Measuring Calcium Chloride

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ajdelange

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We are all plagued with uncertainty concerning the amount of CaCl2 in those little 'prills' that we buy from the LHBS. Are they dihydrate (CaCl2.2H20 which is what I assumed for years) or 78 - 85% (as some flake products are specified to be) and what are the effects of storage exposed to humid air? If you have ever left any out for a day or so when you came back the powder was gone and there was a puddle of syrup in its place. There is a rather elaborate procedure for assaying calcium chloride which contains lots of low level impurities such as sodium, magnesium and hydroxyl ions (or put another way it contains some lime and lye) but the main 'impurity' from our perspective is water of hydration. Anhydrous CaCl2 is obviously 100% calcium chloride. If a recipe calls for 10 grams we weigh out 10 grams (but we had better be quick about it because if the humidity is at all high we can see the scale increment in front of our eyes as the salt takes up that water.

The monohydrate is 86% calcium chloride (assuming no impurities other than water), the dihidrate 75.5% calcium chloride and the hexahydrate only 50.7%. Correspondingly greater amounts of these salts are needed to supply 10 grams of actual CaCl2. A product with 80% calcium chloride content, assuming again that the rest is mostly water, apparently is a mix of mono and di hydrates.

It turns out that we may be able to tell approximately what is in a strong solution of calcium chloride from its specific gravity. To do this we would take a fairly large quantity (a tsp or 2) of the product in question and add it to a small volume of water - small but large enough to float your hydrometer. After dissolving the salt allow the water to cool and then measure its specific gravity with your hydrometer (narrow range hydrometers will give more accurate results). The strength of the solution, in grams per liter CaCl2, is approximately:

g/L = -684.57 + 175.12*SG + 509.45*SG*SG

where SG is the true (the difference between true and apparent can not be seen on a hydrometer) specific gravity 20/20 °C. If you determine that your solution is, for example, 68 g/L, and you need 10 grams for a brew then you must use 10/68 = 0.147 L (147 mL) of this solution.

If you have carefully measured out the salt and the volume then you can determine the concentration of CaCl2 in your powder. For example, I opened a new jar of USP anhydrous CaCl2 and weighed out 7.413 g (I just shoveled it into a 100 mL volumetric flask until it looked like about a tsp) and then added water to the mark. Thus I have 74.13 grams of powder per liter of this solution. It's specific gravity measured 1.0557 which corresponds to a calcium chloride content of 68.14 grams per liter. I can thus guestimate that my powder's calcium chloride is about 68.14/74.13 = 91.9% CaCl2. This corresponds approximately to CaCl2•1/2H2O

On the other hand if I take an old jar of the dihydrate that I first opened years and years ago and chisel the stuff in it apart I find the CaCl2 content to be only 64.5% corresponding to CaCl2•3.5H2O. It looks as if each molecule has picked up over the years an average of 1.5 molecules of water.

[Edit]A new bottle of the prills from LD Carlson assays 96.2% CaCl2 (corresponding to CaCl2•0.24H20) whereas an apparently unopened bottle of the same product that is a couple of years old gave 76% CaCl2 corresponding to CaCl2•1.94H2O.

These results are probably good enough to be useful to home brewers. One supposes that the impurities in ACS, USP or food grades of CaCl2 are not likely to be large amounts of the salts of bismuth and arsenic and that therefore, the assumption that the principal 'impurity' is water may be a fairly good one. That leaves the basic sg vs strength data (which is based on measurements dating back to the 19th century) and the model into which they have been inserted as the principal error source.

Edit: Some people might be interested in knowing the SG of solutions of CaCl2 as a function of their strengths in grams/L. This is

SG = 1 + 0.00083641*gpl -2.7043e-07*gpl*gpl

where gpl are the grams per liter.
 
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mabrungard

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This hygroscopic (sucks up water) behavior of calcium chloride is kind of a pain in the rear. Over time, it will draw water out of the atmosphere and attach that water to the solids. We brewers end up not knowing how much calcium and chloride we are actually adding to our brewing water.

On the large scale commercial side, some industries use food-grade calcium chloride solutions for their use. Once the calcium chloride is saturated with water, it won't change any further. As AJ mentions above, you can easily use a hydrometer to measure the specific gravity of a calcium chloride solution and quickly figure out what its strength is from charts such as this:

http://http://www.prog-univers.com/IMG/pdf/CalciumChloridHandbook.pdf
(see the first 2 columns in Table 7)

Since calcium chloride is easily dissolved in water, we would be well served by taking our solid calcium chloride and dissolving it in distilled water to preserve its strength at a constant level. The next version of Bru'n Water will include an alternate calculation for liquid calcium chloride use. Just insert the % strength of the solution and it will properly calculate the calcium and chloride additions with each mL of added solution.
 
