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A Sparge Water acidification experiment gone awry

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Larry Sayre, Developer of 'Mash Made Easy'
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I figured that my abysmal well water with its 377 ppm of Alkalinity (as CaCO3) at a pH of 7.7 (all per Ward Lab) would make a good test material for checking the accuracy of sparge adjustment as per Mash Made Easy 9.55 and the Brewers Friend online mash calculator. So I gathered and divided some of my well water into 2 x 1 Quart samples, and let them both come up to room temperature over a period of several hours.

Next, when the water was at a measured 69 degrees F. I added 0.6 mL of 88% Lactic Acid to one quart of the well water, and 0.70 grams of Citric Acid crystals to the other quart of well water. I mixed both well and let them sit again. In the mean time I proceeded to check the software pH predictions and calibrate my 'Oakton pH5' pH meter with fresh buffer made from 7.0 and 4.0 pH buffering tablets and distilled water, plus temperature adjust it for the 69 degree water:

Here are the software pH predictions:

Brewers Friend said that:
1) 0.60 mL of 88% Lactic acid would bring my quart of well water to pH 5.10
2) 0.70 grams of citric acid would bring my quart of well water to pH 5.43

Mash Made Easy said that:
1) 0.60 mL of 88% Lactic acid would bring my quart of well water to pH 5.18
2) 0.70 grams of citric acid would bring my quart of well water to pH 5.27

Next I read the actual metered pH for each quart of acidified well water as follows:
1) 0.60 mL of 88% Lactic acid actually brought my quart of well water to pH 5.52
2) 0.70 grams of citric acid actually brought my other quart of well water to pH 4.90

Conclusion:
1) For Lactic Acid: Assigning it a revised percentage concentration of 79% to 80% brings it right close to the measured 5.52 pH
2) I have no real explanation at all for the citric acid. Brewers Friend said that 0.91 grams would be required to achieve pH 4.90, and MME said that 0.84 grams would be required for pH 4.90. Citric Acid crystals must be at least ballpark 20% or more stronger in mEq/gram acidity than predicted by either software.
 
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Assigning a pKa3 value of 5.19 or 5.20 to Citric Acid explains it perfectly. Most online sources say the pKa3 for Citric Acid is between 6.38 and 6.40, but I did find one source which lists it at 5.19, one which lists it at 5.20, and a couple which list it at 5.40.

5.19, 5.20, and even 5.40 are far closer to what I measured than 6.38 to 6.40 for Citric Acids pKa3 value. Strange????
 
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In case you're wondering, both my 88% Lactic Acid and my Citric Acid crystals are LD Carlson brand.

When I adjusted a testing version of MME to a pKa3 value of 5.19 for Citric Acid it predicted 4.90 pH for the addition of 0.70 grams into 1 quart of my 377 ppm Alkalinity and 7.70 pH well water.

I wonder if the 6.38 to 6.4 values for pKa3 seen for such sources as Wikipedia (and others) are measured for the mono-hydrate version of Citric Acid, as opposed to anhydrous? A good while ago I asked this forum if LD Carlson's Citric Acid crystal was anhydrous or the mono-hydrate state, and I never received an answer. My test today confirms to me that it is the anhydrous form. Only the anhydrous form with an assigned pKa3 of 5.19 or 5.20 matches what I witnessed as to its mEq strength per gram.
 
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I just resurrected my Apera pH60 meter from the dead. It turns out that the liquid buffers I was using with it (being a few years old now) were both shot, even when poured from the uncontaminated bottles. When I made up fresh buffers from buffer tablets it calibrated properly and then read 5.54 pH for the quart of well water that had received the Lactic Acid, and 4.91 pH for the quart which had received the Citric Acid. I'm considering this a decent double-check, and now I'm glad that I have two working pH meters.

From now on I will only purchase tablet form buffers.
 
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I'm planning a new test for my 377 ppm Alkalinity (as CaCO3) and pH 7.7 well water as follows:

1) If pKa3 for anhydrous Citric Acid is 5.20, then 0.52 grams of Citric Acid should take 1 quart of my well water to pH 5.38
2) If my current bottle of 88% Lactic Acid is actually 80%, then 0.7 mL of it should take 1 quart of my well water to pH 5.00

Stay tuned for the dual pH metered results for each.
 
