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Message |
    
George Schmidt
Advanced Member
Username: Gschmidt
Post Number: 509
Registered: 08-2004
| | Posted on Saturday, May 14, 2005 - 04:27 am: | 
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I've seen slaked lime mentioned a couple times as a way to remove temporary hardness. I can't really find instructions as such, though. I remember a couple threads in the past on the subject, but I can't find them. I've got a buttload of temp hardness that I've been boiling out, but it seems that lime would be a lot more economical of energy than boiling 10-12 gallons the night before then having to heat it again the next day. It doesn't seem like many people do this, though. Why is this? How hard is it to get this working?
Be wary of strong drink. It can make you shoot at tax collectors -- and miss. ~~Robert A. Heinlein: The Notebooks of Lazarus Long
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Bill Pierce
Moderator
Username: Billpierce
Post Number: 3083
Registered: 01-2002
| | Posted on Saturday, May 14, 2005 - 11:41 am: |
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The best description of the process is on Hubert Hanghofer's site: http://www.netbeer.org/english/lime.htm
Personally, I think boiling the water to remove temporary hardness is simpler. |
Dan Listermann
Senior Member
Username: Listermann
Post Number: 1098
Registered: 03-2004
| | Posted on Saturday, May 14, 2005 - 02:22 pm: |
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What is the best form of slaked lime to buy?
Dan
Listermann Mfg.,Co. www.listermann.com
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ERD
New Member
Username: Rico
Post Number: 17
Registered: 11-2004
| | Posted on Saturday, May 14, 2005 - 03:03 pm: |
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Good article, Bill. Thanks. |
Tom Meier
Member
Username: Brewdawg96
Post Number: 247
Registered: 03-2003
| | Posted on Saturday, May 14, 2005 - 03:05 pm: |
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grocery store pickling lime which is Ca(OH)2, i.e. slaked lime |
George Schmidt
Advanced Member
Username: Gschmidt
Post Number: 510
Registered: 08-2004
| | Posted on Saturday, May 14, 2005 - 03:49 pm: |
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That link is perfect. Thanks, Bill.
Be wary of strong drink. It can make you shoot at tax collectors -- and miss. ~~Robert A. Heinlein: The Notebooks of Lazarus Long
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George Schmidt
Advanced Member
Username: Gschmidt
Post Number: 512
Registered: 08-2004
| | Posted on Tuesday, May 17, 2005 - 05:56 pm: |
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Is hydrated lime from garden supply stores the same as slaked lime?
Be wary of strong drink. It can make you shoot at tax collectors -- and miss. ~~Robert A. Heinlein: The Notebooks of Lazarus Long
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Paul Edwards
Advanced Member
Username: Pedwards
Post Number: 705
Registered: 03-2003
| | Posted on Tuesday, May 17, 2005 - 06:11 pm: |
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Hydrated lime is the same, chemically, but...
Get the pickling lime from the grocery store. It is food grade, while the hydrated lime from the garden supply is for spreading on soil, and is likely not as pure. |
Guy C
Member
Username: Ipaguy
Post Number: 240
Registered: 09-2003
| | Posted on Tuesday, May 17, 2005 - 07:07 pm: |
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Could someone post an example for this treatment? Does anyone have the lime addition in spreadsheet form that shows the pre- and post-effects of the treatment? |
bierslayer
New Member
Username: Bierslayer
Post Number: 22
Registered: 04-2004
| | Posted on Wednesday, May 18, 2005 - 04:35 am: |
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Guy C,
I measured my tap water's alkalinity to see if it agreed with the city water report of 325 (yikes!). I titrated 100 mL of tap water with 0.1N HCl until it reached pH 4.3. The number of mL of acid is equal to the mEq/L of alkalinity. I measured 6.5 mEq/L which is equal to 325 ppm as the city said. (6.5 mEq/L x 60 to get ppm alkalinity) For every mEq of alkalinity, two mEq of calcium is required to neutralize the bicarbonate. So 6.5 mEq/L x 2 = 13 mEq/L of Ca++ required in my case. Multiply that by 20 to get ppm required. (13 mEq/L calcium x 30 =260 ppm calcium needed!)
Take a sip of beer now. Still with me?
