I’m curious about this 10% thing. Are you referring to strike water pH or mash pH? Seems to me that if, in the mash, the pH starts dropping below about 5.7, that means that all the bicarbonate (i.e. residual alkalinity) has been neutralized. If there was any bicarbonate left, it should continue to buffer and resist pH change until it is completely gone. Only then will pH drop. Thus, if your mash pH is 5.4, there is no bicarbonate left to worry about.
This is getting good. Keep going guys. 
There is a concept in Water (the Palmer and Kaminski book) that I find helpful in thinking about alkalinity (pages 67-68).
They refer to the pH of a mash conducted with distilled water as the “normal” mash pH.
Now think of a mash with source water that has some alkalinity. That alkalinity is going to push the pH up from the “normal” value. However, that is going to be counteracted to an extent by the interactions of any calcium and magnesium with phosphates in the malt which will neutralize some portion of the alkalinity (per the mechanism already described earlier in this thread). The amount of alkalinity remaining after that is the residual alkalinity.
If the residual alkalinity has a positive value, the mash pH will be higher than the “normal” value.
If the residual alkalinity has a negative value, the mash pH will be lower than the “normal” value.
I brew mostly pale beers starting with 100% distilled water to which I add some portion of calcium sulfate and calcium chloride. Usually, I need to bring the pH below the “normal” value. The negative RA that results from the calcium additions helps, but I still need to add some small portion of acid malt (alternatively, I could continue adding calcium, but I typically shoot for moderate levels of sulfate and chloride).
For example, I brewed a Belgian-style pale beer last weekend. Bru’n Water predicted a mash pH of 5.64 if conducted with distilled water (the “normal” value). My calcium sulfate and calcium chloride additions brought the predicted mash pH to 5.42. My 1 oz. acid malt addition brought it 5.33 (my actual measured mash pH was 5.34).
The nice thing about low alkalinity water when brewing pale beers is that you don’t need to add a lot of calcium, magnesium, or acid. The adjustments required are pretty minimal.
Regarding the question of whether there’s some noticeable difference between a pale beer brewed with low alkalinity water vs. high alkalinity water that has been neutralized in some flavor-neutral way, I really don’t know. I did wake up at 2:00 a.m last night and ponder the issue for about 90 minutes though.
LOL, I love that last sentence and I have been there. BEEN THERE!
So you use 100% distilled and add some CaCl and CaSO4 and that’s it. I have always wondered if a beer made that way would taste flat or thin or tasted like it was missing something. I have never done it or tasted one like that so I don’t know. Every once in awhile I take a step back and look at one piece of the process or another just to see what others are doing. I am always walking the line between “this is a hobby” and “I want to make the best beer possible” and sometimes I fall in one direction or the other. I have my water ready for tomorrow and I used 2 gallons of distilled + 5.5 gallons of filtered source water that will have 2.8g of CaCl added to it and about 4ml of 88% lactic acid to get the pH in the zip code. Let’s see what happens.
Lets say you need 30 liters of water at 70 mg/L Bicarbonate (which is 57.4 mg/L Alkalinity), and you want to reduce it to 70 mg/L Bicarb from an initial 140 mg/L.
The MW (molecular weight) of the HCO3- bicarbonate ion is 61.01684 grams per mole. The valence of HCO3- is 1. Therefore the equivalent weight (EQ) is also 61.01684 g/EQ, which is equal to 61.01684 mg/mEq (where mEq = milliequivalent).
If you have 140 mg/L bicarb and 30 L of water, then you have 140 mg/L x 30L = 4,200 mg. of HCO3- ions
To cut this in half requires that you remove 2,100 mg. of HCO3-
2,100 mg./61.01684 mg/mEq = 34.4168 mEq’s of bicarbonate to be removed.
If you choose to use 10% Phosphoric Acid, its acid strength relative to pH 6.33 (where on first educated guess removing half of the bicarb will bring your water pH) is 1.2 mEq/mL. Therefore:
34.4168 mEq/1.2 mEq/mL ~= 28.7 mL
The answer is therefore to add 28.7 mL of 10% Phosphoric Acid to every 30 Liters of your water whereby to reduce its bicarb from 140 mg/L (ppm) to 70 mg/L (ppm).
Uh. I like beer.
As to why I guessed at 6.33 being the ballpark pH for water acidified sufficiently to reduce its Bicarbonate from 140 mg/L to 70 mg/L, the standard ‘Carbonate Species’ chart seen below shows the point at which 50% of the HCO3- (bicarbonate) ion species is gone corresponding to about pH 6.33.
Silver_Is_Money: Thank you.
