Using lime to increase alkalinity - when is enough too much?

[ajdelange]

I didn't get a chance to check the math yesterday so here goes. I'll work through the problem manually which, while it may be a bit challenging to read should be informative. The math is what's behind some of the curves in Palmer's book so it may be interesting to readers from that POV.

We start with 100 mg of what we presume is pure Ca(OH)2. This has a GMW of 74.093 grams per millimole so we have 100/74.093 = 1.34966 millimoles of slaked lime each of which contains two milliequivalents of hydroxyl (OH-) ion thus we are going to neutralize 2*1.34966 = 2.69931 milliequivalents of hydroxyl ion and to do that we need the same amount of hydrogen ions:
H+ + OH- ---> H2O.

At any pH there are some hydroxyl and hydrogen ions floating around but at pH 7 they are in equal concentration of 10^-4 milliequivalent/L each and we can neglect their presence. So we specify that we are going to add acid until the lime is all dissolved and the pH is 7. The question becomes "how much phosphoric acid is required to deliver 2.69931 mEq of protons (hydrogen ions) at pH 7.

Phosphoric acid, H3PO4 sheds protons as pH increases. The first shedding is of the first proton: H3PO4 --> H+ + H2PO4-. The ratio of the number of acid molecules that shed this proton to the number that don't is 10^(pH - 2.123) = 10^(7 - 2.213) = 75335.6. Thus, at pH 7, the vast majority of acid molecules will have given up this first proton and been converted to H2PO4-. In turn H2PO4- sheds one of its protons: H2PO4- --> H+ + HPO4--. The ratio of the concentration of those that have shed the second proton (dibasic) to that of those that have retained it (monobasic) is 10^(ph - 7.214) = 10^(7 - 7.214) = 0.610942 i.e. a bit over half of the monobasic ions will have yielded up the second proton. Then the dibasic ions can give up their single protons to become phosphate: H2PO4-- --> H+ + PO4---. The relevant ratio here is 10^(pH - 12.44) = 10^(7 - 12.44) = 3.63078e-06. Very few of the dibasic ions emit their protons at pH 7.

To determine the number of protons from each millimole of phosphoric acid at pH 7 then we assume that at pH 7 there remain P millimoles of H3PO4, note that there would be 75335.6*P millimoles of monobasic phosphate ion and 75335.6*0.610942*P moles of dibasic phosphate and 75335.6*0.610942*3.63078e-06*P moles of phosphate. Each of these species contains 1 phosphorous so the total amount of phosphorous in the solution is P + 75335.6*P + 5335.6*0.610942*P + 75335.6*0.610942*3.63078e-06*P = P*(1 + 75335.6 + 75335.6*0.610942 + 75335.6*0.610942*3.63078e-06) = 121362*P. The fraction of the total which is still phosphoric acid is P/121362*P= 1/121362 = 8.23979e-06. The fraction which is monobasic is 75335.6 times this or 0.620749 and the fraction which is dibasic is 0.610942 times the fraction which is monobasic or 0.610942*0.0.620749 = 0.379242 and the fraction which is phosphate is 3.63078e-06 times the fraction which is dibasic which is an insignificant 1.37694e-06.

Given that the charge on phosphoric acid is 0, the charge on a monobasic ion is -1, the charge on dibasic phosphate is -2 and on phosphate is -3 the total charge on phosphate species at pH 7 is
0*8.23979e-06 - 1*0.620749 - 2*0.379242 - 3*1.37694e-06 = -1.37924 milliequivalents of charge per millimole of phosphate in the solution. Since all the phosphate come from the phosphoric acid we add the total charge is -1.37924 mEq per mmol of added phosporic acid. Since each negative charge comes from the loss of a proton it is clear that if we need 2.69931 mEq of protons we will have to add 2.69931/1.37924 = 1.9571 mmol of phosphoric acid. As the molecular weight of phosphoric acid is 98 mg/mmol we will need 98*1.9571 = 191.176 mg of phosphoric acid. If we use 10% w/w acid we'll need 10 times the amount of solution or 1912 mg i.e. 1.912 grams to get 191.2 mg of the actual acid and as the density of 10% phosphoric acid is about 1.050 g/ml we'll need 1.912/1.050 = 1.82 mL.

I think that's pretty close to what I said earlier. And it checks with my spreadsheet now so I don't think I made a math error. That, of course, still leaves a mystery.

