Measuring Calcium Chloride

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I just finished making up one liter of CaCl2 solution. I added 109.74 grams and carefully brought it up to 1 liter in an Erlenmeyer flask, and when it cooled it measured in at a final SG of 1.077. From this I believe that my "fresh' CaCl2 was about 86.6% by weight as CaCl2, or somewhere between the anhydrite and the dihydrate form.

If I'm calculating this correctly, then 7.95 ml of my solution (call it 8 ml) is the equivalent of 1 gram of the most commonly assumed dihydrate form. Does this sound correct?
 
Without a precision hydrometer, could I use a refractometer instead?

I just tested my 5 month old bottle of LD Carlson Calcium Chloride.

I had about half a bottle (26.8g). I heated 500mL of water (which I weighed out since my flask is not exactly precision), dissolved the CaCl2 into it, and cooled to room temp.

I used a regular calibrated hydrometer to measure it and it looked like it was 1.039. I used my refractometer to double-check it but that only showed 1.029. I'm guessing there's some reason a refractometer doesn't work for this purpose since nobody else mentioned it.

I plugged that SG into your equation and got 46.86g/L.

Since my original g/L was 53.6 it looks like my CaCl2 is ~86%.

Sound about right?
 
A refractometer works just fine for this purpose but you cannot use the Brix scale. You must reduce the Brix reading to a refractive index reading and then look up the calcium chloride concentration corresponding to that refractive index. You can find tables of Brix vs RI and CaCl2 concentration vs RI on line. From them you can obviously prepare a table of Brix reading vs CaCl2 concentration.
 
And it tells us that the data are in the CRC Handbook. Good to know.

While we are at it one can use conductivity as well but in this case one needs to know the conductivity of CaCl2 solutions and the scale that one's particular conductivity meter uses.
 
Can someone help please. Twice now I have made a 1.090 solution with store bought distilled water. After about a month I am getting a lot of calcium, I think it's calcium, dropping out of the solution. When swirling the jar it looks like a snow globe. Is this something to worry about? Am I making the solution to strong? What am I doing wrong?
 
Mine has this noticeable white precipitate also. I don't believe that the CaCl2 beads I bought at the LHBS are very pure. It's as if there is chalk in it. Fore sure it is something that is white and highly insoluble in water.

Could it be that the manufacturer is adding chalk (or something) as an anti-caking agent? Or worse, as a cutting agent to reduce their cost?

I presume that if it isn't dissolved, it also is not throwing the hydrometer SG reading off.

Next time I'm just going to buy Ball's product called "Pickle Crisp", which is far less expensive, and is likely no worse. Purely speculation here, but the company that supplies the LHBS's may be purchasing and re-packaging something similar for all I know. Come to think of it, I've never noticed a white precipitate in pickles processed with 'Pickle Crisp'.
 
ajdelange said:
Take some of the precipitate (filter it out through a paper towel) and drip vinegar on it. Does it fizz?
Click to expand...

Damn I dumped it out already, if it fizzes does it mean calcium? I made the solution with Ball's pickle crisp first time. Then I thought maybe purity issues so I bought a 1lb of CaCl2 at the LHBS, LD Carlson was the name I believe, and it did the same thing.
 
PartagasD4 said:
Damn I dumped it out already, if it fizzes does it mean calcium? I made the solution with Ball's pickle crisp first time. Then I thought maybe purity issues so I bought a 1lb of CaCl2 at the LHBS, LD Carlson was the name I believe, and it did the same thing.
Click to expand...

I wasn't going to mention the brand name, but now that the cat is out of the bag, mine is from the same source.
 
Silver_Is_Money said:
Mine has this noticeable white precipitate also. I don't believe that the CaCl2 beads I bought at the LHBS are very pure. It's as if there is chalk in it. Fore sure it is something that is white and highly insoluble in water.

Could it be that the manufacturer is adding chalk (or something) as an anti-caking agent? Or worse, as a cutting agent to reduce their cost?

