A Brewing Water Chemistry Primer

Yes, 5.4 to 5.5 at room temperature is, IMO, ideal. But you'll see lots of other ranges quoted as being optimum. Unfortunately, you often can't tell what temperature they are talking about. DeClerck states that all his pH's are for room temperature. Other authors omit this statement but I think it is safe to assume that most measurements are at room temperature.

ajdelange said:

Yes, 5.4 to 5.5 at room temperature is, IMO, ideal. But you'll see lots of other ranges quoted as being optimum. Unfortunately, you often can't tell what temperature they are talking about. DeClerck states that all his pH's are for room temperature. Other authors omit this statement but I think it is safe to assume that most measurements are at room temperature.

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Thanks for the clarification.

ajdelange said:

Now if it is a lager that is coming in at 4.2 we might scratch our heads a bit...

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Latest 4.2 post-ferm pH was a Vienna lager using White Labs WLP833 German Bock yeast. It's only ~2 weeks old so it'll be a while before I get around to drinking it. Brewed a Boh Pils with the cake so I'll have another measurement in a couple of weeks (but no taste for a couple of months).

I might 'need' that 4.2 for the Boh Pils though...found a friggin fruit fly in the wort just as I started to rack it! Could have happened during cooling. Luckily it had krausen about 6 hours after pitching (cold) so maybe it'll be OK.

4.2 is low for Vienna - all mine come in around 4.6 so I'm a little puzzled. Meter was calibrated day of measurement? Buffers were fresh? Smells clean and tastes OK (no infections)? What yeast strain?

If a sample tastes clean, then I wouldn't worry. You did everything right and pH's fell in place along the way. The beer should be fine.

ajdelange said:

4.2 is low for Vienna - all mine come in around 5.6 so I'm a little puzzled. Meter was calibrated day of measurement? Buffers were fresh? Smells clean and tastes OK (no infections)? What yeast strain?

If a sample tastes clean, then I wouldn't worry. You did everything right and pH's fell in place along the way. The beer should be fine.

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I assume you mean 4.6?

Yeast was WLP833 German Bock yeast. Meter was cal'd day before and FWIW, when my meter has gone out of cal it has gone high every time. Beer tasted yeasty but OK otherwise.

Yes.

Is there any reason I shouldn't use 5.2 Stabilizer to control pH?

emjay said:

Is there any reason I shouldn't use 5.2 Stabilizer to control pH?

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If you reword the question slightly to "Is there any reason I can't use 5.2 Stabilizer to control pH?" then the answer is yes, there are a couple. They have been set forth so often here that I'll only give the top level answer that the product does not work in most cases (no one with a pH meter on this or any other forum I'm aware of has seen it do what we expect it to do) and it adds a lot of sodium to the brew.

ajdelange said:

If you reword the question slightly to "Is there any reason I can't use 5.2 Stabilizer to control pH?" then the answer is yes, there are a couple. They have been set forth so often here that I'll only give the top level answer that the product does not work in most cases (no one with a pH meter on this or any other forum I'm aware of has seen it do what we expect it to do) and it adds a lot of sodium to the brew.

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Ah, I apologize if it's been asked so much. My PC is dead ATM and I'm posting on my cell phone; my Android browser really doesn't like the search function on this site - it closes the pull-down as soon as I go to type in the text field. I promise I'm not just being obnoxiously lazy

Thanks for the reply though.

Water can ruin your day in 2 ways:
1. It can cause mash pH to be too high. This usually results in dull flavors - an insipid beer.
2. There can be something in the water which causes an off flavor to appear
a)Chloramine - bandaid, plastic,smokey taste
b)Geosmines etc. - musty taste
c)Excess of chloride paired with sodium/potassium - salty taste
d)High content of some metallic ion such as copper, iron or zinc - metallic taste
e)High sulfate content - harsh hops bitterness.
f)Something else I don't know about or have forgotten about.

As he hit mash pH the problem isn't 1) and must be one of the items in 2.

ajdelange said:

Water can ruin your day in 2 ways:
1. It can cause mash pH to be too high. This usually results in dull flavors - an insipid beer.
2. There can be something in the water which causes an off flavor to appear
a)Chloramine - bandaid, plastic,smokey taste
b)Geosmines etc. - musty taste
c)Excess of chloride paired with sodium/potassium - salty taste
d)High content of some metallic ion such as copper, iron or zinc - metallic taste
e)High sulfate content - harsh hops bitterness.
f)Something else I don't know about or have forgotten about.