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ajdelange

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Once the calcium chloride is saturated with water, it won't change any further.
Careful here. Look at Fig 5 in the Dow Handbook. Assume we have a 40% solution on a balmy northern Virginia August day: 35 °C with 90% humidity. Yes, it is that bad! Anyway the saturated vapor pressure of water at 35 is (conveniently from the chart) 42 mmHg 90% of which is 37 mmHg. The vapor pressure over a 40% solution is 20 mmHg so the solution is going to take on water and become more dilute until it reaches about 25% (interpolating by eye). Now consider the vapor pressure of water over a 20% solution. It is much closer to 37 than 20 so the uptake will be limited. We can argue whether one would have his solution outside on a day like this but the message that we want to get across is that dissolving CaCl2 does not protect you completely from the effects of atmospheric moisture but that if you keep your solution fairly dilute the situation is much improved.

The converse is also true. If you open a container of solution in dry air water will leave it in an attempt to equalize vapor pressure. So keep your solutions capped and keep them fairly weak. Under 100 grams per liter (9.26%) should be OK and should be convenient to work with as it would involve dissolving 10 grams of anhydride in 100 mL of water.

One further comment on the Dow Handbook: Note that the table is in percent and that we as brewers are usually interested in grams per liter as we can then easily measure out the number of mL we need to get a desired number of grams. Thus the polynomial is easier to use in that sense (though as the example at 100 grams/L above shows the approximate percentage is just the number of grams/liter divided by 1000 grams if we stick to the low strength which seems to be a good idea if we want to keep strength stable). Also you don't have to interpolate into the table if you use the polynomial. Note that the polynomial specific gravities are for 20/20 °C (the standard in brewing) whereas the table is for 25°C/? where we don't know what ? is (25°C ?).
 
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ajdelange

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A further thought: If you want pure CaCl2 solution you can always make it up from pure CaCO3 ( you can buy 98% pure from Duda Diesel, >99% pure from Spectrum and 99.95% 'Chelometric Grade' from Sigma-Aldrich). It's a lot easier to measure out as it doesn't pick up water the way CaCl2 itself does but for best results it should be dried in an oven before weighing.
 
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ajdelange

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A further note: a quarter liter of 10% CaCl2 solution made from CaCO3 assayed (by the manufacturer) at 99% minimum measured a specific gravity indicating 98% of the calculated strength. This verifies the validity of the polynomial I posted in #1 and that the brand new unopened anhydride I bought is indeed only about 92% CaCl2 and that the dihydrate I spoke of in that post has indeed picked up water over the years.
 

Royski

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Assume we have a 40% solution on a balmy northern Virginia August day: 35 °C with 90% humidity. Yes, it is that bad!
Not quite. That would be a dew point of 33 °C, or over 90 °F. The highest dew point we've seen around DC is about 82 °F which would be 66% relative humidity (RH) at 35 °C. A typical muggy summer day might have a dewpoint in the upper 70s in °F. Dewpoint is much better than percent humidity for expressing the moisture we feel.

Anyway, derail aside, thanks for sharing your findings.
 

trentm

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Assume we have a 40% solution on a balmy northern Virginia August day: 35 °C with 90% humidity. Yes, it is that bad!

Not quite. That would be a dew point of 33 °C, or over 90 °F. The highest dew point we've seen around DC is about 82 °F which would be 66% relative humidity (RH) at 35 °C. A typical muggy summer day might have a dewpoint in the upper 70s in °F. Dewpoint is much better than percent humidity for expressing the moisture we feel.

Anyway, derail aside, thanks for sharing your findings.
Well I wanted to call BS too but this was AJ. So I decided to do some checking and it has happened. Maybe in some microclimates in Virginia as well. http://www.wunderground.com/blog/weatherhistorian/record-dew-point-temperatures
 
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ajdelange

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I don't have extensive records for around here but certainly temperatures above 35°C are common enough and I do have a handfull of measurements made on a pretty irregular basis a couple of summers ago when I was trying to calibrate my weather station. It's been on the fritz which is why I can't go in and look at last summer's data and I'm not sure I trust those capacitor sensors anyway. The cal data was taken with an Assmann (no kidding, that's the name of the thing). The highest dew point I actually ever measured was 21 C when the dry bulb was a relatively modest 27.8 °C for an RH of 67% and a water vapor pressure of 25 mb. According to the web site linked in #7 Virginia beach has seen dew points as high as 30 °C which for a dry bulb of 35 °C corresponds to RH of 75% and a water vapor pressure of 42 mb. Finally, 35°C and 90% RG would yield a vapor pressure of 50.6 mb. Thus I agree that a dew point of 30 °C or more is most unusual the point is that in the DC metro area we encounter rather high levels of water vapor pressure in the air.
 