Using one of Hubert's formulas, multiply your alkalinity in ppm by 0.74/1000 to get the calculated amount of slaked lime required. (325 ppm Alkalinity x .074/1000 =0.24 g/L of slaked lime required) I rounded up to 0.25 g/L. This is the amount of slaked lime I would need to use. (Also, 0.25 g = 250 mg for easier calculations below)
Out of curiosity, I wanted to know how much calcium I had available to complex with the bicarbonates since I know this treatment consumes calcium ions. Given that 40 ppm of Ca++ is equal to 74 mg/L Ca++, 250 mg/L slaked lime divided by 74 mg/L = 3.378 x 40 ppm Ca++ = 135 ppm Ca++ added to my water via slaked lime. Then, 135 ppm divided by 20 = 6.76 mEq Ca++ due to the lime addition.
Take another sip of beer and stay with me now.
We calculated above that we needed 13 mEq/L of Ca++ to neutralize the alkalinity. And we determined that the lime addition will add 6.76 mEq/L of Ca++. That means we still need 6.24 mEq/L of Ca++ (13 mEq/L - 6.76 mEq/L = 6.24 mEq/L needed) My city water report says that the tap water contains 70 ppm of Ca++ which we can divide by 20 to get 3.5 mEq/L that's already in my water. Subtract that from our Ca++ requirement and we're left with a deficit of 2.74 mEq/L Ca++ (13 mEq/L needed - 6.76 mEq/L from lime - 3.5 mEq/L from tap = 2.74 mEq/L Ca++ still needed.) Multiply 2.74 mEq/L x 20 = 54.8 ppm of Ca++ needed.
Now, I used Hubert's method of adding all of the lime to half of the water and then adding the rest of the water to complete the reaction. When I did this my pH of the treated water was 9.9 and I measured the alkalinity as 2.6 mEq/L (or 156 ppm) This shows me that even though I added the calculated amount of lime, I still have some bicarbonate that won't precipitate out of solution because I ran out of calcium ions.
We calculated above that we still needed 2.74 mEq/L of Ca++ and we demonstrated that as well. So this time, I wanted to figure out how much CaSO4 or CaCl2 I would need to add to make up the deficit. (The following numbers are conversions of readily available data that is usually presented in g/gallon instead of g/L)
For CaSO4, 1 gram contributes 234 ppm of Ca++/L. We need about 55 ppm which means we need: 55 ppm divided by 234 ppm = 0.234 g/L of CaSO4. *Remember that you'd also be adding alot of sulfate ions as well *
For CaCl2, 1 gram contributes 275 ppm of Ca++/L. We need about 55 ppm which means we need: 55 ppm divided by 275 ppm = 0.200 g/L of CaCl2. *Remember that you'd also be adding alot of chloride ions as well.*
Slam your beer and crack open another one cuz we're just about finished!
I decided I wanted to have some Ca++ left over after my treatments to assist with dropping mash pH and making enzymes happy so I did a treatment with 0.5 g/L of CaCl2 which contributes 137.5 ppm/L Ca++. We had a deficit of 55 ppm so this treatment results in 137.5 ppm/L - 55 ppm/L = 82.5 ppm Ca++ left over in my treated water. (This is well above the amount of Ca++ I need but I know I'm going to have some alkalinity left over that I won't be able to remove with lime so I'll need the extra Ca++ in the mash later).
Adding 250 mg/L slaked lime to half of the water along with 500 mg/L CaCl2 and then adding the rest of the water resulted in a final water with a measured pH of 8.3 and measured alkalinity of 0.85 mEq/L which is 51 ppm alkalinity (0.85 mEq/L x 60 = 51)
I haven't actually brewed with treated water but in the future would use a combination of CaSO4 and CaCl2 to keep the anion concentrations down. Also, I really don't need to add so much excess Ca++ but I wanted to see how low I could get alkalinity to drop. |
Tom Meier
Member
Username: Brewdawg96
Post Number: 250
Registered: 03-2003
| | Posted on Wednesday, May 18, 2005 - 04:53 am: |
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>>Could someone post an example for this treatment?
grams / Litre = Alkalinity (ppm CaCO3) x 0,74 / 1000
g/L = 200 * 0.74 / 1000 = 0.148 g/L
So for 20L, add 2.96 grams
think that is right, check the math
There is really no need for a spreadsheet. The pre-effect is you start with whatever your water is... the post-effect is you end up with really low alkalinity water regardless of what your water is. |
Bill Pierce
Moderator
Username: Billpierce
Post Number: 3105
Registered: 01-2002
| | Posted on Wednesday, May 18, 2005 - 11:11 am: |
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Very good, bierslayer. I'm impressed with your logic. |
Guy C
Member
Username: Ipaguy
Post Number: 241
Registered: 09-2003
| | Posted on Wednesday, May 18, 2005 - 05:05 pm: |
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Bierslayer, thanks for the detailed example. For some reason boiling and dilution now seems more attractive than ever. 
