Calcium from Calcium Hydroxide, Ca(OH)2, Slaked Lime, Pickling Lime?

[Silver_Is_Money]

How many of you are using calcium hydroxide to both add alkalinity and boost the calcium ion content of your mash water without the need (or when there is not desire) to drive sulfate and chloride levels up? Has this method been working well for you?

54.1% of the mass of Ca(OH)2 is the calcium ion. The richness of this source is likely unsurpassed for adding calcium. And it brings with it zero chloride and zero sulfate. The added alkalinity that comes along with it can easily be moderated or outright eliminated by essentially flavorless phosphoric acid. And for darker brews the alkalinity is likely going to be directly beneficial, as would be raising the calcium ion concentration of the mash water.

The greatest drawback may be the safety hazard associated with using it. Are there other obvious or not very obvious drawbacks that I'm missing?

I've discounted the use of calcium carbonate due to its minimal solubility in water. If it was more soluble it would have many of the benefits and few of the hazards.

 

[Robert65]

I do. I've only used it in beers requiring alkalinity be added to my RO, which are generally dark beers that also benefit from reduced sulfate (less gypsum I still get some calcium from the chloride salt, since to get it all from slaked lime would provide far more alkalinity than needed. But you raise an interesting point, that the excess could easily be neutralized. Something similar has occurred to me just recently. I've been mulling over the idea that I might like to incorporate Sauergut into my pale lagers, as it is a key component of the authentic German flavor profile; but I don't need to acidify my pale mashes. In fact, I need to add a touch of alkalinity to my RO, if any Munich or Cara malts are included, to hit a mash pH of 5.5-5.6. It has occurred to me that I could use slaked lime to counter the acid. It just seems a bit Rube Goldberg to me, and so far I've just accepted that my materials and process differ from a German brewer's dealing with naturally carbonate waters, and so does my beer. But the uses of slaked lime are probably underappreciated.

 

[MaxStout]

I've used it a few times, in dark, malt-forward beers where I wanted to bump up pH in the mash, but not add more sulfate or chloride. I bought a bag of "pickling lime" from the home-canning section of the store. Way more than I'll ever use, for a few bucks.

Ca(OH)2 isn't as caustic as sodium hydroxide, but you should still be careful handling it, as it can burn the skin and eyes. In the amounts you would be using in brewing water, it won't be harmful.

 

[Silver_Is_Money]

For the benefit of those who may be confused, all of the names listed in this threads subject line are simply different names for the exact same chemical.

 

[@RoyalGallon]

Oddly enough I use slaked lime for the complete opposite purpose.

I have very hard water here and so treat it with slaked lime as per this link to drastically reduce the alkalinity and Ca content.

works really well - the amount of precipitated CaCO3 that drops out is amazing!

 

[Silver_Is_Money]

Oddly enough I use slaked lime for the complete opposite purpose.

I have very hard water here and so treat it with slaked lime as per this link to drastically reduce the alkalinity and Ca content.

works really well - the amount of precipitated CaCO3 that drops out is amazing!

Ah, that has me thinking. And concerned. Originally my plan was to utilize a mix of baking soda and slaked lime in the mash water. Now I'm wondering if the baking sodas carbonate ions will combine with the slaked limes calcium ions and immediately precipitate out as CaCO3? That would defeat the purpose of boosting calcium in the mash with slaked lime, and leave one with the same problem of calcium carbonate being virtually insoluble as I had mentioned in post #1 above.

 

[Robert65]

I'm getting in over my head here, but... baking soda's anion is bicarbonate, not carbonate... and isn't there a certain range (minimum requirements to start reaction and limit to how far it can proceed) of both carbonate (not other species) and calcium for lime softening to work... and don't plenty of brewers anyway successfully use a combination of baking soda and pickling lime to raise mash pH? I'll leave it to you who have a better grip on all this...

 

[balrog]

I used to use both slaked lime (for Ca without Cl or SO4) and phosphoric.
As soon as I added the acid to the water in the kettle, it clouded and I subsequently decided that meant something was precipitating out, making my ion calculations suspect. For now, I do not tend to use lime & phosphoric at the same time, but haven't felt the need.