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Here are the actual results:

Apera pH60: Citric Acid sample: pH = 5.33, Lactic Acid sample: pH = 5.29
Oakton pH5: Citric Acid sample: pH = 5.37, Lactic Acid sample: pH = 5.33

Vs. Predicted: Citric Acid = 5.38 (for the case of pKa3 set to 5.20
Vs. Predicted: Lactic Acid = 5.00 (for the case of presuming 80% by weight Lactic Acid)

A whopping success for the Citric Acid. Validates pKa3 dissociation = 5.19 or 5.20 (and not 6.40 as I had originally presumed)
An utter quandary for the Lactic Acid. I mixed it well this time before drawing the 0.7 mL, and now it's even weaker than before.
 
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I suddenly recall reading a relatively recent peer reviewed document in which, try as they may to match Lactic Acid to their mash acidifying math model, Lactic Acid simply would not lower the pH to their predicted target. This is the very same thing I'm observing, sans that I'm acidifying highly alkaline well water. They attributed it to Lactic Acid tossing into the mash a whopper load of its own buffering. But I'm following the AJ model for matching Lactic Acids relative strength to its targeted pH. It appears the pKa issue is biting me here. I have to see how many pKa's Lactic Acid actually has, and what values they have, and see if that corrects the AJ model. Hmmm????

Now I've got to search for that document and read it again. Stay tuned. It may take me awhile to dig this document out of the internet again.
 
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I found it, and as luck has it, it is at the quoted juncture discussing water decarbonation via acids. About Lactic Acid it says this on page 4 left column, about 2/3 down the page:

However lactic acid cannot remove the carbonate hardness completely. This can be traced back to the fact that lactic acid, as a weak acid, cannot be dissociated completely and the carbonates cannot be converted entirely.
 
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Now I'm in a complete funk over this entire matter. I had always presumed that the alkalinity (as CaCO3)/buffering capacity for my homes softened water is one and the same as for my well water itself, so as a shortcut I drew all of the samples used for the tests as seen above from my kitchen sink tap, which is delivered water that has passed through our water softener, believing it makes no difference. But:

I drew 2 quarts directly from the well proper and repeated my 0.52 grams of Citric Acid and 0.7 mL of Lactic Acid added to one quart of well water experiments, and this time, instead of yielding results (confirmed via dual pH meters) of ~5.35 (average) for Citric Acid and ~5.31 pH (average) for Lactic acid (as per post #6 above), both samples measured right close to pH 4.60 (average, Lactic) and pH 4.62 (average, Citric) this time around (also confirmed via dual pH meters). To say the least, I'm totally baffled now. So I've started a thread to see if there is any alkalinity/buffering capacity difference for softener softened vs. straight well water.
 
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I'm next going to draw 4 x 1 Quart water samples direct from the well and add the following to them and then measure their pH in duplicate of everything I've done from softened tap water:

0.52 g. Citric Acid
0.70 g. Citric Acid
0.6 mL Lactic Acid
0.7 mL Lactic Acid

I actually have two ways to draw from my well. Within the same well casing I have a submerged electrical pump drawing at 70 ft. depth (which is our house feed), and a hand pump drawing at 50 ft depth (which provides water when we experience a power failure). The pH 4.6'ish samples came from the hand pump. This time I'm heading for the basement to draw 4 samples pre the softener from delivery via the electric pump. Perhaps there is even a well stratification difference in the Alkalinity whereby drawing from 50 ft. and drawing from 70 ft brings up different water?
 
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It turns out that Lactic Acid has 2 dissociation pKa's, with the first at 3.86 and the second at 15.1. But the 15.1 pKa2 does not alter Lactic Acid's mEq/mL acid strength sufficiently to be noticed out to about 4 decimal places or perhaps even more for the pH targets of importance to brewing, so whether or not one incorporates the 15.1 valued pKa2, it does not seem to alter things.
 

