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

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CadiBrewer

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My last batch was a porter with a really acidic grist. 10 gallon batch, grist was:

Maris Otter - 13 lbs 1 oz
Brown malt - 2 lbs 12 oz
Crystal 75 - 1 lbs 13.5 oz
Chocolate malt - 1 lb 6 oz

In order to get the pH in the right range, I added 8 grams of lime to my 6.7 gallons of mash water. Brewer's Friend predicted pH of 5.46 at that amount of lime. I ended up with a pH at 15 minutes of 5.31. I figure that next time I need to go to 11 grams to hit a pH of between 5.4 and 5.5, which is my target.

With my other salt additions, according to Brewer's Friend 8 grams of lime brought my water to:

Ca 97.7 ppm
Mg 8.3
Na 10.6
Cl 35.4
SO4 54.9

is there a problem adding that much lime to get my pH in the intended range?
 

ajdelange

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Something fishy here. The GMW of Ca(OH)2 is 74.093. Eight grams of that is 1000*8/74.093 = 107.972 millimoles each of which releases 2 mEq of OH- for 216 mEq total. The crystal and the chocolate together raised to pH 5.4 would require about 75 mEq. Don't know what to do with the brown malt but I don't expect its as acidic as either of those but maybe it is. Then you are going to have a proton deficit of about 60 from the base malt. Seems your pH with that much chalk would be about 5.9 treating the brown malt as being about as acidic as Munich II (which may be a mistake).

Anyway, as long as you don't go overboard in the high pH direction with lime you can use as much as you want because the reaction is

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

or put in words: all the hydroxide is turned into water leaving just the calcium ion in solution.
 
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CadiBrewer

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Something fishy here. The GMW of Ca(OH)2 is 74.093. Eight grams of that is 1000*8/74.093 = 107.972 millimoles each of which releases 2 mEq of OH- for 216 mEq total. The crystal and the chocolate together raised to pH 5.4 would require about 75 mEq. Don't know what to do with the brown malt but I don't expect its as acidic as either of those but maybe it is. Then you are going to have a proton deficit of about 60 from the base malt. Seems your pH with that much chalk would be about 5.9 treating the brown malt as being about as acidic as Munich II (which may be a mistake).

Anyway, as long as you don't go overboard in the high pH direction with lime you can use as much as you want because the reaction is

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

or put in words: all the hydroxide is turned into water leaving just the calcium ion in solution.
I too continue to feel that something fishy is going on with my dark grists. B'run Water predicts 6.0 with that much lime, so almost exactly what you estimate. But my test mashes and my actual batches indicate that my pH ends up much lower than that. Brewer's Friend predicts 5.46, which is closer to my results, but still higher.

My practice is to weigh out my salts and put them in the water the night before when I set up my brew stand. That way I can fire up early the next morning. Does something happen with the water chemistry over night that would weaken the effect of the lime?

Also, where do people get their lime? I use Mrs Wages pickling lime from Amazon. Is there a chemical grade calcium hydroxide that might be better packaged and handled than my stuff?
 
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CadiBrewer

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Why would you want your mash to go higher? 5.1-5.3 should be your target.
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.
 

ajdelange

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Why would you want your mash to go higher? 5.1-5.3 should be your target.
I don't think so. 5.1 at room temperature would be under 5 at mash temp and 5.3 at room temp would be between 5 and 5.1 at mash temp. Too low in either case. A good target at room temperature is 5.4 - 5.5 for any beer. If, of course, you experiment and find that 5.3 gives you beer that you like better then use 5.3.
 
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CadiBrewer

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The first question would be "How are you measuring pH?"
With a Hach Pocket Pro+ pH meter. I have been using it for about a year, have done the stability test, and have followed the calibration sticky as closely as I can. In practice, I keep shot glasses in the freezer while brewing. I pull a sample from my mash using the shot glass and cool it down to under 90 degrees fairly rapidly (usually within 30 to 60 seconds). I read the pH and then clean my meter with distilled water using a squirt bottle and dabbing with a paper towel before putting the meter back in the 4.01 buffer. My little meter always settles back to between 3.99 and 4.02 when in the 4.01 buffer. I love the meter but hate the results :)

Though my readings are way off from what is predicted by the software program, I'm fairly confident in my testing process after a year of practice, though not so confident that I don't question that there could be user error somewhere that is causing the fishy readings.

My last three brews have been dark grists and have all had the same issues. I'm going to buy new pickling lime to rule out storage issues with the pickling lime being a culprit. That being said, on at least one test mash, I used baking soda instead of my pickling lime and I had the same under-predicted results, so I'm leery that that is my issue.
 

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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.
 