 

[Silver_Is_Money]

I'm getting in over my head here, but... baking soda's anion is bicarbonate, not carbonate... and isn't there a certain range (minimum requirements to start reaction and limit to how far it can proceed) of both carbonate (not other species) and calcium for lime softening to work... and don't plenty of brewers anyway successfully use a combination of baking soda and pickling lime to raise mash pH? I'll leave it to you who have a better grip on all this...

It may go through some intermediary stages of reaction that I'm not showing here, but (unless I summed it up incorrectly) this is a fully balanced equation:

Ca(OH)2 + 2NaHCO3 = CaCO3 + 2NaOH + CO2 + H2O

I'm almost willing to predict that this is the end game of mixing baking soda and calcium hydroxide. If you add each to water and it immediately clouds up, then bingo!

 

[Silver_Is_Money]

I used to use both slaked lime (for Ca without Cl or SO4) and phosphoric.
As soon as I added the acid to the water in the kettle, it clouded and I subsequently decided that meant something was precipitating out, making my ion calculations suspect. For now, I do not tend to use lime & phosphoric at the same time, but haven't felt the need.

Did your water contain bicarbonate (alkalinity)?

 

[dmtaylor]

I have finally used it in my last batch of London brown ale. Won me a gold medal. So yeah I think it's working! No more mash pH of 5.2. Now I'm aiming for about 5.55. Pickling lime gets me there. I think I used like 1/8 or 1/4 teaspoon or something like that. I too have WAY more from one bag of the stuff from the grocery than I will ever need.

 

[Silver_Is_Money]

I have finally used it in my last batch of London brown ale. Won me a gold medal. So yeah I think it's working! No more mash pH of 5.2. Now I'm aiming for about 5.55. Pickling lime gets me there. I think I used like 1/8 or 1/4 teaspoon or something like that. I too have WAY more from one bag of the stuff from the grocery than I will ever need.

Nice to hear! Did you use only Ca(OH)2, or did you also add some NaHCO3? Did you get a cloudy precipitate (which is a clear sign of CaCO3)?

 

[Robert65]

I too have WAY more from one bag of the stuff from the grocery than I will ever need.

A bag of pickling lime from Walmart is cheap, and that is good, because it may not in fact last a lifetime. It can, over time, revert to (useless!) chalk. I test mine periodically to be sure it is still all calcium hydroxide. Simple: place a couple of drops of acid (lactic or whatever you use for brewing) on a sample of your pickling lime. If it bubbles, it contains some calcium carbonate. If you keep it well sealed up against air and moisture this risk is reduced.

 

[Silver_Is_Money]

My calcium chloride prills dropped out a white precipitate after I liquefied them in distilled water. I've since heard this is rather common. It indicates that some percentage of CaCl2 is actually CaCO3 contamination. And it also indicates low quality control.

 

[Robert65]

My calcium chloride prills dropped out a white precipitate after I liquefied them in distilled water. I've since heard this is rather common. It indicates that some percentage of CaCl2 is actually CaCO3 contamination. And it also indicates low quality control.

I also test calcium chloride when I get it, to check purity, and periodically, to ensure that the weight is still accurate (not taking on more water.) I add some to distilled water, and then use a calcium test kit to check that it's making the expected contribution. I've never seen it precipitate out, always dissolves into a clear solution. But we do have to do our own QC on any "food grade" product, those standards are limited in their intentions. But "food grade" and due diligence is cheaper than reagent grade!

 

[Silver_Is_Money]

I'm no longer advising that a combination of Ca(OH)2 and baking soda be used. When starting with distilled or good RO, use only one or the other, but not both. Otherwise you will see the CaCO3 clouding up the water and eventually dropping out. And you will not know the analytical composition of what remains.

 

[dmtaylor]

Nice to hear! Did you use only Ca(OH)2, or did you also add some NaHCO3? Did you get a cloudy precipitate (which is a clear sign of CaCO3)?