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I may be off base (and tell me if I am), but how well do you know the purity of your reagents? For the Citric Acid, you have mentioned that it will form a mono-hydrate. Even if you have the anhydrous form, just from humidity it will hydrate (just like CaCl2) and my understanding is to do precise work, you need to dry it out immediately prior to use. I think the common procedure it to bake in a 200-250F oven a few hours. The water is driven off at around 175F and the Citric Acid doesn't decompose until around 350F. The molecular weight of the mono-hydrate is about 10% greater than the anhydrous. The question would be whether or not the calculations used are based on the mono-hydrate or anhydrous forms. While you believe you are dealing with the anhydrous form, if it has been exposed to the air, you know that some of it has hydrated just from the humidity.

Perhaps borrowing a page from AJs book on Calcium Chloride is a simpler solution. Per the CRC Handbook, a 30% by weight solution of Citric Acid has a Refractive Index of 1.3744 at 20C. So, dissolving 30 grams of your Citric Acid in enough water to make 100 grams should give you a reading of 26.2 Brix with a Brix refractometer at 20C. If you have the mono-hydrate, it should read 23.6 Birx. Somewhere in between and you have a mix of anhydrous and hydrate forms. Anyway, adjust the concentration until you get it to read 26.2 Brix with the refractometer and you should be at 30% Citric Acid, then just make volume additions. Per the CRC, 30% Citric Acid has a density of 1.1346 g/ml (although I believe that is at 4C, so room temp density will be a tad less).

As far as the Lactic Acid is concerned, my guess is the LD Carlson merely repackages food grade Lactic Acid making no dilutions. At 88%, the refractive index is too high to be able to read on a typical Brix refractometer. However, a 44% Lactic Acid solution should produce a reading of 28.6 Brix with a Brix Refractometer at 20C, so a 50:50 mix of the Lactic Acid and Distilled water should produce a readable concentration, using the typical Brix Refractometer used by homebrewers.
 
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I may be off base (and tell me if I am), but how well do you know the purity of your reagents?
Not well at all. But if my Citric Acid is the mono-hydrate, then the mono-hydrate it is a stronger acid than if it were anhydrous. And somehow I doubt that.
 
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As per #10 above, I did pull 4 new samples, acidify them as per #10, and then let them come up to ~20 degrees C. (20.1 measured). These 4 quarts were pulled in my basement from a tap that precedes both my oxidizer and softener. So they are 70 ft. deep well samples.

While calibrating the meters, I could not bring my Oakton to calibration, try as I may. But the old Apera pH60 calibrated just fine and double checked in the buffers afterward just fine also. So I only used the Apera. Here is what I found for it:

1) For 0.52 grams of Citric Acid, the pH reading was 4.56 (in relatively good agreement with the post #9 result).
2) For 0.70 grams of Citric Acid the pH reading was 3.98
3) For 0.6 mL of Lactic Acid the pH reading was 5.39
4) For 0.7 mL of Lactic Acid the pH reading was 5.18 (not in very good agreement with the post #9 result)

I'm going to surmise from this exercise that:

1) For sure you can't take a softened water sample and presume it will have similar Alkalinity and/or buffering as for a direct well drawn sample. Replacing Ca++ and Mg++ ions with 2 x Na+ ions per each must really toss a monkey wrench into the buffering works.
2) Citric Acid is way more acidic than I could have ever previously imagined. Somewhere around 13.54 mEq/gram will perhaps suffice at the pH's of interest to brewing (5.4 +/- a tad). It could even be stronger than this.
3) My 88% lactic Acid is likely somewhere around 79-80%.
 
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My next contemplated experiment along these lines will be to add Baking soda to 4 quarts of distilled water whereby to achieve ~377 ppm worth of Alkalinity, and then acidify much as was done above in order to see what differences this derivation brings to the table. My thinking is that this may come closer to what was observed for my tests which utilized softened well water, in that the alkalinity will be accompanied with a load of sodium.
 
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OK, time to calculate how much baking soda (NaHCO3) to add whereby to achieve 377 ppm Alkalinity (as CaCO3) in 1 Quart of distilled water. First assuming full dissociation, and then assuming partial dissociation at pH 5.40. Two parts are presented below, followed by a refutation and a conclusion.