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With a Hach Pocket Pro+ pH meter.
OK - that's solid then.

My last three brews have been dark grists and have all had the same issues. I'm going to buy new pickling lime to rule out storage issues with the pickling lime being a culprit.
You can test the 'strength' of you lime quite easily. Put 100 mg in a liter (or any convenient volume) of water and measure the pH. Then add phosphoric acid until pH 7 is reached. It should take 1.95 mL of 10% acid to hit 7. If it does, the strength of your lime is 100%. If it takes half 1.95 mL then it is 50% and so on.

Now note that even if the lime has been exposed to CO2 and fizzes when acid is dropped on it that it's alkalinity (strength) is still the same (wrt mash pH) but that in the mash that strength might not be realized quickly. This is why we use lime instead of calcium carbonate.

That being said, on at least one test mash, I used baking soda instead of my pickling lime and I had the same under-predicted results, so I'm leery that that is my issue.
That suggests that something (one of the malts) is appreciably more acidic than we think it is. This is quite possible.
 
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OK - that's solid then.



You can test the 'strength' of you lime quite easily. Put 100 mL in a liter (or any convenient volume) of water and measure the pH.
I will definitely do this tonight when I get home and post the results. How do I measure 100mL of powdered lime?
 
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You can test the 'strength' of you lime quite easily. Put 100 mg in a liter (or any convenient volume) of water and measure the pH. Then add phosphoric acid until pH 7 is reached. It should take 1.95 mL of 10% acid to hit 7. If it does, the strength of your lime is 100%. If it takes half 1.95 mL then it is 50% and so on.
It turns out that the lime is the problem. 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.

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.

My lime seems to be just below 50% effective. Ordering new lime now. Hopefully this solves the issue I've been having with my darker grists and it also gives me a way to check my lime in the future if I'm having issues. Thanks AJ. Seriously, thanks. I brewed back to back batches of a dark mild before this porter that had me pulling what little hair I have left out because my pH was unfathomably low even on the re-brew.
 

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So having established that half of this stuff isn't either calcium hydroxide or calcium carbonate (did it fizz?) the obvious question becomes 'What is it?' and the not so obvious question is then 'Is this typical of the pickling limes we buy at the super market?'. Perhaps we should be advising all users to do this test and the fizz test on any lime they contemplate using.
 

trentm

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It did get answered in #2. In case that isn't clear the answer is no, there is no problem with adding as much lime as you need to get pH right unless you are concerned about high calcium.
Thanks and sorry I missed your earlier response. My brain will not accept chemistry and I guess it turned off before it got to part I was looking for.
 

mabrungard

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So having established that half of this stuff isn't either calcium hydroxide or calcium carbonate (did it fizz?) the obvious question becomes 'What is it?' and the not so obvious question is then 'Is this typical of the pickling limes we buy at the super market?'. Perhaps we should be advising all users to do this test and the fizz test on any lime they contemplate using.
AJ, what about a solubility test to assess lime purity? With the high pH of a distilled water solution of pickling lime, any chalk impurity should be prevented from going into solution and it should deposit itself on the bottom of the vessel.

If we use a vessel like a glass test tube, you should be able to see any sediment in the tube following an appropriate shaking time to dissolve the lime. I haven't worked out an appropriate lime dosing rate, but something like a gram or so per 100 mL of distilled water is probably sufficient.

This is not a very quantitative test, but it would serve as a qualitative test like the acid fizz test.

This whole pickling lime purity issue is troubling. It sure looks like baking soda use may be a surer way to add alkalinity to mashes.
 

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This whole pickling lime purity issue is troubling. It sure looks like baking soda use may be a surer way to add alkalinity to mashes.
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.

Perhaps we should be advising all users to do this test and the fizz test on any lime they contemplate using.
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?
 

ajdelange

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AJ, what about a solubility test to assess lime purity?
Why not? With calcium carbonate having a solubility of a little over a mg/100 mL and Ca(OH)2 being soluble to the extent of 189 mg/100 mL solubility should separate any calcium carbonate from lime.

I haven't worked out an appropriate lime dosing rate, but something like a gram or so per 100 mL of distilled water is probably sufficient.
We'd want people to keep it to about 100 mg/100 mL so the lime can dissolve.

This is not a very quantitative test, but it would serve as a qualitative test like the acid fizz test.
The acid test is good but with a solubility test all you need is water.

This whole pickling lime purity issue is troubling. It sure looks like baking soda use may be a surer way to add alkalinity to mashes.
It certainly is puzzling. I can understand how lime gets converted to CaCO3 if exposed to air (the traditional bag of lime in the dunny for example) but in a sealed jar? If it converts to CaCO3 then the calcium isn't available for crisping (that's done before the vinegar is added) so the product would be innefective for its intended purpose. Etc.
 