 

[Tutsbrew]

Perhaps we need to account for atmospheric CO2 absorption...

 

[ajdelange]

If the 100 mg powder consisted of 50mg Ca(OH)2 and 50 mg CaCO3 the acid requirement would still be 1.3 mL phosphoric. That's half a mL less which should certainly be detectable. The fizz test would confirm the presence of carbonate. I'll be back at my lab in a couple of weeks and will test some lime with hydrochloric acid of known strength. Maybe that will clear things up a bit.

 

[GotPushrods]

AJ -- I haven't yet taken the time to work through your math, but it might be easier to work this problem as 2 solutions. I don't remember the exact problem, but I believe you said 100 mg of calcium hydroxide in 1L of water, titrated to pH 7.

1. 0.100 g of Ca(OH)2 in 1000 mL water = 0.00135 M solution
2. 10% (w/w%) H3(PO)4 with your choice of volume = 1.1 M solution (pardon the lack of significant digits)

I will cheat and enter these into a calculator first just to see if I can corroborate with minimal work...

I like http://www.webqc.org/phsolver.php for stuff like this.

Once your learn their syntax, it's not too bad. Since this is for solutions, I entered the following:

For 1 ml of acid, I enter:

Ca(OH)2 c=0.00135 v=1000 pKb=2.37
H3(PO)4 c=1.1 v=1 pKa1=2.148 pKa2=7.198 pKa3=12.319

pH = 6.68

... and for 2 ml:

Ca(OH)2 c=0.00135 v=1000 pKb=2.37
H3(PO)4 c=1.1 v=2 pKa1=2.12 pKa2=7.21 pKa3=12.67

pH = 3.18

Since this doesn't jive with the original estimate, and does closely correlate to the experiments, my next step would be to verify this solution math by hand, then try and figure out what is off with the mEq method of calculation.

 

[GotPushrods]

Putting 100 mg of lime into 1000 ml of water gave me a pH of 11.15. Adding 1 ml of 10% phosphoric acid brought the pH down to 6.8. The full 2 ml of acid brought the pH down to around 3.38.

Just for fun, let's enter a few more of your observations and see how close they are to this concentration calculator.

Ca(OH)2 c=0.00135 v=1000 pKb=2.37

pH = 11.03

Pretty close. So what if your lime were only 50% pure?

Ca(OH)2 c=0.000675 v=1000 pKb=2.37

pH = 10.77

Your observed pH was actually higher than the 100% pure calculation, so far so good.

For kicks and giggles, I started over and I doubled the lime to 200 mg and ran a new test. 2ml of acid brought the pH down to 6.46.

Let's run that one too...

Ca(OH)2 c=0.0027 v=1000 pKb=2.37
H3(PO)4 c=1.1 v=2 pKa1=2.148 pKa2=7.198 pKa3=12.319

pH = 6.67

 

[ajdelange]

Deleted. Duplicate post.

 

[ajdelange]

1 L of 10% phosphoric acid weighs about 1050 grams and contains, therefore, 105 grams of phosphoric acid. As the molecular weight of phosphoric acid is 98 this is 105/98 = 1.07 molar. If we assume we are going to titrate to the second pH of phosphoric acid, 7.214 at 20 °C, then we know that every molecule (nearly every molecule) of phosphoric acid released a proton and that half of the resulting H2PO4- ions released a second proton. Thus each molecule of acid produces 1.5 protons and 1.07 M becomes 1.6 N. Note that at mash pH 10% H3PO4 is about 1 N which is handy. If I need 2*100/74.093 milliequivalents of OH- then I'll need (2*100/74.093)/1.6 = 1.687 mL of 10% phosphoric acid to realize it.

Now plugging this into your on line pH calculator as

Ca(OH)2 c=.00135 pKb1=1.4 pKb2=2.43 v=1000
H3PO4 c=1.07 pKa1=2.12 pKa2=7.21 pKa3=12.67 v=1.687

I get back
pH = 7.2025096546188

which is consistent. The mistake you made is not telling the program that Ca(OH)2 has two hydroxyls. It does not dope this out from the fact that you type Ca(OH)2. You get the same answer if you tell it it is xxxxZ. You must specify the second pKb as I have done above. It took me a while to figure this out. Then I asked myself "Why does it ask me for the pK's when I've already told it I'm using phosphoric acid?" Answer: it doesn't know it's phosphoric acid until you enter the pK's.