Some givens we may find useful:
MW NaHCO3 = 84.00661
MW Na+ = 22.98977
MW HCO3- = 61.01684
MW CaCO3 = 100.0869
EQW CaCO3 = MW/2 = 100.0869/2 = 50.04345

Part 1: The case for full dissociation

Weight of 1 Qt pure DI water = 946.35 Grams at STP
(but weight of water plus NaHCO3 will clearly be a tad more, so let's permit a bit of a jump ahead 'cheat' here and call it 947 Grams)

377 grams per million grams = 377/1,000,000 = 377 ppm

377/1,000,000 x 947 grams = 0.357019 grams of CaCO3 required for "377 ppm as CaCO3"

84.00661/50.04345 x 0.357019 grams = 0.59932 Grams of NaHCO3 (lets round this to 0.60 Grams)

Answer = 0.60 Grams of Baking Soda added to 1 Quart of distilled water = 377 ppm of Alkalinity as CaCO3

----------------------------------------------------------------------------------------------------------------------------------------------

But now, per AJ, we are faced with rather poor dissociation issues for Baking Soda:

Part 2: The case for partial dissociation

Baking Soda delivers 11.9038 mmoles per added gram and since its valence is 1 this = 11.9038 mEq/gram for 100% dissociation

But at a pH of 5.4 it only partially dissociates (ionizes into Na+ and HCO3-), so some therefore must remain behind as NaHCO3

The partial dissociation for NaHCO3 at specifically pH 5.4 is such that only 10.7053 mEq/Gram effective mEq strength is realized

Therefore:

Per AJ (presuming that I'm interpreting him correctly, which I obviously may not be), one should add 11.9038/10.7053 x 0.60 Grams = 0.667 Grams of NaHCO3 (lets call it 0.67 Grams)

Answer when factoring in dissociation = 0.67 grams of Baking Soda must be added to better equate it to 377 ppm as CaCO3

NOTE of Refutation: I highly dispute this on the basis that when acid is added to bring the solution to pH 5.40 and some CO2 gas is thereby liberated, whereby some CO2 also obviously (temperature dependent) remains in solution, but some also evolves out as a gas and escapes, so clearly in reality the chemical reaction between the acid and Baking Soda can not fully travel in each direction as AJ contends, which would only be possible if zero CO2 ever escaped the solution as a gas and thereby vanished. The real answer at room temperature clearly therefore must lie somewhere between 0.60 grams and 0.67 grams of baking soda to be added. And at mash temperature the CO2 remaining in solution is going to be far less, as more is driven off as a gas and is never to return to the solution whereby to reverse the reactions direction and thereby partialize the dissociation. So at mash temperature the dissociation problem will be far less of a problem.

Conclusion: Based upon all of the above I don't know how much NaHCO3 to actually add to achieve a perfect 377 ppm as CaCO3 for which to titrate at room temperature, but I know it will lie somewhere between the extremes of 0.60 grams and 0.67 grams, so I will simply guess and add 0.63 grams. My baking Soda is not likely 100% pure anyway, so in reality this means I will be actually adding a smidge less than 0.63 grams.
 
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Here's what I'm currently in the process of making up into 1 volume measured quart of distilled water per each for the next round of testing:

Sample 1: 0.63 g. Baking Soda, & 0.47 g. Citric Acid
Sample 2: 0.63 g. Baking Soda, & 0.47 g. Citric Acid
Sample 3: 0.63 g. Baking Soda, & 0.6 mL Lactic Acid
Sample 4: 0.63 g. Baking Soda, & 0.6 mL Lactic Acid

Since one of my pH meters is 'potentially' on the fritz (whereby I will make an effort to use it again if it will confidently calibrate and double check, or pass it by if it doesn't), I've decided to do duplicates of the sample make up, which will also test my rather crude equipments ability to deliver the quantities as indicated above. I'm using a tiny budget gram scale which reads out to 2 decimal places for the citric acid, and a 1 mL medicine dispensing syringe for the Lactic Acid.
 