As he hit mash pH the problem isn't 1) and must be one of the items in 2.

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AJ, I meant to ask you before- at what point does lactic acid or sauermalz become discernible? 3% for the malt? What about for the lactic acid?

I know my sparge water is too alkaline, that's why I ask. Not so much for mash pH, but for sparging. Can a wee bit of lactic acid be added for the sparge to keep the pH under control during the sparge?

Sauermalz is discernible at 3% but not for sourness but rather a subtle complexity. IMO that is one of the advantages to using it, at least in lagers. It really adds something to the taste. I'm certain one could go higher before sourness was detectable but I'm not sure how much higher. I expect it depends on the beer and the taster. As Weyermann recommends 8% (IIRC) for a pseudo Berliner Weiße I think we're safe in assuming that's too much (unless trying to do a Weiße). I certainly think you are OK adding a bit of lactic acid to the sparge water for pH control.

Pardon me for being a moron, but...

Is calcium chloride dihydrate the same as calcium chloride?

CaCl the same as CaCl2?

If really delving down into the nitty-gritty of water chemistry one has to consider the formation of ion pairs such as CaCl+ which come about because a solution of calcium chloride has mobile doubly charged calcium ions and mobile singly charges chloride ions loose and moving about. The positive charge on a calcium ion will attract a negatively charged chloride ion forming the pair but just as the bonds between Ca++ and the pair of chloride ions are easily disrupted during solvation with water so would a single bond with a single chloride ion and the pair wouldn't be long lived.

If you filled a container with calcium vapor and admitted 1 chlorine atom I suppose 1 molecule of CaCl would be formed as a chlorine radical will be pretty grabby for an electron from the first calcium it encounters.

CaClF exists though.

Guess I should have stated the reason for asking...

The prescribed formula lists "calcium chloride dihydrate (what your beer store sells)" Online beer stores offer calcium chloride. Just want to make sure we are talking about the same thing.

Googling "calcium chloride dihydrate" gave results listing CaCl2. Online vendors don't offer any designation.

arturo7 said:

The prescribed formula lists "calcium chloride dihydrate (what your beer store sells)" Online beer stores offer calcium chloride. Just want to make sure we are talking about the same thing.

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Turns out I'm glad you asked that question. I've always used the dihydrate obtained from a chemical supplier because it is more stable i.e. it doesn't pick up water from the air which makes it easier to weigh accurately, it won't turn soupy on you if you leave it out etc. Now I happen to have some of LD Carlson's from the LHBS which is labeled Calcium Chloride without saying anything about the waters of hydration so as common sense would dictate that it's also the dihydrate for the reasons given above I always assumed that it had to be. But your question got me wondering. Had I ever checked it? Couldn't remember. I hit on the idea of putting the same amount of the stuff I know to be the dihydrate and the LD Carlson stuff into equal volumes of water and measuring the conductivity (no need to fiddle with all the stuff for a calcium hardness or chloride ion test). To my great surprise the LD Carlson measured 35% greater conductivity. A given weight of anhydrous calcium chloride (CaCl2) has 32% more calcium and chloride ions in it than the dihydrate (CaCl2.2H2O). So the LD Carlson offering appears to be the anhydrous. To confirm this I put about half a gram on the balance and left it exposed to room air for about 45 min. It gained 3% in weight in this time (room humidity 23% - it doubtless picked up most of the moisture when I was near the balance exhaling water vapor).

So if you get it from LD Carlson it is anhydrous. The LD Carlson stuff is in the prill (little sphere) form. That may protect it from water pick up so that it doesn't turn to soup as soon as you open the jar. If you get it from a different packager then it's anyone's guess. If it is in the prill form it is probably from the same source as LD Carlson's and probably anhydrous.

The recommendation "1 tsp per 5 gallons" is pretty approximate. It assumes the dihydrate and that the dihydrate weighs about 5 grams per tsp. This would yield about 72 mg/L calcium ion concentration and and 127 mg/L chloride. Using the same amount of anhydrous salt would increase calcium and chloride by about 32% to 95 and 169.

So they are not the same, but they work the same at equal quantities?

ajdelange said:

I've always used the dihydrate obtained from a chemical supplier because it is more stable i.e. it doesn't pick up water from the air which makes it easier to weigh accurately, it won't turn soupy on you if you leave it out etc.