I presume that if it isn't dissolved, it also is not throwing the hydrometer SG reading off.

Next time I'm just going to buy Ball's product called "Pickle Crisp", which is far less expensive, and is likely no worse. Purely speculation here, but the company that supplies the LHBS's may be purchasing and re-packaging something similar for all I know. Come to think of it, I've never noticed a white precipitate in pickles processed with 'Pickle Crisp'.
Click to expand...

I tried both and both did the same thing. It's clear as water for a few weeks then I noticed the white precipitate collecting at the bottom. I'll just go back to measuring dry if I can't prevent it. A pound of CaCl2 at the LHBS was better priced for me than the pickle crisp
 
PartagasD4 said:
I tried both and both did the same thing. It's clear as water for a few weeks then I noticed the white precipitate collecting at the bottom. I'll just go back to measuring dry if I can't prevent it. A pound of CaCl2 at the LHBS was better priced for me than the pickle crisp
Click to expand...

The precipitate (or its precursor) must also be present when it is in the dry (dihydrate) beads state. Any particles not in solution will not to my knowledge alter the solutions SG, so it is still well more precise to use it as a liquid.
 
My thinking is that it might be chalk formed when CO2 from the air dissolves in the solution. Were I correct in this thinking the precipitate would, of course, fizz when exposed to an acid (such as vinegar). I'm not too comfortable with this thesis as dissolved CO2 leads to carbonic acid with only very, very little carbonate resulting unless the pH were raised by something such as lime (Ca(OH)2) which is, often, a contaminant of calcium chloride. As a first check I'd take a new solution and check its pH. If it is high (and I have seen this) then that's indicative of the presence of Ca(OH)2 and the precipitate is quite likely CaCO3. If this is bothersome the fix would be to neutralize the solution to pH 7 with hydrochloric acid but I think I'd rely on Silver's observation that if it is lying on the bottom it isn't effecting the SG of the solution and so just recheck the SG of the solution and use that value to calculate the CaCl2 concentration.
 
Thank you for the help and knowledge. I made a weaker 1.070 solution and will go with don't worry about it as long as the gravity remains 1.070.
 
ajdelange said:
My thinking is that it might be chalk formed when CO2 from the air dissolves in the solution. Were I correct in this thinking the precipitate would, of course, fizz when exposed to an acid (such as vinegar). I'm not too comfortable with this thesis as dissolved CO2 leads to carbonic acid with only very, very little carbonate resulting unless the pH were raised by something such as lime (Ca(OH)2) which is, often, a contaminant of calcium chloride. As a first check I'd take a new solution and check its pH. If it is high (and I have seen this) then that's indicative of the presence of Ca(OH)2 and the precipitate is quite likely CaCO3. If this is bothersome the fix would be to neutralize the solution to pH 7 with hydrochloric acid but I think I'd rely on Silver's observation that if it is lying on the bottom it isn't effecting the SG of the solution and so just recheck the SG of the solution and use that value to calculate the CaCl2 concentration.
Click to expand...

I experienced the same thing quite some time ago and stopped using liquid CaCl2 because of it.

Do you know if this reaction reaches an equilibrium? I.e. could you make up a batch, let it sit for X hours/days/months, decant/filter the liquid off the precipitate, measure, and then keep the resulting liquid? Or will it continue to precipitate calcium carbonate for a long time?
 
Please keep in mind that the reaction is, at this point, a hypothetical one. Given that the hypothesis is true then it would proceed to the point where the carbonic acid has neutralized all the OH- in the solution and then stop. Or, as I noted above, you could jump the gun and neutralize the OH- with HCl when you make up the solution.... if this is indeed what is causing the precipitation.
 
The vinegar test produced a steady stream of gas bubbles when poured onto the precipitate.
 
schematix said:
The vinegar test produced a steady stream of gas bubbles when poured onto the precipitate.
Click to expand...