In my earlier calculations I just assume that Ca(OH)2 is a strong base (pKb1, pKb2 < -10) and go on that basis. Using the two values 2.43 and 1.4 which I found on some website gives 99.9996% dissociation which is close enough to 100% for government work.

So this program seems to vindicate my calculations. Big relief.

 

[GotPushrods]

1 L of 10% phosphoric acid weighs about 1050 grams and contains, therefore, 105 grams of phosphoric acid. As the molecular weight of phosphoric acid is 98 this is 105/98 = 1.07 molar. If we assume we are going to titrate to the second pH of phosphoric acid, 7.214 at 20 °C, then we know that every molecule (nearly every molecule) of phosphoric acid released a proton and that half of the resulting H2PO4- ions released a second proton. Thus each molecule of acid produces 1.5 protons and 1.07 M becomes 1.6 N. Note that at mash pH 10% H3PO4 is about 1 N which is handy. If I need 2*100/74.093 milliequivalents of OH- then I'll need (2*100/74.093)/1.6 = 1.687 mL of 10% phosphoric acid to realize it.

Now plugging this into your on line pH calculator as

Ca(OH)2 c=.00135 pKb1=1.4 pKb2=2.43 v=1000
H3PO4 c=1.07 pKa1=2.12 pKa2=7.21 pKa3=12.67 v=1.687

I get back
pH = 7.2025096546188

which is consistent. The mistake you made is not telling the program that Ca(OH)2 has two hydroxyls. It does not dope this out from the fact that you type Ca(OH)2. You get the same answer if you tell it it is xxxxZ. You must specify the second pKb as I have done above. It took me a while to figure this out. Then I asked myself "Why does it ask me for the pK's when I've already told it I'm using phosphoric acid?" Answer: it doesn't know it's phosphoric acid until you enter the pK's.

In my earlier calculations I just assume that Ca(OH)2 is a strong base (pKb1, pKb2 < -10) and go on that basis. Using the two values 2.43 and 1.4 which I found on some website gives 99.9996% dissociation which is close enough to 100% for government work.

Awesome, good catch! I knew you would find something. I have enough chemistry background that this stuff all makes sense to me, but my drawback is that I write software for a living. Sometimes it's hard to imagine that something might be done differently than I would obviously do it! Upon further inspection, it definitely is not parsing the chemical name. You can type "Pluto" for Ca(OH)2 and get the same result.

This is buried in the middle of the page:

For strong acids enter pKa=-10
For strong bases enter pKb=-10

I had tried that first and got terrible output as well, which explains a lot. You NEED both pKbs.


The more I think about it, this experiment might be tough for many people as the target pH is so close to our acid's pKa2. The measurements are going to be critical, although I still wouldn't expect 100% error in the amount of acid required.

 

[CadiBrewer]

1 L of 10% phosphoric acid weighs about 1050 grams and contains, therefore, 105 grams of phosphoric acid.

I have one more thought. I'm better understanding this stuff after reading the Water book twice, but I'm still far, far, far from understanding chemistry so forgive me the silly question, if in fact it is silly. Could my 10% phosphoric acid be more concentrated than 10%? I thought about it after doing the experiments. Since I switched to RO water and have been using sauermaltz instead of acid, I haven't used acid in forever. In fact, the bottle I have might be five or more years old. It has been sealed the entire time, but does the stuff become more concentrated if not stored correctly?

 

[CadiBrewer]

The more I think about it, this experiment might be tough for many people as the target pH is so close to our acid's pKa2. The measurements are going to be critical, although I still wouldn't expect 100% error in the amount of acid required.

Assuming my previous post about the acid being concentrated more than 10% is silly, I think the measurements are what's causing my problems. It might not be 100% error, but when trying to measure .1 grams of calcium hydroxide and 1.82 ml of acid, there is plenty of room to swing using my scale and graduated syringe. I was going to test my measurements by upping everything tenfold and use 1 gram and 18.2 ml in 10 l of water to see how measurements affect my findings.

The upside to all of this is that I'm learning as I go and it is quite fun.

 

[ajdelange]

The more I think about it, this experiment might be tough for many people as the target pH is so close to our acid's pKa2. The measurements are going to be critical, although I still wouldn't expect 100% error in the amount of acid required.