FWIW I’m sharing some notes from the brewday today. A very pale beer that I made with 10 lbs of grain (pilsner, some Munich and flaked corn) and 7.5 gallons of water… 2 of which were distilled. I hit my mash ph of 5.4 at room temp (which, according to another thread here translates to 5.2 at mash temp… please advise if that’s incorrect) but my sparge pH was a little low. Even though a higher percentage of the sparge water was distilled and I lowered the acid (from 2ml to 1.5), the pH was still around 5. so I need to go easier on that addition. Mid-boil pH reading was 5.3 and I added just a ½ ml of acid to lower it knowing that there was less bicarb in the water because of the distilled. I also just realized that my wife was chasing me out of the kitchen so she could cook something for an event today and I forgot to measure the post boil pH after the acid addition. I doubt it went from 5.3 to 5 but I assume it did drop. I also got a good hot break at the end of the boil and the wort was very clear going to the fermenter so I think this “2 gallons of distilled” is going to be a regular thing on most batches. Certain styles may not need it.
This is something I don’t understand.
Once dissolved, carbonate can exist in water in three different forms (aqueous carbon dioxide/carbonic acid, bicarbonate, and carbonate) depending on the pH of the water. The chart you posted shows this, and there is a similar chart on page 66 of Water. That chart also shows that below a pH of 4.3, all of the carbonate has converted to aqueous CO2/carbonic acid.
However, I don’t think that it follows to say that you cannot eliminate all of the bicarbonate in the mash (whether through reactions between phosphates in the malt with calcium and magnesium that release hydrogen protons which react with dissolved carbonate to form water and CO2 gas or through the addition of acid with supplies the necessary hydrogen ions) without bringing the pH down to 4.3.
I can prepare a mash with distilled water (that contains no bicarbonate) and have the pH settle at 5.4. Similarly, I can prepare a mash with water that has 100 ppm bicarbonate and add enough acid to bring the pH to the same 5.4. In that second scenario, I have neutralized all the alkalinity that was due to the bicarbonate content of the strike water, haven’t I eliminated all of the bicarbonate?
Or we talking about two different things here (the form in which carbonate exists in water given a particular pH vs. bicarbonate, residual alkalinity, and its effect on mash pH) or am I missing part of the larger picture?
This is quite easy to explain.
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Distilled water simply has little to no bicarbonate species to begin with. It settles in at about 5.8-5.9 pH due to the CO2 in the atmosphere reaching an equilibrium state with the water whereby entering CO2 combining with water and forming Carbonic Acid (H2CO3) is in perfect balance with Carbonic Acid dissociating back into CO2 and water. CO2 + H2O ⇆ H2CO3
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Natural water with true HCO3- Bicarbonate ions acts as I’ve previously outlined. It’s Bicarbonate is sourced primarily from extremely long term exposure to Calcium Carbonate and Magnesium Carbonate bearing rocks.
Right. This is what I originally asked about. I am not sure if I neutralize the bicarb with acid and get to a mash pH of 5.4 if the bicarb is GONE or if it’s just at a lower ppm.
In the end, I no desire or reason to build my water up. The numbers are modest enough to work for almost all styles and an easy dilution with distilled seems easy to do when necessary.
Fortunately two old (one circa 1950’s - the other circa 1960’s) peer reviewed brewing industry documents I’ve uncovered had addressed this and concluded that at even as high as 50 mg/L (ppm) Bicarbonate in ones mash and/or sparge water it is most likely harmless, and that by 25 mg/L it is 100% assuredly harmless. We have generally been led astray by some more recent and unfounded (as in amateur level and/or not peer reviewed, yet touted long enough to become gospel) notion that bicarbonate is something that must (for mainly the brewing of light colored lagers and ales) be fully eliminated. The same (generally home brewing oriented) literature generally presumes (incorrectly) that such complete elimination is achieved via acidifying ones source water to pH 5.4 - 5.6.
Since pH 5.4 more factually represents ~90% bicarbonate reduction, ones source water would need to exceed 500 mg/L bicarbonate whereby to also exceed 50 mg/L of same when acidified to pH 5.4. But where there is that much bicarbonate there are also far and away most likely going to be massive amounts of the likes of magnesium and calcium and iron and manganese, so using such water (which will generally be well water) “as is” (albeit for acidified to pH 5.4) is in most cases not an option for reasons that are other than the acid addressable high bicarbonate.
So permit me to pull us deeper down this rabbit hole. Most of the discussion regarding bicarbonate, and the equilibrium of carbonic acid/bicarbonate/carbonate with respect to pH has been looking at strike water and mash pH. But if we’re concerned with the flavor impact of bicarbonate on beer, then you also can’t overlook the fact that we boil after the mash. This drives off all of the CO2 that is dissolved as carbonic acid. And since there is an equilibrium, this leads to more bicarbonate being converted to carbonic acid, and more of that CO2 being boiled off, etc. I’m sure that the chemistry is not as straightforward in boiling wort as it is in water, but generally you end up in the 50-80ppm range for bicarbonate when you pre-boil brewing liquor IIRC. The remainder of CO2 either boils off or binds with calcium (as carbonate) and precipitates out as chalk.