Tap water plus Ca(OH)2. I didn't really notice clouding, but probably wasn't paying too close attention either.

 

[Robert65]

Tap water plus Ca(OH)2. I didn't really notice clouding, but probably wasn't paying too close attention either.

Well, empirically you now know how to hit your pH, and ultimately produce the desired beer, whether all of the salts you add are going into solution or not, which is all that really matters in the end.

 

[Robert65]

Little experiment. 400 mL RO. Pickling lime and baking soda in arbitrary quantities that seem a bit large, but within the realm of what might be used in brewing, 0.078g each. First picture is immediately after dissolving pickling lime. Third picture is after stirring in the baking soda. But the second picture is after letting the lime solution stand for a few minutes while I weighed the soda. Not sure how well you can see it, but it is already a bit cloudy, though the lime does not bubble when tested with acid and initially dissolved clear. Something is already happening. Now, this RO was run for brewing a couple of weeks ago and has been sitting in a mostly empty jug since. So it has probably taken up atmospheric CO2 as carbonic acid, and likely had a minuscule amount of native bicarbonate left after RO treatment (measured 5ppm TDS at the time.) Not sure what this means, but interesting.

EDIT: Does suggest that if you are adding calcium chloride or gypsum which may contain trace contaminants, this could affect the amount of calcium you get. But I wonder if the effect on mash pH would even be measurable? What would be the delta RA?

EDIT again: I have never seen this premature cloudiness in actual brewing water, even though I treat my RO in the MT the night before brewing. It's always stayed crystal clear overnight. Really wonder what happened in picture 2.

 

[Silver_Is_Money]

I also test calcium chloride when I get it, to check purity, and periodically, to ensure that the weight is still accurate (not taking on more water.) I add some to distilled water, and then use a calcium test kit to check that it's making the expected contribution. I've never seen it precipitate out, always dissolves into a clear solution. But we do have to do our own QC on any "food grade" product, those standards are limited in their intentions. But "food grade" and due diligence is cheaper than reagent grade!

Would that be a Salifert Test Kit?

 

[Silver_Is_Money]

Little experiment. 400 mL RO. Pickling lime and baking soda in arbitrary quantities that seem a bit large, but within the realm of what might be used in brewing, 0.078g each. First picture is immediately after dissolving pickling lime. Third picture is after stirring in the baking soda. But the second picture is after letting the lime solution stand for a few minutes while I weighed the soda. Not sure how well you can see it, but it is already a bit cloudy, though the lime does not bubble when tested with acid and initially dissolved clear. Something is already happening. Now, this RO was run for brewing a couple of weeks ago and has been sitting in a mostly empty jug since. So it has probably taken up atmospheric CO2 as carbonic acid, and likely had a minuscule amount of native bicarbonate left after RO treatment (measured 5ppm TDS at the time.) Not sure what this means, but interesting.

EDIT: Does suggest that if you are adding calcium chloride or gypsum which may contain trace contaminants, this could affect the amount of calcium you get. But I wonder if the effect on mash pH would even be measurable? What would be the delta RA?

EDIT again: I have never seen this premature cloudiness in actual brewing water, even though I treat my RO in the MT the night before brewing. It's always stayed crystal clear overnight. Really wonder what happened in picture 2. View attachment 648224View attachment 648225View attachment 648226

My bet is that in image #2 you are seeing:
CA(OH)2 + CO2 = CACO3 + H2O

Image #3 definitely shows that combining baking soda (or bicarbonate alkalinity) and slaked lime is not a good idea.

 

[Silver_Is_Money]

I've just completed modifying all of my recipes to mash at a target of pH 5.6, and to (where needed) add either baking soda or slaked lime (but never both together) to raise the mash pH to 5.6. I build from a base of RO water that tests at only a few ppm TDS.

 

[Robert65]

Would that be a Salifert Test Kit?

Yes.

