH it's easy enough to calculate the relative amounts of carbonic, bicarbonate and carbonate in the solution and hence its harness and alkalinity (I'll put that stuff at the end) but you need to know the total amount of carbo (carbonic plus bicarbonate + carbonate) dissolved and calculating that is a bit tricky even if you know the partial pressure of CO2. So a more practical approach is to decide how much bicarbonate you want and the pH you want it at, then calculate the distribution and from that the carbo. This tells you how much CaCO3 to add. But you can't just add it. You have to dissolve it and get the solution to the desired pH and that requires protons. In nature, as noted earlier, these come from CO2. Home brewers tend to think that any acid will do (and it will in one sense i.e. it will dissolve the stuff) but in the sense of duplicating nature it doesn't do at all. What one does is suspend the chalk in the water and sparge with CO2 until the water is clear and the pH on target or a bit lower. The resulting water will have the desired properties but will not be at equilibrium (any more than the water from my well is at equilibrium in my HLT). Gradually, CO2 will escape, the pH will rise and the chalk will precipitate out. "Gradual" is at room temperature. In an HLT it can happen in minutes which, of course, raises the question "why bother" and the answer is really that there are precious few times when it is worth going to all the trouble of sparging with CO2. A possible motivation might be that you wish to make authentic beer using authentic water and methods including water treatment as carried out by the original brewers of this style.

The calculations I've hinted at are quite capable of being incorporated into a spreadsheet (they are in the one I use) but the interpretation of the spreadsheet and its use would have to be considered "advanced". The popular spreadsheets run into trouble because they ignore all this (if you are not asked to input source water pH, target water pH and alkalinity titration end point pH you are dealing with a model that will not properly handle carbonate or bicarbonate additions) but as long as carbonate additions are avoided and water pH is below 8.3, the simpler models will generally do.

As for adding CaO or Ca(OH)2 to overcome excessively low pH in the mash tun - certainly. Or NaOH or Na2CO3 (all food grade, mind you). But if you find you are having to do that you are doing something wrong. The only time I have ever had to do that was when I slipped with acid (i.e. I did something wrong). Most beers come in at higher than desired mash pH - not lower. OTOH CaCO3 should do the job. The amounts of CaCO3 required to move a mash from 4.9 to 5.2 should dissolve and be effective. Of course the nice thing about Ca(OH)2 (I don't think anyone should be using quicklime) is that it does not leave bicarbonate residue. Bicarbonate doesn't taste that great (IMO).

Distribution of species: You will find this in tables and graphs in texts implying there is great mystery about it. It's actually pretty simple (if you ignore Debye-Huckel and even that doesn't complicate things much):

r1 = 10^(pH - 6.38) = (moles of bicarbonate)/(moles of carbonic)
r2 = 10^(pH - 10.38) = (moles of carbonate)/(moles of bicarbonate)

d = 1 + r1 + r1*r2

f1 = 1/d = fraction of total carbo which is carbonic
f2 = r1*f1 = fraction of total carbo which is bicarbonate
f3 = r2*f2 = fraction of total carbo which is cabonate.

That, in 6 lines, is at the core of brewing water chemisty. But it's not in most spreadsheets.

At pH 6.38 about half the carbo will be carbonic and most of the rest bicarbonate with a trace of carbonate. At pH 8.3 about 98% of carbo molecules will be bicarbonate with about 1% each carbonic and carbonate. At pH 10.38 a bit less than half will be carbonate, the same fraction bicarbonate and a trace carbonic.
ajdelange
 
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