It's an indication of acid neutralizing calcium carbonate and releasing CO2.
Vinegar is generally a weak solution of acetic acid in water. A lot of CaCl2 is process from limestone beds and I wouldn't be surprised if some limestone impurities were tagging along with the calcium chloride crystals.
 
That pretty much confirms the hypothesis that
1)The supplied calcium chloride contained impurities in the form of hydroxides: Ca(OH)2, Na(OH)...
2)These are being neutralized by carbonic acid resulting in the precipitation of calcium carbonate: Ca++ + 2(OH)- + H2CO3 --> CaCO3 + 2H2O


The fixes would be
A) Ignore but check the specific gravity of the clear liquid to account for the calcium lost (probably barely detectable if at all).
B) Open and close the solution container as briefly as possible
C) Neutralize the (OH)- with hydrochloric acid when the solution is prepared.
 
Might the Ca(OH)2 contaminant present within the CaCl2 beads be of sufficient quantity to throw mash pH's noticeably off to the high side? Is there any logical and/or justifiable reason for it to be there in the first place? This certainly brings mash pH measurements into question, and since not accounted for, will throw a monkey wrench into the works when using software for mash pH. The software will inevitably take the blame, albeit unjustifiably.

Is there another source for CaCl2 beads that does not exhibit this irritating contamination? i'm not happy at all with LD Carlson in regard to this matter. It makes me wonder about the quality of their other products.

I've only opened mine for the briefest of time periods, and only three times since making it, and the precipitate is pronounced. Should I toss it out and start over?

I don't have any HCl. Can I use 88% lactic acid to neutralize my CaCl2 solution?
 
I'm gonna make up a fresh batch this weekend and will observe the change in specific gravity over a month or so.

Will be interesting to see if it changes significantly or not.
 
The batch I made yesterday has already formed a haze at the bottom of the jar. No big solid pieces but I'm sure it won't be long.
This is going off an old guy's memory and it's been a 1/2 year or so but I believe when I tested the gravity of my solution it remained at 1.090 throughout about 6 months or so. I just dumped it because I didn't like the big chunky looks of it and I was not adding the calcium I thought I was. Of course now I'm reading it was more lime/chalk than calcium. I think I will just run it through a coffee filter next time instead of dumping/wasting.
 
Silver_Is_Money said:
Might the Ca(OH)2 contaminant present within the CaCl2 beads be of sufficient quantity to throw mash pH's noticeably off to the high side?
Click to expand...

Probably not but it should be easy enough to check. Place a known amount in some distilled water and check the pH. If the pH is high (and I only think I have observed this once) titrate with acid to neutrality. The amount of acid used is the alkalinity of the known mass of calcium chloride. Use it in the mash pH estimation algorithm as you would use any other alkalinity.


Silver_Is_Money said:
Is there any logical and/or justifiable reason for it to be there in the first place?
Click to expand...

Yes. Most CaCl2 comes as a byproduct of the Solvay process in which a solution of ammonium chloride is treated with quick lime:

2 NH4Cl + CaO → 2 NH3 + CaCl2 + H2O

CaO is a strong base and if the reaction does not go completely to the right there will be hydroxyl ions present from the reaction of the quicklime with water.

Silver_Is_Money said:
This certainly brings mash pH measurements into question, and since not accounted for, will throw a monkey wrench into the works when using software for mash pH. The software will inevitably take the blame, albeit unjustifiably.

Is there another source for CaCl2 beads that does not exhibit this irritating contamination?
Click to expand...

Fisher specs their ACS grade calcium chloride to produce solutions with pH 4.5 - 6.5. Clearly you would not have this problem with that material.


Silver_Is_Money said:
i'm not happy at all with LD Carlson in regard to this matter. It makes me wonder about the quality of their other products.
Click to expand...
You are not buying ACS grade from LD Carlson. It is presumably FCC. The specs for one supplier's FCC CaCl2.2H20 that I found in a net search promise alkalinity as less than 0.3% as Ca(OH)2. So apparently some level of alkalinity is allowed under the FCC regs for this chemical.