Callacave

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If you have carefully measured out the salt and the volume then you can determine the concentration of CaCl2 in your powder. For example, I opened a new jar of USP anhydrous CaCl2 and weighed out 7.413 g (I just shoveled it into a 100 mL volumetric flask until it looked like about a tsp) and then added water to the mark. Thus I have 74.13 grams of powder per liter of this solution. It's specific gravity wa 1.0557 which corresponds to a calcium chloride content of 68.14 grams per liter.
Forgive my stupidity AJ, but could you explain how you calculated the concentration at 68.14 g/L?
 
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ajdelange

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Forgive my stupidity AJ, but could you explain how you calculated the concentration at 68.14 g/L?
No problem. The basic formula is
g/L = -684.57 + 175.12*SG + 509.45*SG*SG

For a measured specific gravity of 1.0557 that's

g/L = -684.57 + 175.12*1.0557 + 509.45*1.0557*1.0557 = 68.0875

Note rounding error here.
 

trentm

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I measured the concentration of my CaCl2 last night, it came in at 59.73%. So I am really glad this has been posted. Still I have a few questions:

1) Do the water calculators (I use Bru'n Water) assume a 100% concentration? (I saw where Martin plans on updating his to use a known liquid concentration)

2) When using AJ's method to measure concentration, do you measure the volume of water and add the CaCl2 or add the water to the quantity of CaCl2 for a total of say 100 ml?

3) Is there a consensus that our best management practice for storing CaCl2 is to mix up a 10% w/v working solution?
 
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ajdelange

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1) Do the water calculators (I use Bru'n Water) assume a 100% concentration? (I saw where Martin plans on updating his to use a known liquid concentration)
Not AFAIK. I do all my calculations base on the anhydride and then adjust the anhydride weight by the CaCl2 % of whichever salt I am using. For yours, for example, I would, if the anhydride calcuation showed 1 gram was required, use 1/0.5973 gram.

2) When using AJ's method to measure concentration, do you measure the volume of water and add the CaCl2 or add the water to the quantity of CaCl2 for a total of say 100 ml?
The latter. If you put x grams in a flask, dissolve it and make up to 100 mL you then know your solution is 10*x grams per liter. The formula tells you, from the SG calculation, the g/L you got from what you dissolved. If it is 5*x grams/L then you must, to be precise, use twice as much of the solution as you would if you had 10*x grams/L. Approximately you would have to use twice as much of the powder as you would were it 100% CaCl2.

3) Is there a consensus that our best management practice for storing CaCl2 is to mix up a 10% w/v working solution?
As this discussion is fairly new I'd hardly say there was a consensus. 10% seems a reasonable concentration in terms of preparation and storage. The ideal solution strength from the storage POV would be that which results in a water vapor pressure equal to the water vapor pressure in the area where the bottle is to be opened. While that is true practically speaking you are orders of magnitude better off with a 10% solution than you are with the powder in terms of water pickup from the air.
 
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ajdelange

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I went by the LHBS today and picked up a couple of 2 Oz jars of calcium chloride prills. They assay 76% CaCl2 by this method and are thus effectively the dihydrate. I have edited the original post to indicate that so people new to the thread won't have to read all the way down here to get this essential piece of info.
 
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ajdelange

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It is with some embarrassment that I have to report that when I went downstairs after posting #16 the other day that I found 2 brand new bottles of LD Carlson CaCl2 sitting in front of my laptop. I grabbed the bottle in the lab assuming it was one of the new ones (wondering all the time what I'd done with the second bottle) when in fact it was some bottle of indeterminate age which I didn't even know I had. When I assayed one of the new bottles I found it to be 96% CaCl2 so it appears that the LD Carlson product is the anhydride which, over time, will, even in a closed bottle, pick up some moisture from the air. It looks as if the best advice is to either buy new CaCl2 each time you need some or assay it as described in #1 each time you draw from an older bottle. Once you have determined the g/L from the SG it is, of course, a simple matter to calculate the number of mL of that solution you need to get a desired number of grams of actual CaCl2.
 

smakudwn

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Maybe a stupid question but would it help if we vacuumed sealed the CaCL2 into bags between uses?
 

mabrungard

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I still have to contend that converting solid calcium chloride to a liquid solution that you can measure with your own hydrometer makes a lot of sense. This is a problematic mineral due to its hygroscopicity.
 

orangehero

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You can bake whatever CaCl2 you have at 200°C for 2 hours to make it anhydrous.
 