If you’re looking into the flavor contribution of bicarbonate, you would want to consider the amount in the finished beer, which should be considerably lower than what you started with.
Guys, thank you again. Yes, the flavor contribution (or I could say “mouthfeel” contribution) interests me. In that Gordon Strong link, he mentioned disliking “alka-seltzer” beers (or someone else mentioned that) and I wonder what that is referring to… high bicarbonate or something else? If it’s something else then what is it? If it’s high bicarb then how high are we talking? Is he referring to a pale lager made with source water that had 400+ppm of bicarb?
Also, for the record I’ll post my numbers. They’re all modest except for bicarb so there is not a ton of Magnesium, etc. Ca 34, Mg 13, Na 12, Cl 21, SO4 27 (that’s 9x3 so maybe that’s SO4-S) and bicarb 138ppm. I have made two pale beers in the past month where I diluted with 2 gallons of distilled water and I saw better hot break and I had an easier time moving the pH with less acid. I think this is a good step and getting gallons of distilled water is easy and inexpensive. I might start doing this on every batch going forward and pay close attention to the results. Eventually I might find that it’s not necessary on all styles. Cheers and thanks again.
For readers of this thread, be aware that Gordon Strong is not detesting the flavor of bicarbonate, he’s detesting the effects of bicarbonate in beer. Bicarbonate is essentially flavorless in water. Now let’s move on to the real issues.
Having a modest bicarbonate content is okay for most brewing. I had a level similar to Ken’s in my Tallahassee water supply. I never detected lactate in my resulting beer that I made there. But I have a typical taste threshold for lactate. There are some people that are supertasters with respect to lactate and they might have had another opinion. That level is getting up there.
If YOU feel that you can taste lactate when employing lactic acid in any of its forms, you should limit the lactic addition and employ another acid to supplement or replace. Using other acid is okay, but I find that using some form of lactic acid is helpful in creating authentic tasting Continental European beer styles. British and American styles can be made without lactic acid.
I feel like this question is still outstanding, and also there has been some divergence in the discussion between (i) acidifying strike water and (ii) residual alkalinity/mash pH.
If you reduce the mash pH to at or below the pH that would have resulted if you had used 100% distilled water with no bicarbonate as your strike water, what is the mechanism by which bicarbonate can still be present? Put another way, if you have reduced residual alkalinity to zero (or below), have you converted all the bicarbonate to water and CO2?
I get that it doesn’t really matter for all practical brewing purposes, but I want to understand it just for the sake of understanding it.
Martin, thank you for joining and replying. I do not taste the acid, AFAIK.
And since Narcout likes divergence, please allow me to add one last thought. Imagine something pilsnery or maybe an amber lager (SRM 8-9 ish) where I have added 2-3 ounces of something noble like Hallertau, Tettnanger, Spalt Select, etc. and I get a little roughness, harshness in the finish. Is this related to the rest of this topic? Are those noble hops reacting to leftover bicarb, my 27ppm of sulfate, etc? That piece might not even be closely related but my untrained and unscientific mind wonders if there is a connection and I also seem to remember a story (probably unsubstantiated) about noble hops and either sulfate or bicarb clashing. Cheers again.
The pH scale for the carbonate/bicarbonate/carbonic acid equilibrium are:
4.3 - below this pH, everything is in the form of Carbonic acid/dissolved CO2. There are no bicarbonate or carbonate ions in solution
4.3-8.3 - everything exists as a mixture of carbonic acid and bicarbonate, with the proportion of bicarbonate increasing as pH increases
6.3 - this is a buffer with equal concentrations of carbonic acid and bicarbonate
8.3 - at this value everything exists as bicarbonate
8.3-12.3 - everything exists as a mixture of bicarbonate and carbonate, with the proportion of carbonate increasing as pH increases
10.3 - this is a buffer with equal concentrations of bicarbonate and carbonate
12.3 - above this pH, everything is in the form of carbonate, no bicarbonate or carbonic acid is present
Note that these values are directly related to the pKa values for H2CO3 <> H+ + HCO3-; and HCO3- <> H+ + CO3(2-). This isn’t related to the pH of distilled water. Also, the presence of ions like calcium and mag make this a bit less straightforward, as they can bind to carbonate and bicarbonate ions, sequestering them from this equilibrium.
I get the chart, what I don’t get is the how (or if) bicarbonate can still be present in the mash if you have reduced residual alkalinity to zero (which you can do in the normal mash pH range and without reducing the pH of the mash all the way to 4.3).
RA (ppm as CaCO3) = Alkalinity (ppm as CaCO3) - [(Ca (ppm)/1.4) + (Mg (ppm)/1.7]
If RA is zero, how is bicarbonate still present?