Given your base water profile this would be a good target profile for a Dunkleweizen:
Actually, that's not correct. It's the additions needed to reach the target profile that need to be balanced, and they are. It's his starting water profile that appears unbalanced but there are several reasons why this may be. Usually, it's due to the fact that not all ions are reported in the water analysis. For instance, if a source of Ca has an associated anion that is not of the type that's reported, the Ca would show on the report but not the anion. This would cause the ions to appear to be unbalanced. The additions, however, do need to be balanced. Given his initial water profile, the target profile above is, in fact, achievable.
You should see my water profile if 100% carbon filter. Water here is very hard. Without a softner it turns utensils that are plastic white. Our drinking water is on hard water with the filter. If I just do RO Water from the store. Not sure what to use besides Calcium Chloride and since german wheats do not need sulfate(gypsum). Sulfate brings out harshness in bittering. That is why English Pale Ales use it.
You certainly don't want the sharp bitterness that a Pale Ale would have but their sulfate levels are generally around 225-275ppm, way more than the 80ppm I suggested. The generally accepted sulfate levels for a Dunkleweizen are around 50-150ppm.
Just reread the question and realized that it was a little bit more involved. I don't know any calculator that will give you concentrations of individual salt solutions directly out of a profile without fiddling around with the amounts of each salt. Problem is there's more than one way to get there depending on what salt you have on hand and your starting border profile.
Ahh, that makes sense.
Not CaCl?
I always mix them up... Thanks.
CaCl2 comes from the factory as the anhydride. By the time you open up the bottle and use it just once, it will already be down to about 94%-96% CaCl2, so perhaps you can sort of ballpark figure that with each use it looses about 5% due to water adsorption. But you can't ever buy it as the dihydride (to my knowledge). When it is down to about 75.5% it is at the dihydride state. But it doesn't stop adsorbing water at that juncture. It continues until it turns into a mush water/salt goo.
Yes that's true. However, my plastic bag seals very well and since years it's not changing much, judged simply by the looks of it. I know it continues to attract water but I close it quickly and push all the air out before so it seems to stay within limits that are ok. I just overdose it by about a quarter or a third compared to the anhydrate and so far I seem to be doing pretty fine with that approach.
The *label* on my 1kg tub of CaCl2 says "CaCl2.2H2O". Of course, that doesn't meant it is, for reasons you explained. I guess I could bake it or get a hydrometer that's actually accurate and calculate what it really is (= was), but as long as it's a reasonably loose solid powder, I don't think it's off enough for me to care. (I rarely use Cl salts anymore anyway, after I started using HCl for pale beers).
I could be wrong about no one retailing the dihydrate. But it seems like it would take a lot of effort (not to mind expense) to arrive at precisely the dihydrate state, and what might be gained in doing so?
Yes - that is exactly how you do it. Plug the gravity into a calculator and it will give you the concentration. Enter the concentration into your preferred water calculator.
None of your described methods is effective.
Since when was boiling ineffective at removing chlorine?
Yes, well, efficient is very subjective. Everything is give and take, boiling may not be efficient on energy, but it's efficient with time. Meanwhile campden tablets are cost and energy efficient, but definitely not time efficient. Regardless, Miraculix said effective. Since they do the job at removing chlorine, they are effective. I'm not really worried about "efficient" for reasons mentioned above.
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