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Here are the results:

Once again my Oakton proved to be wonky but my Apera held tight to calibration on recheck after sample testing each time. Results for the Apera were read in triplicate, and are as follows:

Sample 1): pH avg. = 5.53 +/- 0.03 extreme variation
Sample 2): pH avg. = 5.63 +/- 0.02 extreme variation
Sample 3): pH avg. = 5.56 +/- 0.02 extreme variation
Sample 4): pH avg. = 5.57 +/- 0.01 extreme variation

Averaging between samples yields:

0.47 g. Citric Acid in 1 Qt of 'nominally' ~377 ppm alkalinity (via 0.63 g. Baking Soda) = 5.58 pH average
0.6 mL Lactic Acid in 1 Qt. of 'nominally' ~377 ppm Alkalinity (via 0.63 g. Baking Soda) = 5.565 pH average

Once again this seems to confirm that my 88% Lactic Acid is actually closer to 80% Lactic Acid. Since the pH differences between Citric and Lactic samples are so close, if I accept that at 5.57 pH my 80% Lactic Acid's strength is thereby 10.3404 mEq/mL, then:

0.6 mL / 0.47 g. x 10.3404 mEq/mL ~= 13.2 mEq/Gram for Citric Acid (which is in fair general agreement with my earlier ~13.54 estimate)

I'm going to assume henceforth that a decent mEq/Gram value for Citric Acid is going to be somewhere around ~13.2 on the low end to likely not higher than ~14'ish at the high end (see below) within the ballpark vicinity of typical mash pH targets. I'm going to tentatively therefore call it ~13.6 mEq/Gram, since my 0.63 grams of Baking Soda = 377 ppm Alkalinity is an admitted guess, and if AJ is correct and it indeed requires 0.67 grams to hit 377 ppm Alkalinity, then:

0.67/0.63 x 13.2 mEq/Gram calculated = 14.04 mEq/Gram for Citric Acid at ballpark mash pH target(s)

YMMV, so I await independent confirmation.
 
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Interestingly enough, with 'Mash Made Easy' set to 13.6 mEq/Gram for citric acid, when 377 ppm alkalinity and pH 7.7 are input into the sparge water calculator, and the sparge water pH target is set to 4.56 pH (in direct correspondence with post #14 above) the predicted Ciric Acid requirement that the software now delivers is 0.52 Grams, which is a bingo hit match to the findings for Citric Acid in post #14 using my actual well water. A very good check as I see it in confirmation of 13.60 mEq/Gram acid strength being right close to the truth for citric acid.

So in the end it was not the experiment(s) that went awry, it was rather my three early "false" presumptions that went awry:
1) Softened water and well water share the same Alkalinity and buffering characteristics, which I now see as a complete falsehood. And for which all of the titrating I did using softened tap water above must be discarded, and only my actual well water and the simulation of it using baking soda are to be trusted.
2) Citric Acid was initially presumed to be way less acidic that it actually proved out to be.
3) A presumption that citric acid strength could be reliably modeled off of its 3 x pKa values. I now believe it can not be so modeled.

The initial testing was in the end not awry, but very informative. As was the latter testing which set things straight.
 
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I just checked the 'Five Star' website for Citric Acid, and in the stability section their SDS says it is Hygroscopic, whereas in the same location the SDS for the LD Carlson supplied Citric Acid indicates Stable. I'm going to have to lean toward it being hygroscopic over time. That means it will eventually reach the monohydrate state.
 
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Brief summary: When I rigorously attempted to apply dissociation constants to citric acid whereby to model its acid strength, the computed quantities to be added consistently led me to measured results such as pH 4.90, pH 4.60, and pH 3.98 when I was shooting for higher and more practical and applicable to brewing pH values. When on the other hand I finally assumed a fixed acid strength for citric acid the far more simplistic and straight forward math model predictions of this approach lead to pH's that were within the realm of usefulness to brewing as well as being right close to actual measured observation. This is indeed a quandary (I.E., the source for my funk). But in the end the task of science is to provide models (math models) that are a better fit to actual applied scientific measurement. So by this it has become clear to me that the best math model is the one that will usefully predict an outcome that will be verified in the real world of brewing by actual testing. And if a theoretical math model does not even come close to predicting real world verifiable results one must abandon it as being flawed for the application. At least until such time as an explanation for such deviation can be introduced whereby to rectify it. So for now, for me, it's out with dissociation constants for citric acid as applied to beer brewing.