Tom, this treatment depletes calcium in water where much of the hardness comes from hydrogencarbonates, right? All those calculations bierslayer did above seem to indicate that, and a spreadsheet application would seem to be a useful option for determining the amount of calcium depleted.
(Message edited by ipaguy on May 19, 2005) |
bierslayer
New Member
Username: Bierslayer
Post Number: 23
Registered: 04-2004
| | Posted on Thursday, May 19, 2005 - 02:08 am: |
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Thanks Bill.
Guy C... I would agree that dilution with RO is easiest but with large volumes of water it gets to be a bit of work unless you've got a system at home. The slaked lime treatment is really simple in my opinion. Figuring everything out is a little intimidating but once you've got the calculations done it's not too tricky. Maybe I'll work on a spreadsheet and see what I can come up with
Incidentally, I boiled my tap water and the alkalinity post-boil was 2.8 mEq/L (168 ppm) and pH was 9.1. So my starting water was at an alkalinity of 6.5 mEq/L - 2.8 mEq/L remaining = 3.7 mEq/L removed. That tells me that all of my Ca++ was probably depleted during the boil but still some additonal bicarbonate was precipitated out. If you choose to pre-boil, I'd consider adding more calcium to make it more effective. I'll check this out sometime and report back. |
bierslayer
New Member
Username: Bierslayer
Post Number: 24
Registered: 04-2004
| | Posted on Friday, May 20, 2005 - 04:18 am: |
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Please note that I made some stupid math errors in the original post and have corrected them below. I tried to edit the post above but it won't let me. Sorry for any confusion and for doubling up on the extra long post.
I measured my tap water's alkalinity to see if it agreed with the city water report of 325 (yikes!). I titrated 100 mL of tap water with 0.1N HCl until it reached pH 4.3. The number of mL of acid is equal to the mEq/L of alkalinity. I measured 6.5 mEq/L which is equal to 390 ppm as the city said. (6.5 mEq/L x 60 to get ppm alkalinity) For every mEq of alkalinity, two mEq of calcium is required to neutralize the bicarbonate. So 6.5 mEq/L x 2 = 13 mEq/L of Ca++ required in my case. Multiply that by 20 to get ppm required. (13 mEq/L calcium x 30 =260 ppm calcium needed!)
Take a sip of beer now. Still with me?
Using one of Hubert's formulas, multiply your alkalinity in ppm by 0.74/1000 to get the calculated amount of slaked lime required. (390 ppm Alkalinity x 0.74/1000 =0.289 g/L of slaked lime required) This is the amount of slaked lime I would need to use. (Also, 0.289 g = 289 mg for easier calculations below) I used 250 mg/L as a starting point for my treatment for no particular reason other than measuring convenience.
Out of curiosity, I wanted to know how much calcium I had available to complex with the bicarbonates since I know this treatment consumes calcium ions. Given that 40 ppm of Ca++ is equal to 74 mg/L Ca++, 250 mg/L slaked lime divided by 74 mg/L = 3.378 x 40 ppm Ca++ = 135 ppm Ca++ added to my water via slaked lime. Then, 135 ppm divided by 20 = 6.76 mEq Ca++ due to the lime addition.
Take another sip of beer and stay with me now.
We calculated above that we needed 13 mEq/L of Ca++ to neutralize the alkalinity. And we determined that the lime addition will add 6.76 mEq/L of Ca++. That means we still need 6.24 mEq/L of Ca++ (13 mEq/L - 6.76 mEq/L = 6.24 mEq/L needed) My city water report says that the tap water contains 70 ppm of Ca++ which we can divide by 20 to get 3.5 mEq/L that's already in my water. Subtract that from our Ca++ requirement and we're left with a deficit of 2.74 mEq/L Ca++ (13 mEq/L needed - 6.76 mEq/L from lime - 3.5 mEq/L from tap = 2.74 mEq/L Ca++ still needed.) Multiply 2.74 mEq/L x 20 = 54.8 ppm of Ca++ needed.