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So would a good way to find out what you have be to leave a small amount sitting around in the open for some time (hours/days?) and see if it turns soupy?

At the risk of excessive horse beating, I'm trying to reconcile the original "baseline" rule of thumb with "conventional brewing wisdom" (using EZ 2.0)...

My water is sub 20ppm for everything but CaCO3 which is 90ppm. Diluting 2:1 with RO and adding back 5 grams of calcium chloride basically gives me 90ppm Ca, 140ppm Cl and almost nothing else. So no Mg, Na, or SO4 are really needed? This also results in a chloride to sulfate ratio of over 34 for me. We're really chucking that whole chloride / sulfate concept in the toilet, aren't we? Maybe that's a good thing...

I do get a good predicted pH of 5.31 with 2% sauermalz on a simple 2-row pale ale, but even with no RO dilution, my predicted pH is still 5.35. Why the 35ppm cap on CaCO3 when (at least in my circumstance) there doesn't appear to be much difference between 30ppm and the un-diluted 90?

ajdelange said:

The recommendation "1 tsp per 5 gallons" is pretty approximate. It assumes the dihydrate and that the dihydrate weighs about 5 grams per tsp. This would yield about 72 mg/L calcium ion concentration and and 127 mg/L chloride. Using the same amount of anhydrous salt would increase calcium and chloride by about 32% to 95 and 127.

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I assume this was a typo and you meant 95 and 169. If this is the case, using the anhydrous LD Carlson, the "baseline" would be 3.8 grams in 5 gal, correct?

arturo7 said:

So they are not the same, but they work the same at equal quantities?

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No (unless you consider 32% difference essentially the same and given the approximate nature of the recommendation that isn't so unreasonable). Using equal weights of the anhydrous and dihydrate the dihydrate will yield 32% less calcium and chloride as the anhydrous. Or, put another way, you can use 32% less anhydrous than dihydrate. Given that my recommendation is a teaspoonful based on the dihydrate you can use 3/4 of a tespoonful of the anhydrous.

DeafSmith said:

So would a good way to find out what you have be to leave a small amount sitting around in the open for some time (hours/days?) and see if it turns soupy?

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Probably not as they both take up water - it's just that the anhydrous takes it up faster. I tried this this morning. After about an hour and 3/4 in a room at 26% relative humidity the weight of a sample of the dihydrate increased 1.5% whereas the weight of a sample of the anhydrous increased twice this (3.2%). I took both samples into the bathroom where I took a long hot shower and pondered conditions in the Middle East without running the fan. The weight of the dihydrate increased (relative to its starting weight) by 10.4% while the total weight increase of the anhydrous was 13.6%. Both samples were starting to glisten. Don't know if the anhydrous would have turned to soup eventually while the dihydrate was spared or if both would have gone to soup eventually.

So what is a good way to tell them apart? Anything that lets you asses either the water content of the calcium content of a known weight. A 10 grams/L solution of the anhydrous form has:

SG: 1.0077 Cond: 13.82 mS/cm (8.86 g/L NaCl equiv.) RI: 1.3352

for the dihydrate

SG: 1.0057 Cond: 10.55 mS/cm (6.74 g/L NaCl equiv.) RI: 1.3347

i.e. all are about 30% higher with the anhydrous form (with RI I'm comparing the difference relative to water (nD = 1.3329) and the same with the SG.). Thus if you can weigh out 1 gram of a salt accurately and dissolve it in DI water to make 100 mL of solution and can measure SG or RI or conductivity you could compare to the numbers I have just given. The fact that the ratio of the "points" in the SG above (1.35) is pretty close to the ratio of the molecular weights (1.32) suggests that points are near linear WRT concentration so that if a brewer were to make up a solution of 5 g salt in 100 mL solution the anhydrous solution might be expected to have SG 1.0385 and the dihydrate 1.0285. The difference here is 10 points so determination should be more accurate.

Any of these methods require the ability to measure salts accurately which many home brewers do not have. There is another simpler method and that is to heat some of the prills or powder over a gas flame. A little water will come off the anhydrous (because it picks up some every time you open the jar) but a lot will come off the dihydrate which will fizz and splatter. If you can heat both and compare it will be easy to tell which is which. If you only have one it may be more difficult. A new bottle of the anhydrous shouldn't yield much water.