Right near a pK value the strength of the acid is most variable with pH. Between two pK's it is less so. Doing the titration (because that's what it really is) to a pH closer to mash pH would reduce error from this source somewhat.

 

[ajdelange]

Could my 10% phosphoric acid be more concentrated than 10%? I thought about it after doing the experiments. Since I switched to RO water and have been using sauermaltz instead of acid, I haven't used acid in forever. In fact, the bottle I have might be five or more years old. It has been sealed the entire time, but does the stuff become more concentrated if not stored correctly?

It could be. If purchased from a home brew supply source we doubt that it is carefully assayed. Over time if it loses water it could become more concentrated but if it were indeed 10% when you bought it half the water would have to evaporate for it to become 20%. More likely the concentration is 10 ± a couple of percent.

 

[ajdelange]

I was going to test my measurements by upping everything tenfold and use 1 gram and 18.2 ml in 10 l of water to see how measurements affect my findings.

That's probably a good idea.

 

[trentm]

The more I think about it, this experiment might be tough for many people as the target pH is so close to our acid's pKa2. The measurements are going to be critical, although I still wouldn't expect 100% error in the amount of acid required.

I did the acid test for my lime last night.

Methods and Materials:

Weight measurement: Old set of triple beams
Volume measurement: 1000 and 100 microliter pipettors
Acid: Lactic 88%
Calculated (thanks AJ) acid to titrate 1 g of Ca(OH)2: 2.294 mL

One gram of Ca(OH)2 was dissolved in 1 L of RO water in a 2 L erlenmeyer flask on a stir plate for ~5 minutes. The initial volume of acid was added and with the stir plate set to a low speed the pH probe was lowered into the solution to test the pH. Subsequent acid additions were added directly to the solution with the pH probe in place. The test was repeated once.

Results: Below

Discussion:
The only calibration I can do with my old balance is to set it to 0. While my pipettors are <1 year old, they are student grade. That said, I am more confident in the pipettors than the balance. The results of the two tests varied (I would guess significantly). The initial pH (before acid addition) was ~12.5. I do not trust my pH meter at the high range as I cannot calibrate it at pH 10. Only a small drop in the pH was observed after the initial dose of acid; after which only minute amounts of acid had apparently significant effects. I feel the difference in the 2 tests were due to variation in the balance and may account for the greater than expected amount of acid to titrate to pH 7. It is also worth noting that in the first test the final acid addition of 0.025 mL resulted in a drop in pH of 1.85 and in the second test the final acid addition of 0.05 mL dropped the pH by 2.76. These volumes of acid are equivalent to about a half drop and a drop, respectively.

 

[ajdelange]

The worth of my suggestion depends on having an acid that is of known strength. If you have to use more acid than what I calculated for 100% Ca(OH)2 that suggests that perhaps the acid is not really of 88% strength. To be sure you would need to use a standardized acid or standardize the acid you are going to use against a known base. Calcium carbonate is good for this but it would have to be known to be pure. In reality here we are comparing two food grade, not analytical grade, products in which sense perhaps my suggestion wasn't such a good one. It would be much better and one would have much more confidence in the test result if something like this http://www.hach.com/sulfuric-acid-standard-solution-1-000-n-1-l/product?id=7640199730&callback=qs were used. Then again there is the issue of measurement error.

 

[trentm]

The worth of my suggestion depends on having an acid that is of known strength. If you have to use more acid than what I calculated for 100% Ca(OH)2 that suggests that perhaps the acid is not really of 88% strength. To be sure you would need to use a standardized acid or standardize the acid you are going to use against a known base. Calcium carbonate is good for this but it would have to be known to be pure. In reality here we are comparing two food grade, not analytical grade, products in which sense perhaps my suggestion wasn't such a good one. It would be much better and one would have much more confidence in the test result if something like this http://www.hach.com/sulfuric-acid-standard-solution-1-000-n-1-l/product?id=7640199730&callback=qs were used. Then again there is the issue of measurement error.

