View source for Building brewing water with dissolved chalk

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In order to build brewing water with temporary hardness, chalk (CaCO<sub>3</sub>) is needed. Chalk, however, is not very soluble in water and most of what is added remains suspended or settles to the bottom. This article describes how that chalk can be dissolved if that is desired. 

= About chalk and the carbonate system = 

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''If you do not want to be bothered with the chemistry behind dissolving chalk you can safely skip ahead.''

Chalk, also known as calcium carbonate (CaCO<sub>3</sub>), is a salt of carbonic acid. It's ions are calcium (Ca<sup>2+</sup>) and carbonate (CO<sub>3</sub><sup>2-</sup>). If we know one thing about chalk in brewing it is that it doesn't dissolve very well in water. This is because the solubility product between calcium and carbonate is a very small number. Chemists write it like this:

:[Ca<sup>2+</sup>]*[CO<sub>3</sub><sup>2-</sup>] ≤ K<sub>sp</sub> = 3.36×10<sup>-9</sup>

The names in brackets ([]) stand for the concentration of the particular ions. What that formula means is that water can only hold calcium and carbonate in solution as long as the product of their concentrations is less or equal to 3.36×10<sup>-9</sup>. If the product is greater calcium carbonate will precipitate until enough calcium and carbonate have been removed from the water to satisfy the equation. 

But how can we get calcium carbonate dissolved in water? We do this by lowering the carbonate concentration. To understand this we have to take a look at carbonic acid and the carbonate system. Carbonic acid is formed when carbon dioxide is dissolved in water. Not all the dissolved carbon dioxide (CO<sub>2</sub>) will form carbonic acid but some does according to this chemical equation:

:CO<sub>2</sub> + H<sub>2</sub>O ↔ H<sub>2</sub>CO<sub>3</sub>

[[Image:Carbonic_acid_to_carbonate.gif|frame|right|'''Figure 1''' - The transformation of carbonic acid to carbonate over bicarbonate and back to carbonic acid]] 

The "↔" means that there is a constant back and forth between both sides. I.e. carbonic acid is formed and falls apart into CO<sub>2</sub> and water constantly. But on average there is always a portion of CO<sub>2</sub> bound as carbonic acid. The interesting thing about carbonic acid is that it is a weak acid that can loose up to two protons (H<sup>+</sup>). We looked at weak acids in [[An_Overview_of_pH#Weak_acids_and_bases|Weak Acids and Bases]] where we saw that the loss or acceptance of protons depends on the pH of the environment. When a carbonic acid molecule looses one proton it becomes a bicarbonate ion:

:H<sub>2</sub>CO<sub>3</sub> ↔ HCO<sub>3</sub><sup>-</sup> + H<sup>+</sup>

This process is reversible. I.e. if a bicarbonate ion picks up a proton it becomes carbonic acid. 

When a bicarbonate ion looses a proton it becomes carbonate:

:HCO<sub>3</sub><sup>-</sup> ↔ CO<sub>3</sub><sup>2-</sup> + H<sup>+</sup>

That process is illustrated in Figure 1. As mentioned earlier, the relative concentration of these 3 carbonate species, as they are called, [deLange] depends on pH since it is pH which determines how much H<sup>+</sup> is available. The more there is the more carbonic acid will exist and the less there is the more carbonate will be present. Figure 2 plots these relative concentration as % of the total carbonic acid + bicarbonate + carbonate amount over a pH range of 0 to 14.

[[Image:Carbonic_acid_dissociation.gif|frame|right|'''Figure 2''' - The relative concentration of carbonic acid, bicarbonate and carbonate based on the pH of the solution]]

But how does this tie into dissolving chalk? Well, earlier we saw that we need to lower the concentration of carbonate in order to allow chalk to dissolve. This is done by converting the carbonate, which is already in the water (either present or from a small amount of chalk that already dissolved or part of the water's alkalinity) into bicarbonate. This process consumes protons. It is these protons that we have to supply to dissolve more chalk than the water can hold without any help.

To get these protons we need to add an acid. After all that's what an acid does, it adds protons and as a result lowers the pH of a solution. For acids we have basically two choices carbonic acid or other acids. Let's have a look at what happens when we add an acid other than carbonic acid. Hydrochloric acid for example which, in water, falls apart into a proton (H<sup>+</sup>) and a chloride ion (Cl<sup>-</sup>):

:2H<sup>+</sup> + 2Cl<sup>-</sup> + CaCO<sub>3</sub> → H<sub>2</sub>O + CO<sub>2</sub> + 2Cl<sup>-</sup> + Ca<sup>2+</sup>

The result of the reaction between hydrochloric acid and the chalk is water, CO<sub>2</sub> (which escapes), Chloride and calcium. We could have done this water modification much simpler by just adding calcium chloride as a salt. One of the reasons why we want to dissolve chalk in water is increasing its alkalinity which is caused by bicarbonate and carbonate ions. If the bicarbonate or carbonate content of the water doesn't change its alkalinity won't change either. 