 

[balrog]

I'm no longer advising that a combination of Ca(OH)2 and baking soda be used. When starting with distilled or good RO, use only one or the other, but not both. Otherwise you will see the CaCO3 clouding up the water and eventually dropping out. And you will not know the analytical composition of what remains.

Did your water contain bicarbonate (alkalinity)?

I do have nearly RO water from tap
I do not keep that good a set of notes to be able to say when heating strike water clouded
I did have 3 batches during the time I remember thinking it odd to see cloudiness, and they all had baking soda, pickling lime, as well as phosphoric acid. I have changed my procedure since then to not have lime/baksoda with phosphoric together.

 

[enkamania]

My bet is that in image #2 you are seeing:
CA(OH)2 + CO2 = CACO3 + H2O

Image #3 definitely shows that combining baking soda (or bicarbonate alkalinity) and slaked lime is not a good idea.

Looks like I have to change my plan for Friday.

 

[Silver_Is_Money]

The "intermediate step" I eluded to in post #9 above is:

2NaHCO3 (when added to water, and given time) = Na2CO3 + CO2 + H2O

And from there you get:

Na2CO3 + Ca(OH)2 = CaCO3 + 2NaOH

Or you can just look at it as posted in #9 above.

 

[dmtaylor]

All we need now is to add fish and we'll have lutefisk.

 

[Silver_Is_Money]

All we need now is to add fish and we'll have lutefisk.

How many mEq/Kg of alkalinity does that have?

Actually, having never tried it, I shouldn't be knocking it. Seems quite popular up around Minnesota.

 

[Robert65]

Seems quite popular up around Minnesota.

Where they preserve the fish so they can eat it all winter, and then spend all winter fishing through a little hole anyway... [emoji848] Oh wait, fishing year round = drinking beer year round. Got it now.

 

[MaxStout]

All we need now is to add fish and we'll have lutefisk.

Or make pretzels.

 

[Silver_Is_Money]

My calcium chloride prills dropped out a white precipitate after I liquefied them in distilled water. I've since heard this is rather common. It indicates that some percentage of CaCl2 is actually CaCO3 contamination. And it also indicates low quality control.

In rethinking this, the precipitate I noticed was only seen days to weeks later. The CaCl2 contaminant is therefore more likely to be Ca(OH)2, which with use and exposure to CO2 drops out as CaCO3. If CaCl2 prills are mixed with water and the pH of the water rises then this is indeed the case. And to think that people add CaCl2 to lower pH. This might be just another reason why it's so hard to get mash pH prediction software to hit things right.

Ca(OH)2 + 2HCl = CaCl2 +2H2O (or alternately, CaCl2.2H2O, which is the dihydrate state of CaCl2)

My guess is that the above balanced chemical formula is representative of how CaCL2 is commercially produced. And if the mix is a bit off and there is more Ca(OH)2 in the reaction vessel than need be, then the remaining Ca(OH)2 stays behind as a contaminant when the mix is evaporated to dryness to form prills.

 

[dmtaylor]

In rethinking this, the precipitate I noticed was only seen days to weeks later. The CaCl2 contaminant is therefore more likely to be Ca(OH)2, which with use and exposure to CO2 drops out as CaCO3. If CaCl2 prills are mixed with water and the pH of the water rises then this is indeed the case. And to think that people add CaCl2 to lower pH. This might be just another reason why it's so hard to get mash pH prediction software to hit things right.

But CaCl2 will still lower pH in the mash. The wee bit of Ca(OH)2 or CaCO3 that may form in the water while it heats up shouldn't make much difference unless it's really bad quality CaCl2. When used for the mash, the remaining dissociated Ca++ (probably like >95% of it) will immediately be available to react with phosphates and whatever else in the mash to effectively lower pH. Close enough for me, maybe not for you but at this point I question whether you're just overthinking all this. We're not really dealing with plain water here, we're ultimately always dealing with a mash.

EDIT: And hopefully, if CaCl2 contains impurities for one of us, hopefully it's all of a consistent enough quality for all of us that it doesn't affect a pH model too terribly for any of us. As long as we're consistent in our inconsistency or whatever.