A couple of observations:
•At the level at which we use CaCl2 in brewing a fraction of a % contamination with Ca(OH)2 isn't likely to be significant.
•It is easy enough to neutralize any alkalinity with hydrochloric acid (requires a pH meter or test strips) and eliminate the problem.
•Exposing the solution to air (bubbling air through it) will neutralize the alkaliity.


Silver_Is_Money said:
I've only opened mine for the briefest of time periods, and only three times since making it, and the precipitate is pronounced. Should I toss it out and start over?
Click to expand...
No, it's fine. As just noted the alkalinity has been neutralized when the precipitate forms. Just check the SG of the supernatant to see if it has changes appreciably. Use the SG of the supernatant in the CaCl2 concentration calculation.


Silver_Is_Money said:
I don't have any HCl. Can I use 88% lactic acid to neutralize my CaCl2 solution?
Click to expand...
You can use any reasonable acid (no HFl please!) but it seems to me the easiest would be carbonic as that happens automatically anyway,
 
PartagasD4 said:
... I was not adding the calcium I thought I was. Of course now I'm reading it was more lime/chalk than calcium.
Click to expand...
It is chalk and chalk contains calcium so you are not adding all the calcium you originally put into the mix. The amount you have lost to the precipitate is, however, insignificant based on the reports of undetectable decrease in SG.
 
The reason I recommend HCl for the neutralization is, if you all haven't already figured this out, that the contaminant Ca(OH)2 gets converted to the desired CaCl2.
 
ajdelange said:
A refractometer works just fine for this purpose but you cannot use the Brix scale. You must reduce the Brix reading to a refractive index reading and then look up the calcium chloride concentration corresponding to that refractive index. You can find tables of Brix vs RI and CaCl2 concentration vs RI on line. From them you can obviously prepare a table of Brix reading vs CaCl2 concentration.
Click to expand...

Hi AJ and thank you for sharing your vast knowledge.

I'm a newbie home brewer and a really know nothing about it yet. So please give me some slack if this sounds stupid:

Let's say i measure my solution with a refractometer and get 16 Brix. I then use this

http://web2.airmail.net/sgross/fermcalc/FermCalcJS.html

calculator and it gives me a refractivity value of 1,3573.

Does this mean that i can use this next chart and find the concentration level of my solution? (around 10% by the chart)

https://www.mt.com/au/en/home/supportive_content/concentration-tables-ana/Calcium_Chloride_re_e.html

Thank you!
 
That is great as i think my hydrometer is somewhat inaccurate but my refractometer is just fine!
Thanks again.
 
So i decided to make my first CaCl2 solution:

As i had posted the question both here and on a youtube channel i subscribed a few months ago, i will reproduce my comment there:

"Ok i have the numbers: Used 60 grams of CaCl2 (crystals) and added RO water to 500 ml mark, just as you did in the video. I just added the crystals first and water after to the 500 ml mark.
Results:
-Hydrometer : 1071
-Scale: 535 grams which gives (535gr/500ml) 1,07. Hydrometer is close or dead on.
-Refractometer: 13 Brix. This gives me a 1,3525 RI, and by checking the CaCl2 RI table i get 8%.

Using g/L = -684.57 + 175.12*1,071 + 509.45*1,071*1,071 i get 87,34 gr/L which is 8,7% w/vol.
That is a 0,7% difference and i really cannot say which one is correct. My Hydrometer came with a kit and is crappy, my scale is not a precision one (rounds grams) and my refractometer is also a cheap chinese one.
If i'm thinking right, if i want 2 grams of calcium chloride in my beer, ii would add:
Hydrometer reading: 2000/87,34= 22,9 ml of solution
Scale reading: 2000/86,07= 23,23 ml of solution
Refractometer reading: 2000/80= 25 ml of solution

Am i thinking right?

Edit: i think i am making a mistake and in fact the refractometer reading is w/w. If so a 1070 gravity would mean i had to add in 23ml*0,998203x1,070 =24,8 grams. That is almost dead on the refractometer reading.