Now, I used Hubert's method of adding all of the lime to half of the water and then adding the rest of the water to complete the reaction. When I did this my pH of the treated water was 9.9 and I measured the alkalinity as 2.6 mEq/L (or 156 ppm) This shows me that even though I added the calculated amount of lime, I still have some bicarbonate that won't precipitate out of solution because I ran out of calcium ions.
We calculated above that we still needed 2.74 mEq/L of Ca++ and we demonstrated that as well. So this time, I wanted to figure out how much CaSO4 or CaCl2 I would need to add to make up the deficit. (The following numbers are conversions of readily available data that is usually presented in g/gallon instead of g/L)
For CaSO4, 1 gram contributes 230.4 ppm of Ca++/L. We need about 55 ppm which means we need: 55 ppm divided by 230.4 ppm = 0.239 g/L of CaSO4. *Remember that you'd also be adding alot of sulfate ions as well *
For CaCl2, 1 gram contributes 270.3 ppm of Ca++/L. We need about 55 ppm which means we need: 55 ppm divided by 270.3 ppm = 0.203 g/L of CaCl2. *Remember that you'd also be adding alot of chloride ions as well.*
Slam your beer and crack open another one cuz we're just about finished!
I decided I wanted to have some Ca++ left over after my treatments to assist with dropping mash pH and making enzymes happy so I did a treatment with 0.5 g/L of CaCl2 which contributes 135.15 ppm/L Ca++. We had a deficit of 55 ppm so this treatment results in 135.15 ppm/L - 55 ppm/L = 80.15 ppm Ca++ left over in my treated water. (This is well above the amount of Ca++ I need but I know I'm going to have some alkalinity left over that I won't be able to remove with lime so I'll need the extra Ca++ in the mash later).
Adding 250 mg/L slaked lime to half of the water along with 500 mg/L CaCl2 and then adding the rest of the water resulted in a final water with a measured pH of 8.3 and measured alkalinity of 0.85 mEq/L which is 51 ppm alkalinity (0.85 mEq/L x 60 = 51)
I haven't actually brewed with treated water but in the future would use a combination of CaSO4 and CaCl2 to keep the anion concentrations down. Also, I really don't need to add so much excess Ca++ but I wanted to see how low I could get alkalinity to drop. |
bierslayer
New Member
Username: Bierslayer
Post Number: 25
Registered: 04-2004
| | Posted on Friday, May 20, 2005 - 04:58 am: |
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Hmm....three posts in a row. I guess I'm talking to myself now!
I made a quick spreadsheet that you could use to figure out your lime additions. The colored cells are where you need to enter your relevant water values. Note that the numbers you get are starting points and not absolutes.
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George Schmidt
Advanced Member
Username: Gschmidt
Post Number: 514
Registered: 08-2004
| | Posted on Friday, May 20, 2005 - 03:32 pm: |
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bierslayer, what format is the spreadsheet in?
Be wary of strong drink. It can make you shoot at tax collectors -- and miss. ~~Robert A. Heinlein: The Notebooks of Lazarus Long
|
Geoff Buschur
Advanced Member
Username: Avmech
Post Number: 761
Registered: 06-2004
| | Posted on Friday, May 20, 2005 - 03:52 pm: |
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That didn't work
(Message edited by avmech on May 20, 2005) |
Geoff Buschur
Advanced Member
Username: Avmech
Post Number: 762
Registered: 06-2004
| | Posted on Friday, May 20, 2005 - 03:56 pm: |
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Another try.
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bierslayer
Junior Member
Username: Bierslayer
Post Number: 26
Registered: 04-2004
| | Posted on Friday, May 20, 2005 - 05:01 pm: |
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Thanks Geoff. I forgot to add the .xls to the name of the file. |
Guy C
Member
Username: Ipaguy
Post Number: 244
Registered: 09-2003
| | Posted on Friday, May 20, 2005 - 05:52 pm: |
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Thanks for the spreadsheet! |
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slaked lime calculator