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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.
Tried the fizz test (using 88% lactic acid) and solubility test on my supply of pickling lime that has been stored in the original bag for ~6 months. The Mrs Wages bags do not seal well as indicated by a squeeze test. I do live in a relatively dry climate and the past 6 months have been in our dry season.

Results: No fizz, however the lactic acid was very afraid of the lime. I placed the lime on a surface and placed a couple of drops of lactic acid on the lime but the acid ran away, so I mixed them with the pipette tip, no fizz. For the solubility test: Because my best balance is a set of triple beams (0.1 gram minimum) I was leery about the accuracy so I used 1g of lime in liter of water in an E. flask. Shook for 1 minute and transferred 100 mL to a 100 mL graduated cylinder. After 5 minutes, a light dusting could be observed at the bottom of both vessels. No further precipitation occurred. Then I added another gram to the liter (2 g total) and followed the same process. This time a thin layer of precipitate was apparent.

I don't have phosphoric acid and my LHBS does not carry it. Is it possible to do the "strength test" with some dilution of 88% lactic acid?
 

ajdelange

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The Mrs Wages bags do not seal well as indicated by a squeeze test. I do live in a relatively dry climate and the past 6 months have been in our dry season.
I'm not sure moisture is required. The reaction is

CO2 + Ca(OH)2 --> CaCO3 + H2O

This is the reaction by which mortar sets. Does it continue once the mortar is dry? At a slower rate? Don't know.

Results: No fizz, however the lactic acid was very afraid of the lime. I placed the lime on a surface and placed a couple of drops of lactic acid on the lime but the acid ran away, so I mixed them with the pipette tip, no fizz.
You could try it with diluted lactic acid or vinegar.

For the solubility test: Because my best balance is a set of triple beams (0.1 gram minimum) I was leery about the accuracy so I used 1g of lime in liter of water in an E. flask. Shook for 1 minute and transferred 100 mL to a 100 mL graduated cylinder. After 5 minutes, a light dusting could be observed at the bottom of both vessels. No further precipitation occurred.
There shouldn't be any precipitation. It's a question of whether or not it all dissolves.

Then I added another gram to the liter (2 g total) and followed the same process. This time a thin layer of precipitate was apparent.
This is to be expected as the solubility limit is 1.89 g/L. In going to 2 g you have exceeded that. The 'solution' should have been cloudy ('milk of lime') and the undissolved stuff should have settled out pretty quickly). Even at 100 g/L not all the material may have dissolved as when saturation is approached it takes a lot of stirring to get everything into solution.

I don't have phosphoric acid and my LHBS does not carry it. Is it possible to do the "strength test" with some dilution of 88% lactic acid?
Yes. Perhaps the easiest approach to explaining how is to note that 100 mg of Ca(OH)2 requires 243.3 mg of lactic acid to neutralize it (to pH 7.0). If your lactic acid is 88% then you would need 243.3/0.88 = 276.5 mg of this acid which, as it has density at 88% of 1205.3 mg/mL is 276.5 / 1205.3 = 0.2294 mL. That's going to be tricky to measure out unless you have a micro pipetter. Perhaps you want to neutralized a gram of Ca(OH)2 requiring you to measure out 2.294 mL or 2.765 g. The measured out lactic acid and Ca(OH)2 can both be dissolved in convenient amounts of water.
 

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1) You could try it with diluted lactic acid or vinegar.


2) This is to be expected as the solubility limit is 1.89 g/L. In going to 2 g you have exceeded that. The 'solution' should have been cloudy ('milk of lime') and the undissolved stuff should have settled out pretty quickly). Even at 100 g/L not all the material may have dissolved as when saturation is approached it takes a lot of stirring to get everything into solution.


3)Yes. Perhaps the easiest approach to explaining how is to note that 100 mg of Ca(OH)2 requires 243.3 mg of lactic acid to neutralize it (to pH 7.0). If your lactic acid is 88% then you would need 243.3/0.88 = 276.5 mg of this acid which, as it has density at 88% of 1205.3 mg/mL is 276.5 / 1205.3 = 0.2294 mL. That's going to be tricky to measure out unless you have a micro pipetter. Perhaps you want to neutralized a gram of Ca(OH)2 requiring you to measure out 2.294 mL or 2.765 g. The measured out lactic acid and Ca(OH)2 can both be dissolved in convenient amounts of water.
1) OK

2) So, did my lime fail the solubility test? The solution was cloudy and at the 1g/L concentration the undissolved stuff was barely visible, I had to hold it to a light source and it was still hard to see.