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ajdelange

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Yes, you can do that. Or at least I think you can. The last water molecule is bound tighter than the first and may require more heat but I've seen 200 °C listed elsewhere as adequate. What you must do, however, is get the stuff into a dessicator right out of the oven because it really picks up water fast. It's 35% RH in the house today and a sample of 96% CaCl2 (LD Carlson) has been gaining water from this relatively dry air at the rate of 2% of its original weight per hour for the last couple of hours.

It seems much simpler to me to just dissolve in some RO water and do the hydrometer test. For example, if you needed 25 grams you could dissolve the contents of a new 2 0z bottle (56.7 grams) in half a liter of water (note that I didn't say dissolve in water and make up to half a liter but if you do it that way you can assay the strength of what's in the bottle more easily)*, measure the SG when it cools and stuff that into the g/l formula. Say it turns out to be 1.086. Stick that into the g/L polynomial and you know you have 106.6 grams per liter and for 25 grams you would need 25/106.6 L =234 mL of the solution you made. You take that and proceed to add it to your HLT and you are done. No ovens, no 2 hr wait, no dessicator. CaCl2 is cheap enough you could discard the other half of the solution or knowing that new CaCl2 is pretty close to the anhydride you could add half a bottle to 250 mL of water etc. Or you could save the remainder adding more CaCl2 and water to it next time you need some and checking the strength of the mix when you do that. In a tightly capped bottle the strength of the solution shouldn't change much over time but as you will be checking the strength it doesn't matter if it does.

* To assay the strength of the CaCl2 note that you put 56.7 grams of powder into 500 grams of water obtaining a solution weighing 556.7 grams with a volume of 556.7/(1.086*0.998203) = 513.5 mL (at room temperature) which, at 106.4 g/L means there is 54.6 grams of CaCl2 in it. You used 56.7 grams of powder so your powder must be 100*54.6/56.7 = 96% Calcium chloride (anhydride).
 

mclaughlindw4

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I didn't know any of this before. I've had the same bottle of CaCl2 that I purchased from LHBS about a year ago. It is the kind in the prill form. So i decided to grab a handful of the little dessicant packets and bake for 2 hours at 200 degrees F. I'll store the bottle in a tupperware with the dessicant.

Before I baked, I weighed what I had, then I weighed after baking. I didn't get much of a change. It was 37.2 grams before and 36.8 grams after.

So i guess my CaCl2 didn't change much from my initial purchase?
 

mabrungard

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At 200F, you didn't really do anything. As pointed out above, the temperature needs to be pretty high to get the hydrated water out of the molecules. Give it another try at 200C.
 

schematix

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I am going to atttempt to make a CaCl2 solution for my brew this weekend and wanted to make sure I understood how to do this. I have a several year old packet of CaCl2 that has been stored air tight and dry but I question its original form and current purity. I assume its somewhere between 75-95% anhydrous at this point but i don't really know. My goal is to make a maximum 20% w/w solution.

This is my proposed procedure:

1. Measure (by weight) all of the CaCl2. Let's say it weighs 50g.

2. To get maximum 20% strength, I will weigh 5x50g=250g of distilled water and add the CaCl2 and mix until dissolved. I will then chill to room temp and measure the SG with my hydrometer. Let's say it measures at 1.150. Using some of the provided formulas this would mean the real w/w is about 16%, and that corresponds to about 190g/L.

3. Assuming I need 3.2g of CaCl2 to hit my numbers, I'd measure out 16.84mL (3.2g/190g/L) of this solution, and add it to the main water.

4. Done. Go drink a beer in celebration.

Will this procedure yield my desired results?
 

schematix

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I followed my procedure above to create a ~20% solution yielded a ~12.5% solution. So my CaCl2 was holding onto a lot of water and throwing my weight measurements into question. Glad I did this!
 

schematix

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So i went back to this CaCl2 solution I made some weeks ago and there is quite a bit of a white powdery material in the bottom of the jar.