P.S: I must say that no credit of this is mine (if this is right of course). I just gathered info in Homebrewtalk and Google and decided to try it! Special thanks to user ajdelange from HBT forum."

Hope to be correct on the calcs. If it were not for this forum i would simply weight X grams of whatever calcium chloride i had and add it to my beer.
 
Yes, the Mettler Toledo table does give CaCl2 content in w/w. We can still live with that. When I originally undertook this work I wanted to give people the means to determine CaCl2 content from hydrometer, refractometer or TDS meter readings. As I thought most would use hydrometers I posted that first and never finished up on the other two parts. This Mettler Toledo table gives us the path to a formula for grams CaCl2 per liter (the handiest form as one determines the amount of liquid to measure out just by dividing the requirement in grams by the grams/L strength of the solution.

First off, the Mettler Toledo table can be summarized by the formula

w/w = 426.203*(RI - 1.333) -760.013*(RI - 1.333)^2 + 1953.22*(RI - 1.333)^3

or

w/w = -7609.05 + 15157.4*RI -10217.6*RI*RI +2347.26*RI*RI*RI

The first one returns exactly 0 for RI = 1.333, the second a slightly negative number ( -0.00073). For most purposes it doesn't matter which you use.

These formulas give us the grams CaCl2 per gram of solution from the RI. We have a formula (OP) for calculating grams CaCl2 per liter of solution from the SG and as the density is simply 0.998203 times the SG we can easily find the grams CaCl2 per gram solution by dividing g/L by 0.998203. This gives us grams/gram as a function of SG. We invert this function and so are able to calculate SG from g/g obtained from the Mettler Toledo table fit equation. Inserting this SG back into the original (OP) formula we have a path from RI to grams/L. We calculate a bunch of pairs and do a curve fit. The result is

grams CaCl2/Liter solution = 341613 -1.00106e+06*RI +1.09765e+06*RI^2 -536044*RI^3 + 98840*RI^4

This is the Brew Science forum but in case anyone isn't familiar with the notation -1.00106e+06 (which would, anywhere but the US, be written -1,00106e+06) means -1001060 i.e. the decimal point is shifted by the number of places after the e, plus meaning to the right.

The fit is fourth order which it needs to be in order to get accuracy of 0.01 gram per liter.
 
I wouldn't think that would be the form he'd want as most home brewers have refractomers that do not read RI but rather Bx. The formula

grams CaCl2/Liter solution = 5.9392*Bx + 0.050098*Bx^2 -0.00018856*Bx^3 + 4.7399e-06*Bx^4

in which Bx is the Brix reading of a refractometer calibrated for sucrose (what most home brewers use) would be more valuable
 
ajdelange, have you updated the baseline measurements from the main post here? Are you still recommending that everything still be halved as compared to what is currently posted?
 
urg8rb8 said:
ajdelange, have you updated the baseline measurements from the main post here? Are you still recommending that everything still be halved as compared to what is currently posted?
Click to expand...


That's my understanding. I've been a big fan of the primer for a year or so. The only time I get really aggressive with salts is for the NEIPA style. Martin has a nice neipa profile in the summer issue of Zymurgy. It made for a really nice Beer.

I have found that less is more with CaCL and Gypsum in most cases.
 
Dcpcooks said:
That's my understanding. I've been a big fan of the primer for a year or so. The only time I get really aggressive with salts is for the NEIPA style. Martin has a nice neipa profile in the summer issue of Zymurgy. It made for a really nice Beer.

I have found that less is more with CaCL and Gypsum in most cases.
Click to expand...

I am about to make a West Coast style IPA for the second time. First time, I treated the water with the original primer. But now that I am going to do it again, I want to see if I can detect a difference if I half everything.
 
I found it more pleasant at half of the recommendation from the original primer with west coast IPA's . It's subjective imo but I liked the beer better. I do think you'll notice a difference.
 

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