For me, all of your suggestions were a good and fun. I will continue to use lime for mash pH adjustment. Here is the way I summarize the issue: We have came to realize that pickling lime can lose it's strength and caution should be taken when using this product for water adjustment. Although the "strength test" may not be practical for many homebrewers (myself at least), I have learned through this thread how to reasonably judge the quality of lime. Although the strength test proved unreliable in my situation, it was a good learning experience. By using the "fizz" test and observing the final solution for good formation of "lime milk" and excessive precipitation, I feel I can be confident enough to use this product to adjust mash pH. Although the unreliability of product strength adds another variable to predicting mash pH, there are plenty of other uncontrolled variables to suggest measurement and final adjustment. This stuff is under $5 at my local grocery. For that price I can replace my supply once or twice a year without even feeling it. For me it remains the best option for increasing mash pH. Thanks to the OP for asking the question and all involved in finding a solution.

 

[CadiBrewer]

For me, all of your suggestions were a good and fun. I will continue to use lime for mash pH adjustment. Here is the way I summarize the issue: We have came to realize that pickling lime can lose it's strength and caution should be taken when using this product for water adjustment. Although the "strength test" may not be practical for many homebrewers (myself at least), I have learned through this thread how to reasonably judge the quality of lime. Although the strength test proved unreliable in my situation, it was a good learning experience. By using the "fizz" test and observing the final solution for good formation of "lime milk" and excessive precipitation, I feel I can be confident enough to use this product to adjust mash pH. Although the unreliability of product strength adds another variable to predicting mash pH, there are plenty of other uncontrolled variables to suggest measurement and final adjustment. This stuff is under $5 at my local grocery. For that price I can replace my supply once or twice a year without even feeling it. For me it remains the best option for increasing mash pH. Thanks to the OP for asking the question and all involved in finding a solution.

I completely echo this. All of your suggestions and help has led to quite a bit of learning on my part. Thanks AJ for your guidance.

 

[craigevo]

The Water book and the information pages on B'run Water suggest that 5.4 to 5.5 is a good pH target for darker beers. It rounds out the roasted flavors and makes them not as harsh.

Does the book say how/why this rounding out works at this pH specifically?

I wonder if this is opposed to a higher or lower pH. Could it be more due to the final beer pH rounding out roasted flavors rather than the mash pH itself (?)

 

[ajdelange]

5.4 - 5.5 is a good mash pH range for dark beers and also a good mash pH for light beers. I don't recall the book having much to say about the why's and wherefore's. Based on what people have posted in this forum and others and my own experience it seems that as pH is lowered flavors, all flavors, seem to become 'brighter' (that's not my word but I just can't think of a better one) as pH is decreased. Clearly there has to be some point beyond which they become too bright, whatever that means, and I have had some conversations with John in which he expresses that opinion and, IIRC he uses terms like 'sharp'. I, for some reason, have not wound up with a mash pH too low....yet. Thus I can't confirm from personal experience what the consequences of too low a mash pH might be other than the obvious less than optimum.

What you taste will, of course, depend to some extent on the pH of the final beer and I know at least one brewer, Gordon Strong, who adjusts the final pH of his beer. Left alone beer goes to a pH that depends more on yeast strain selection and fermentation health than it does on mash/wort pH though, of course, they do have some effect.

 

[craigevo]

Thanks for the info AJ.

I do see people referencing a specific mash ph range for dark beers quite often, and I cannot remember once anyone saying why only for dark beers. Maybe I am reading too much into it.

 

[CadiBrewer]

5.4 - 5.5 is a good mash pH range for dark beers and also a good mash pH for light beers. I don't recall the book having much to say about the why's and wherefore's. Based on what people have posted in this forum and others and my own experience it seems that as pH is lowered flavors, all flavors, seem to become 'brighter' (that's not my word but I just can't think of a better one) as pH is decreased. Clearly there has to be some point beyond which they become too bright, whatever that means, and I have had some conversations with John in which he expresses that opinion and, IIRC he uses terms like 'sharp'. I, for some reason, have not wound up with a mash pH too low....yet. Thus I can't confirm from personal experience what the consequences of too low a mash pH might be other than the obvious less than optimum.

What you taste will, of course, depend to some extent on the pH of the final beer and I know at least one brewer, Gordon Strong, who adjusts the final pH of his beer. Left alone beer goes to a pH that depends more on yeast strain selection and fermentation health than it does on mash/wort pH though, of course, they do have some effect.