The question now becomes: What will be the 'applied science' measured validity of utilizing the complexity of dissociation constants (Ka's, or pKa's) with regard to brewing for other (and more common) pH altering reagents such as Lactic Acid, Phosphoric Acid, and Baking Soda? Will there be yet more quandary/funk?
 
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One way to logically approach what on the surface appears to be the illogical, leading to my funk (I.E., that predictions based upon dissociation constants for Citric Acid are somehow '"partially" broken at the local level of my actual testing) is to look at other entrenched laws of science that are routinely locally broken as well. The most classic example here involves entropy. On earth things tend to often violate this inviolable law on a rather routine basis, and we must be glad that they do, else we would not exist. Things can indeed go from a state of less order to a state of more order on our earth, seemingly in direct violation of the law of entropy. Does this change the law of entropy. Of course not. The reason for this case is that the energy from the sun alters the local system. I have already explained how the exiting of CO2 gas from the local system of the mash (or any other local environment wherein acid is introduced) "partially" breaks the otherwise unbreakable rule of dissociation constants based math modeling firmly dictating pH outcomes for NaHCO3 (Baking Soda), rendering a strict application of the science of NaHCO3's dissociation constants to some degree invalid to brewing. All that needs to be determined within local systems for the addition of Citric Acid whereby to similarly "partially" break the inviolable law of dissociation is to discover the mechanism extent whereby to alter the local system from its requisite closed state, open it thereby, and thus logically violate that which at first appears inviolable. Any 'open system' mechanism which would disrupt the free reversal of the reaction of Citric Acid with any/all basic substances (which is dictated by its dissociation constants) will (and in fact must) suffice to permit such violation of the math model to occur. It may yet again simply be related to the escape of CO2 from the local system as a gas, or it may be some other mechanism, or a combination.

One means to test this might be to highly pressurize a system in which the HCO3- bicarbonate ion is reacted with Citric (or any other) Acid, and see if while under pressure (which keeps more CO2 in solution, permitting reaction reversal) the pH measures higher than when pressure is later reduced (permitting the liberation of CO2 as a gas from the system and altering the dissociation equilibrium).

If CO2 is indeed the "partial" dissociation constant breaker (to some measure) for Citric Acid, it logically should prove to be (to a differing "partial" extent) the dissociation constant math model breaker for any other weak acid one may choose, including Lactic or Phosphoric.
 
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Another mechanisn whereby to disrupt theoretical pH based outcomes based upon dissociation constants is to form within a chemical reaction a component (or components) which is (are) highly to completely insoluble. Once such a reactant product is formed, being insoluble, it drops out of solution, thereby forbidding the reaction to run in reverse, and promoting its forward only motion despite the presence of pKa's desire to reverse the reaction.

So now we have both levels of evolution (as a gas) and levels of insolubility as "partial" dissociation disruptors.

Any such disruption mechanism will drive the reaction of a weak acid further along than predicted without factoring in such disruption, thereby driving the pH lower and effectively increasing the acids strength beyond that predicted via the strict adherence to pKs's (dissociation constants).

Indeed it is a complex environment within the sparge water, the mash, boil, fermentation, etc... for beer brewing, and it is naive to:
1) Ignore dissociation and compute resultant pH via full theoretical normality based acid strength
2) Ignore dissociation disruptors which "partially" to some unique measure drive acid strength above pKa based prediction
(it is also a changing or evolving environment, and not a static environment)

And then lastly there are buffering compounds potentially being formed....
 