3) I have a micro pipetter but I'm balance poor. So I'll go with 1g of Ca(OH)2 and 2.294 mL of 88% lactic acid in say 1 L of RO water to expect a pH of 7, correct?

Lastly, sorry about the numbering, I can't figure out the multi-quote thing. Can anyone advise?
 

ajdelange

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2) So, did my lime fail the solubility test? The solution was cloudy and at the 1g/L concentration the undissolved stuff was barely visible, I had to hold it to a light source and it was still hard to see.
I don't think so. If half or a quarter of the material did not dissolve that would be one thing but I am guessing that if you put it on a stir plate eventually this haze would disappear.

3) I have a micro pipetter but I'm balance poor. So I'll go with 1g of Ca(OH)2 and 2.294 mL of 88% lactic acid in say 1 L of RO water to expect a pH of 7, correct?
Yes, unless I made a math error.
 
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So having established that half of this stuff isn't either calcium hydroxide or calcium carbonate (did it fizz?) the obvious question becomes 'What is it?' and the not so obvious question is then 'Is this typical of the pickling limes we buy at the super market?'. Perhaps we should be advising all users to do this test and the fizz test on any lime they contemplate using.
AJ, sorry to keep bugging on this one, but now I'm really stumped. I bought some new lime from Amazon.com and tried the test again with almost the same exact results of just under 50% "effectiveness" of the lime. The brand is an aquarium brand instead of Mrs. Wages pickling lime, it is called Kalkwasser, and the only ingredient listed is calcium hydroxide.

Looking at the back of the bottle, it indicates that the guaranteed purity of calcium is min of 535 mg and max of 540 mg per gram. Does that equate to the 50% effectiveness of the product? If not, what the heck is going on? I don't mind having to double the lime in my water additions (after running a test mash, of course) now that I know the approximate effectiveness, but that seems like something a newbie like me would do and not a real solution. Any thoughts?
 

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Looking at the back of the bottle, it indicates that the guaranteed purity of calcium is min of 535 mg and max of 540 mg per gram. Does that equate to the 50% effectiveness of the product?
Exactly how is that worded? Were it pure slaked lime the calcium content would be 0.540 g calcium per gram of the powder. This suggests that it is indeed the calcium content they are talking about especially as no other ingredients are listed.

If not, what the heck is going on? .... Any thoughts?
The only one that comes to mind is that I made a math error. I'll check into that tomorrow.
 
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Exactly how is that worded? Were it pure slaked lime the calcium content would be 0.540 g calcium per gram of the powder. This suggests that it is indeed the calcium content they are talking about especially as no other ingredients are listed.



The only one that comes to mind is that I made a math error. I'll check into that tomorrow.
Here's a picture of the front and back of the bottle for wording.

1413676225616.jpg


1413676237857.jpg
 

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I'm following this thread with great interest....
One question though....how does one conduct a test mash?
Do you just take a little bit of your grain bill for your recipe and mash that? Or do you buy a separate batch of grains?
 
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I'm following this thread with great interest....
One question though....how does one conduct a test mash?
Do you just take a little bit of your grain bill for your recipe and mash that? Or do you buy a separate batch of grains?
You need to buy a separate batch of grains. You can mash a pound of grain, total, in the same proportions as your grain bill.
 

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Do you just take a little bit of your grain bill for your recipe and mash that? Or do you buy a separate batch of grains?
Yes, exactly. Probably the biggest thing to watch out for is that the test mash grains represent the same mixture as the full mash so that, for example, if mashing 50 lbs of base malt with 2 lbs of caramel I, two pounts of Caramel II and a pount of sauermalz one would grind say 51 lbs of the base and put the ground stuff in a separate bucket, and a bit extra of the caramels and sauermalz, again keeping them separate. He would then take 500 grams of the base malt (about a pound) 20 grams of each of the caramels and 10 of the sauermalz and mix that up for the test mash. Or 5 Oz, .2 Oz, .2 Oz and 0.1 Oz or 10 Oz, 0.4Oz, 0.4 Oz and 0.2 Oz or whatever is convenient just as long as the ratios are correct.

You could, of course, buy the grain for the test mash separately but why do that? Taxing your in house supply a bit is certainly more convenient and has the advantage that you are assured that you are testing the same lot numbers as the malt you intend to use in the main mash.
 

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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.
 

ajdelange

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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.
 

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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:
Code:
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:
Code:
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.
 

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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.

Code:
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?

Code:
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...

Code:
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

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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

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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.

:confused:
 
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CadiBrewer

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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?
 
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CadiBrewer

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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.

:confused:
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.
 
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