I suspect, but cannot assert, that the dihydrate is not sold in prill form. But there certainly lots of people selling anhydrous prills. So it's probably true that if it's in the form of prills its anhydrous. Prills that emit no (or very little) water on strong heating are very probably anhydrous.

I also suspect that the whole world (except me) is using the anhydrous so that this is not something you really need to worry about. I do wish the vendors would label their products though.

I took the LD Carlson stuff (anhydride) I've been experimenting with and using the teaspoon in a cheap set of cooking measuring spoons measured level teaspoonfuls using the time honored method of leveling by passing a spatula edge along the top of the spoon got (besides CaCl2 prills all over the place) 3.19 ( call it 3.2) grams with standard deviation about 0.08 grams. Leveling by this method rolls some prills out below the level of the rim of the teaspoon.

Repeating without the leveling by spatula but rather filling until the teaspoon looks full and filled to the rim (there is some rise in the center but not much - it looks pretty level) gave me a mean of 3.644 but the standard deviation jumped to 0.25 gram.
The lowest I measured with the eyeball leveling was 3.2 grams which is consistent with the spatula leveling average and yields 61.4 mg/L Ca++ and 108.6 mg/L Cl-. The largest I measured was 3.87 gram and that produces 73.8 mg/L Ca++ with 130.6 Cl-. These are all in the range we'd like calcium and chloride to be.


So yes, I'd say these are consistent with the guideline.

It is a very curious situation. I was under the impression that the dihydrate was the typical version supplied to homebrewers, but AJ has cast doubt on that.

This last finding just adds more creedence to that.

Bumping this question I asked in post #230 - I'm still wondering if anyone knows for sure if Pickle Crisp is anhydrous or dihydride. See post #230 for my refractometer results, which I think would indicate anhydrous based on what ajdelange has posted. But I'd feel better about using that fact in my calculations if I could be certain.

When I dissolve 0.5 grams of the anhydride in 10 grams of water I get a Bx reading of 5.1 in an instrument calibrated against sucrose. Dihydrate would be less than this. But you measured appreciably greater than either anyhydrous or dihydrate so there is something wrong with your measurement (or mine). IOW, knowing what I know at this point I can't tell.

ajdelange said:

When I dissolve 0.5 grams of the anhydride in 10 grams of water I get a Bx reading of 5.1 in an instrument calibrated against sucrose. Dihydrate would be less than this. But you measured appreciably greater than either anyhydrous or dihydrate so there is something wrong with your measurement (or mine). IOW, knowing what I know at this point I can't tell.

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AJ - thanks for taking the time to measure that. I thought that maybe I had made an error because I used a different scale to measure the water than the one I used for the CaCl2, so I repeated my measurements, this time using 0.5 grams (7.72 grains) of Pickle Crisp and 10 grams (154.3 grains) of distilled water, using the same scale for both. This is the scale I used:

https://shop.rcbs.com/WebConnect/Ma...creenlabel=index&productId=3009&route=C08J154

I calibrated my refractometer (has ATC) with distilled water, giving it about 30 seconds for the temperature to stabilize, then dried that off with a clean cloth, rinsed the window with the solution, applied more solution, closed the cover, and measured 7.2 Brix. I am sure the weight measurements are correct, so either my refractometer is bad (but seems give results for wort consistent with my hydrometer), or Pickle Crisp isn't pure CaCl2 (but the label lists the ingedients as "Calcium Chloride" - nothing else listed.

The difference in results suggests that AJ's CaCl might have picked up water. Is that possible? What sort of drying temperature is needed to drive the bound water out of the dihydrite to form anhydrate?

Deaf's higher Brix reading implies that his sample was less hydrated than AJ's, but I suppose there may be a refractometer error also? Impurities are also a possibility. Some sort of dessicant?

I'm sure that both the anhydride and the dihydrate I am using have, given the amount of time I've had them, picked up water. Even though they have been kept in jars with tightly sealing lids every time they get opened air gets in there and around these parts in the summer that air is moisture laden.

Or I could have made a math error. I'll check it again when I get a moment.

Just to eliminate my refractometer as a problem, I dissolved 10 grains of ordinary table sugar into 190 grains of distilled water to make a 5% solution. After calibrating my refractometer with distilled water, I measured the solution at 5 Brix on the nose.

The Pickle Crisp I have probably hasn't picked up much water. I bought it only a few months ago, and I have opened it only about a half dozen times, and then only very briefly, making sure to recap it tightly. Also, because this has been during the winter, indoor humidity has been low.