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After a lengthy search I finally came across this chart:

Citric_Quantity.png

Per the very bottom line of the above chart: 0.26 grams of Citric acid when added to 1 gallon neutralizes 1 mEq/L of Alkalinity

I was testing 1 quart at a time, so 0.26/4 Qts/Gal = 0.065 grams of Citric acid added to 1 Quart neutralizes 1 mEq/L of Alkalinity

377 mg/L (ppm) Alkalinity / 50.04345 mg/mEq of CaCO3 = 7.533453mEq/L Alkalinity for my well water
(and also for my Baking Soda and distilled water make-up equivalent of same)

1 Qt. = 0.94635 Liters

0.94635 L/Qt. x 7.533453 mEq/L = 7.1293 mEq/Qt. of Alkilinity in 1 Quart of my well water

7.1293 mEq/Qt. x 0.65 grams/mEq/Qt = 0.463405 grams of citric acid required to hit pH 5.8

I added 0.47 grams per Qt. and I hit an average of 5.58 pH thereby (so 0.01 grams more added for 5.58 pH vs. 5.8 pH)

Seems like a pretty close match, and I feel this strongly vindicates my acid strength estimation for Citric Acid.
 

jseyfert3

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These 4 quarts were pulled in my basement from a tap that precedes both my oxidizer and softener.
I'm going to surmise from this exercise that:

1) For sure you can't take a softened water sample and presume it will have similar Alkalinity and/or buffering as for a direct well drawn sample.
So a) I'm not a chemistry guy and b) I'm a complete beer noobie but I am currently stuck in bed, and ran across this. I see you are assuming softening won't change alkalinity, but you don't have just a softener, you have an oxidizer too. What does this do? I did a quick Google search and found something I seem to remember off This Old House, it puts a bunch of air across the water (aeration)?

While I said I wasn't a chemistry guy, I do have a pool and more importantly a hot tub. I forget the exact reason, but aeration of water will raise the pH, the higher the total alkalinity the faster the pH will rise for a given amount of aeration, up until the point at which the pH reaches equilibrium with the atmosphere, the exact pH this equilibrium happens at is dependent on the TA level. This has something to do with the carbonates in the water and atmospheric CO2. Not sure if the TA changes by aeration but I would suppose it does.

Where this comes in with hot tubs is they have lots of aeration, our water is roughly 350 ppm TA (as CaCO3), so the pH rises extremely quick. To keep this in the mid to upper 7's, acid is needed. So typically on a fresh fill I will cycle aeration to raise pH, then acid to lower pH, and continue until this cycle of aeration and acid addition has lowered the TA to something around 50-60. Around this point the pH rise due to aeration has now been drastically reduced, so acid additions are not needed nearly so often.

Anyhow, random thoughts that may or may not be helpful or accurate.
 
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I never thought of the impact of the oxidizer, but mine does not have an acid stage so it could be the reason behind the rise in alkalinity. To be honest, I simply don't know. All I know is that from now on when I want to run tests based upon the Alkalinity of my well water (as reported by Ward Labs) I will use well water and not oxidizer plus softener modified well water.

I can however state that I've measured both for pH, and my tap and well water are pretty close in pH. There are several types of oxidizers working upon differing principles. If one raises the pH while performing its function, it may be of a differing nature than mine.

That plus the issue whereby water pH and water Alkalinity do not correlate very well.
 
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As to the presumed flavor impact of Citric Acid in beer, I recall that forum moderator @passedpawn has made several 10 gallon batches into which he has added 10 grams of Citric Acid during the mash, and as I recall, he reports that he has never detected it's presence. Perhaps he will drop in to confirm this, and to let us know if he is still using Citric Acid to reduce mash pH.
 
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As to the presumed flavor impact of Citric Acid in beer, I recall that forum moderator @passedpawn has made several 10 gallon batches into which he has added 10 grams of Citric Acid during the mash, and as I recall, he reports that he has never detected it's presence. Perhaps he will drop in to confirm this, and to let us know if he is still using Citric Acid to reduce mash pH.
That is correct! I did not continue as I get pretty bored with water chemistry efforts. But there was a time during which I was acidifying as you described. Never detected it in the finished beer. I know what it tastes like :) I also clean my coffee pot with citric acid, every 100 batches, and I have to send 3 additional full batches of water through it before the residual citric acid is diluted below detection threshold.
 
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