Finally got to my next batch. Every test confirms that my new bottle of pickling lime is about 50% effective. Brewed an Oatmeal Stout. Bru'n Water predicted that I would need 3 grams of pickling lime to get the pH to 5.4 for the 10 gallon batch. I did a test mash, and then ended up using 6.3 grams for the 10 gallon full batch and I think I nailed the pH. Tested at 7 minutes it was 5.54. At 15 minutes, it was 5.45. I didn't get a chance to test again until 60 minutes, at which time the pH was 5.3. I believe that this indicates the pH was in the right range for the greatest amount of time during the conversion.

Reading and testing has caused me to learn that the pH doesn't stay at one spot during the entire mash. Seems obvious now that I think of it, but a year ago when I started fiddling with water, I figured I'd measure the additions, test the mash pH, and it would magically be at the right pH. Now I realize that the pH starts high for my dark beers and then works down as the grain buffers the water pH. For the lighter beers, it is less dramatic, but the pH starts low and works up through the sweet spot. Sorry to ramble but this learning is a revelation to me.

As suggested multiple times by AJ, I'll be doing test mashes from here on out for almost all batches, unless I'm confident in my grain (not just the grain bill) and my salts. Seems like a small step to make brewing day have a greater chance of success.

 

[ubnserved]

This might be a case where the pickling lime purity is suspect. Pickling lime can degrade to chalk if it is in contact with moist air. Check the lime purity by adding a drop of strong acid onto a small amount of lime. If there is any 'fizzing', it means that there is chalk in the lime and its strength is compromised.

Strong acid. .... like what? Sorry if this is covered later in the thread I have to ask while I am thinking about it or I will forget.

 

[ubnserved]

Please, don't give up on pickling lime. I recently tried an unorthodox use of pickling lime to bring my pH up in order to add enough lactic acid to a Saison to bring out a wonderful tartness after the primary fermentation. Then the beer was finished with Brettanomyces in the secondary. I do not have the language skills required to properly illustrate the effects of this process. I can say, after 33 years of brewing, I have never had anything this close to perfection.

I fear baking soda would not be as advantageous due to issues with high sodium. Certainly, if we all work together we can find a solution to the quality issue.



I will try these tests on my supply of pickling lime and report the results. I also use "Mrs Wages" brand and have noticed it seals with a zip type closure that may not be effective in protecting the product from exposure to air. In the future I will look for better containment. Any suggestions?

Is moisture or air the cause of problems with the picking lime?.... both issues could be removed by using a vacuum pump and sealer. I would suggest that when you purchase the product break it down to smaller packages and reseal in brew size portions. I would also suggest only buying this brand in late summer when product would be restocked at the store. When I was younger I worked at a grocery store and the packages left after canning season would sit on the shelf for a year.

Additionally if the product is turning into chalk and the chalk is removed with the trub what would be the problem with adding more to get your desired numbers?

Again if this has been answered sorry. If I don't ask questions as I go I will forget and I am just starting to study about chemistry part of brewing. Boredom and brain injury don't go well together.

 

[ubnserved]

I never thought that when I started reading this thread I would need a chemical engineering degree. Hats off to you guys that understand all of this, I am glad you are there doing the hard work I will read it all again and will eventually be able to remember and understand the math and symbols involved in the calculations. Right now my memory is just too sketchy.

 

[ajdelange]

As suggested multiple times by AJ, I'll be doing test mashes from here on out for almost all batches, unless I'm confident in my grain (not just the grain bill) and my salts. Seems like a small step to make brewing day have a greater chance of success.

After brewing a particular beer several times you will have that confidence and a test mash will not be necessary. When you get to that point it is still wise to check the pH in the mash tun (the mash itself becomes the test mash). If there is a discrepancy (because you used 60L crystal from maltster B instead of from Maltster A) just note it in your log and make an adjustment next time you use maltster B's stuff. A typical mash will have buffering of around 20 mEq/Lb-pH so that if you mash 10 lbs of grain and are off by 0.1 pH you will need approximately 20*10*.1 = 20 mEq of acid or base to adjust out that 0.1 pH. Note that 88% lactic acid is about 12 N and 10% phosphoric acid about 1 N. Thus you would need 20/12 ml of lactic or 20 mL of the phosphoric to make the 0.1 pH adjustment lower. Going in the other direction it would take 61 mg of sodium bicarbonate or 40 mg of lye to absorb 1 mEq of protons so you would need 20*61 =1.22 grams of sodium bicarbonate or 20*40 = 800 mg of lye to raise the pH by 0.1.

 

[ajdelange]

Is moisture or air the cause of problems with the picking lime?