Any idea what this could be? The solution was completely dissolved when i first made it, minus maybe a few grans that wouldn't dissolve.
 

schematix

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Where did the calcium carbonate come from? Has my solution changed in strength?
 

schematix

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Hmm. It has been in a tightly sealed canning jar.

I'm not a chemist by trade (I prefer controlling electrons), but stretching my memory back to college it doesn't look like that reaction is balanced. Is it reacting with atmospheric oxygen and carbon dioxide?

Is this reaction reversible under any conditions or is this solution now a throw-away?
 
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ajdelange

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It isn't because I left out the CO2 which does indeed come from the air. If it's not chalk then I don't know what it is. Try decanting the liquid. That is still useable. The strength probably did not change very much but you could always recheck SG. Try adding some vinegar to the residue to see if it fizzes.
 
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ajdelange

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I'm not a chemist by trade (I prefer controlling electrons)...
Actually, that is exactly what a chemist does.

More to the point, I had another thought. Perhaps it is some calcium chloride which has come out of solution. Try shaking the jar. If the stuff redissolves then that's what it doubtless is. If it doesn't redissolve and it doesn't fizz under acid then I don't know what it is.
 

JKoravos

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Great post. I'll never forget the first time I tried to weigh out some CaCl for a batch. It was the height of summer and I'd just bought a brand new pouch from the LHBS. I brought it out to our prototyping lab at work to use the gram balance. Our prototyping lab is in a smaller outbuilding from our main facility and is not climate controlled (no A/C!). I zipped it open and started weighing, before I could even get enough on the scale to hit my number I was already dealing with a goopy mess. I couldn't believe it.

I can't argue with the method in the OP, but my tack on this would likely be to create the anhydride in the oven, get a baseline mass of anhydrous and calculate the remaining quantity each time I pulled from the jar on an anhydrous basis.

For example, say cooked my CaCl2 to create 100g on anhydrous and then put it away.

Next time I brew I measure the weight of the jar and it's 105g. I know that each gram I use will have (100/105=) .952g of anhydrous. If I want 2g on anhydrous, I would weigh out (2/.952=) 2.10g of partially hydrated CaCl2. Now I know I have 98g of anhydrous remaining (and 102.9g of total mass remaining).

You can continue to repeat this process as long as you track your anhydrous content.

I can certainly see the merit in both methods, though. Now I'm tempted to go bake my CaCl2 to see what I've actually been using.
 

mclaughlindw4

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Great post. I'll never forget the first time I tried to weigh out some CaCl for a batch. It was the height of summer and I'd just bought a brand new pouch from the LHBS. I brought it out to our prototyping lab at work to use the gram balance. Our prototyping lab is in a smaller outbuilding from our main facility and is not climate controlled (no A/C!). I zipped it open and started weighing, before I could even get enough on the scale to hit my number I was already dealing with a goopy mess. I couldn't believe it.

I can't argue with the method in the OP, but my tack on this would likely be to create the anhydride in the oven, get a baseline mass of anhydrous and calculate the remaining quantity each time I pulled from the jar on an anhydrous basis.

For example, say cooked my CaCl2 to create 100g on anhydrous and then put it away.

Next time I brew I measure the weight of the jar and it's 105g. I know that each gram I use will have (100/105=) .952g of anhydrous. If I want 2g on anhydrous, I would weigh out (2/.952=) 2.10g of partially hydrated CaCl2. Now I know I have 98g of anhydrous remaining (and 102.9g of total mass remaining).

You can continue to repeat this process as long as you track your anhydrous content.

I can certainly see the merit in both methods, though. Now I'm tempted to go bake my CaCl2 to see what I've actually been using.

I like this. I just don't want to be bothered with making solutions of things at home. I did the baking thing and have just been storing it in a air tight container with a crap load of desiccant packets in it.
 

schematix

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After 2 brews using the CaCl2 solution, I am a convert. The solid form picks up water from the air at a ridiculously fast rate. You can put it on an precision scale and watch the weight climb.

Baking it also a hassle. It takes a few hours to dry it out, and as soon as its out of the oven it's collecting water again. Not to mention a big waste of energy running the oven that hot for that long.

The newest Brun Water also includes calculations for CaCl2 solution.
 

mclaughlindw4

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I have a toaster oven. I can't imagine it uses that much energy and I don't think there is much moisture for it to pick up in my little dessicator.

If I had a good means of measuring smaller volumes and a precision hydrometer I might do the solution thing but I don't right now. So doing that probably wouldn't be any more accurate then what I am doing.
 
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