I almost lumped them all together (except chalk) but then started thinking that while a gram of gypsum per liter is a lot it puts "only" 558 mg/L sulfate in the water which is not that much greater than what is seen in some recipes, water profiles etc. Lets say that 400 mg/L is the most sulfate anyone would want to use. 558 is 0.14 orders of magnitude above that but a gram per liter is only 0.3 orders of magnitude less than its solubility. So maximum sulfate addition to solubility limit in cold water is 0.44 orders of magnitude (less in hot water) and its been noted here that getting gypsum to dissolve can be a bit of bear - especially if one tries to dissolve it in hot water.
By contrast, a gram of calcium chloride (monohydrate) would produce about 550 mg chloride. That's a lot of chloride. Lets say 250 mg/L is the most anyone would ever want. That's 0.34 orders of magnitude less. But the solubility of calcium chloride is around 700 grams/L (at room temperature) which is 2.8 orders of magnitude more than 1 gram per liter. Ratio of maximum addition to solubility is thus over 3 orders of magnitude whereas for the gypsum its less than one. I guess in my mind calling a ratio "approximately infinite" is more supportable when it has a few 0's in it (to the left of the decimal place). That's the main reason I separated gypsum (and calcium hydroxide) out form the others but I probably had the scheme I use to add salts to the HLT which is to add all the salts to water to make up 36 mL of solution because my HLT is 36" tall and add 1 cc of that solution for every inch of water I put into the HLT. This works just fine for calcium chloride but would not work for gypsum. I guess just having it is suspension would be OK as long as I remembered to shake the tube before drawing an aliquot.
Yes I am and it's because I have seen so many postings, recipes and recommendations in which huge quantities of it are called for in situations where it is more probable that acid would be of benefit. If brewers were a little afraid to use for fear that the ghost of A.J. will come to haunt them that would be OK with me.
The apparent problem here is predicting the alkalinty. Where chalk has been dissolved by CO2 about half the alkalinity is contributed by the chalk and half by the CO2 as long as the pH is middling. I was troubled by Kai's method because once you open the bottle CO2 escapes and the pH rises and unless you know the pH you can't know the alkalinity and while the you are measuring pH CO2 is continuing to escape and eventually enough will escape that the chalk will start to precipitate back out. It was only last night that it dawned on me (Kai might have already figured this out) that it doesn't matter what happens in the soda bottle as long as you get the liquid into the liquor tank before precipitation occurs. At that point you have chalk dissolved by CO2 and even though you don't know how much CO2 may have escaped if you have the pH of the water and the amount of chalk you dissolved with the CO2 you can solve for the alkalinity. To do that requires iterative solution of a set of nonlinear equations but, and this also dawned on me last night, the solution is the same every time you do it, so if the solutions as a function of pH can be fit with a simple function you are there.
alk = [CaCO3]*( 1.0167 -0.39579*exp( -(pH - 4.7)/0.43063 ) )
Here [CaCO3] is the amount of chalk you put in the soda bottle (in mg) divided by the number of liters in the final mix i.e. the water you added the soda bottle contents to plus the water in the soda bottle. pH is the pH of the mixture after addition of the soda water bottle contents. The alkalinity is defined in terms of a titration to an endpoint at pH 4.3. This fit is valid over the range 4.7 < pH < 9. Note that at pH > 6.5 the alkalinity is approximately 101.7 ppm as CaCO3 per 100 mg/L CaCO3 in the mix and this is, of course, why we use ppm as CaCO3 as the unit of alkalinity. Another caveat that goes with this is that the solutions which this fit describes are for the situation where the only acid in the system is carbonic. It would be possible to do the same this if other acids were present but as the diversity of possible other acid types and blends is infinite pre-computation as I've done here isn't practical.
I've from time to time said that all this water chemistry stuff would be a lot easier if one could buy calcium bicarbonate powder. This may be the next best thing - calcium bicarbonate solution and I even though briefly about offering such a thing as a product. Essentially soda water with some chalk dissolved in it (but no sodium chloride).
No argument there. The reason we have trouble with chalk is that we can't easily handle it as mother nature does. She uses CO2 to dissolve it and we must too if we want similar results.