It is the CO2 in the air that is the root of the problem. Whether moisture is required or not I don't know. In a lime kiln the reaction is
CaCO3 ---> CaO + CO2
which reaction can go in the opposite direction as well, if slowly. If moisture is added to the CaO:
CaO + H20 ---> Ca(OH)2.

Clearly
Ca(OH)2 + CO2 ---> CaCO3 + H2O
is possible as it is the reaction by which cement sets.

Additionally if the product is turning into chalk and the chalk is removed with the trub what would be the problem with adding more to get your desired numbers?

The problem is that calcium carbonate reacts very slowly but it does react. Just yesterday I was experimenting with this and observed, a bit to my surprise, a turbid suspension of CaCO3 with pH 3! Now that pH was climbing over time but climbing pretty slowly. By continuing to add acid the turbidity was removed and I has a clear solution at pH 3 but the pH continued to climb. That means that there were still tiny, tiny crystals of CaCO3 absorbing hydrogen ions.

Undissolved CaCO3 will continue to dissolve (and in so doing absorb protons) for the duration of the mash and the lauter. It is probably that some of the tiny microcrystals referred to above will make it through the grain bed into the beer and certainly biacarbonate ion derived from the chalk in the lauter tun will.

 

[CadiBrewer]

After brewing a particular beer several times you will have that confidence and a test mash will not be necessary. When you get to that point it is still wise to check the pH in the mash tun (the mash itself becomes the test mash). If there is a discrepancy (because you used 60L crystal from maltster B instead of from Maltster A) just not it in your log and make an adjustment next time you use maltster B's stuff. A typical mash will have buffering of around 20 mEq/Lb-pH so that if you mash 10 lbs of grain and are off by 0.1 pH you will need approximately 20*10*.1 = 20 mEq of acid or base to adjust out that 0.1 pH. Note that 88% lactic acid is about 12 N and 10% phosphoric acid about 1 N. Thus you would need 20/12 ml of lactic or 20 mL of the phosphoric to make the 0.1 pH adjustment down. Going in the other direction it would take 61 mg of sodium bicarbonate or 40 mg of lye to absorb 1 mEq of protons so you would need 20*61 =1.22 grams of sodium bicarbonate or 20*40 = 800 mg of lye to raise the pH by 0.1.

This is really great info. Thanks AJ

 

[ajdelange]

Strong acid. .... like what? Sorry if this is covered later in the thread I have to ask while I am thinking about it or I will forget.

"Strong acid" is somewhat ambiguous. It can mean concentrated but that actually is not a requirement here. 'Strong' here means it has at least one pK appreciably less than the strongest pK of carbonic acid which is 6.38 so anything with a pK less than that, even if dilute, will do. For example household vinegar which is only 5% acetic acid is OK because the pK of acetic acid is 4.76. Either the 88% lactic acid or 10% phosphoric acid sold by home brew suppliers or the hydrochloric acid sold by hardware stores or battery acid (sulfuric) would qualify.

 

[mabrungard]

The problem is that calcium carbonate reacts very slowly but it does react. Just yesterday I was experimenting with this and observed, a bit to my surprise, a turbid suspension of CaCO3 with pH 3! Now that pH was climbing over time but climbing pretty slowly. By continuing to add acid the turbidity was removed and I has a clear solution at pH 3 but the pH continued to climb. That means that there were still tiny, tiny crystals of CaCO3 absorbing hydrogen ions.

Undissolved CaCO3 will continue to dissolve (and in so doing absorb protons) for the duration of the mash and the lauter. It is probably that some of the tiny microcrystals referred to above will make it through the grain bed into the beer and certainly biacarbonate ion derived from the chalk in the lauter tun will.

Thanks AJ!

As I suspected, suspended chalk could make it through the mashing stage into the kettle where it eventually dissolves and pushes the wort pH higher than desired. Not a good thing.

Note to All: Stay away from chalk additions...they definitely don't work as intended.

 

[CadiBrewer]

Thanks AJ!

As I suspected, suspended chalk could make it through the mashing stage into the kettle where it eventually dissolves and pushes the wort pH higher than desired. Not a good thing.

Note to All: Stay away from chalk additions...they definitely don't work as intended.

If my lime is 50% effective and I add twice the lime that Bru'n Water calls for to get my pH in the sweet spot, will that create too much chalk in the kettle?