Acids other than carbonic acid don't work. So what about carbonic acid. When carbonic acids becomes bicarbonate it looses a proton which can convert the carbonate ions added by the chalk to bicarbonate. Here is what this looks like:

:H<sub>2</sub>CO<sub>3</sub> → HCO<sub>3</sub><sup>-</sup> + H<sup>+</sup>
:HCO<sub>3</sub><sup>-</sup> + H<sup>+</sup> + CaCO<sub>3</sub> → Ca<sup>2+</sup> + 2HCO<sub>3</sub><sup>-</sup>

The chalk molecule fell apart into a calcium ion and a bicarbonate ion. The other bicarbonate ion comes from the carbonic acids which lost a proton. Because this proton was lost to dissolving chalk we can say that the one molecule of chalk was also responsible for the creation of that bicarbonate ion. In the end one molecule of chalk, when dissolved with carbonic acid, creates one calcium ion and two bicarbonate ions.

[[Image:CaCO3_solubility.gif|frame|right|'''Figure 3''' - The CO2 pressure needed to dissolve a certain concentration of chalk in water. Note that both the x and the y axis are labeled with a logarithmic scale (data source [http://en.wikipedia.org/wiki/Calcium_carbonate#Solubility Wikipedia - Calcium Carbonate Solubility]]]

Getting carbonic acid into water is as easy as adding CO<sub>2</sub> to it. We saw earlier how some of the CO<sub>2</sub> dissolved in water forms carbonic acid. To dissolve a substantial amount of chalk, however, much more CO<sub>2</sub> than the atmosphere is providing is needed. From carbonating beer we know that the CO<sub>2</sub> content of beer depends on its temperature and head space CO<sub>2</sub> pressure. The same is true for water: the higher the CO<sub>2</sub> pressure the higher the amount of dissolved CO<sub>2</sub>. Using a number of chemical equilibrium equations the relation between the CO<sub>2</sub> pressure over water can be correlated with the amount of chalk that can be dissolved in that water [deLange]. Figure 3 shows how the CO<sub>2</sub> pressure correlates with the amount of chalk water can hold. Note that both the x and the y axis are logarithmic and that the seemingly linear relationship is actually close to a cubic function. In practical terms that means that dissolving more than ~800 ppm CaCO<sub>3</sub> in water is unpractical since it requires a substantial amount of CO2 pressure which we don't generally have access to. This needs to be kept in mind for the proposed procedure of dissolving chalk which initially dissolves the chalk in a volume smaller than the actual brewing water volume. 

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=Why dissolve chalk=

[[Image:Chart_pH_over_chalk_concentration.gif|frame|right|'''Figure 4''' - The mash pH for mashes made with distilled water and various levels of chalk addition. One set of experiments dissolved the chalk with CO2 while the other set left the chalk undissolved]]

Now that we know how to dissolve chalk in water we need to ask us why to spend this effort? Especially since the mash will provide an acid environment that should dissolve any chalk that wasn't dissolved in the water. At mash pH (~5.5) less than 0.00005 % of the carbo species are carbonate.

For some reason, however, [http://braukaiser.com/documents/effect_of_water_and_grist_on_mash_pH.pdf mash pH experiments] conducted with suspended and dissolved chalk showed that suspended, i.e. undissolved chalk, is limited in its ability to raise the mash pH. This is shown in Figure 4. The most dissolved chalk was to be able to raise the mash pH was about 0.2 units. This was the base for both Pilsner and Munich malt and does not seem to depend on the acidity of the malt. Dissolved chalk, on the other hand, showed a nice fairly linear relationship between chalk concentration and pH.

I tested this in side-by-side where I brewed a [http://braukaiser.com/lifetype2/index.php?op=ViewArticle&articleId=132&blogId=1 Schwarzbier] with water that had suspended chalk and water that had dissolved chalk. Based on the mash pH research I knew that, when dissolving it, I could use only half the chalk and would still get the same mash pH. That was validated in the experiment.

And if you want to brew according to the German ''Reihnheitsgebot'' you will have to dissolve the chalk, since the addition of salts in the mash or the use of water that doesn't meet drinking water standards is not allowed.

All this being said, dissolving chalk to build brewing water should be considered a technique for the advanced brewer. I have brewed many excellent beers by adding chalk directly to the mash or by not dissolving it in brewing water and just recently started to dissolve the chalk for my brewing water. as with many other techniques, a brewer will have to see for himself if it is worth the additional effort.

=How to dissolve chalk using carbonated water= 

To brew an upcoming batch of [[Kaiser Alt]] I want to mimic the water profile of the city of Düsseldorf where most of the Alts are brewed in Germany. One of the most well known Alt breweries there, ''Zum Uereige'', does not use city water. According to this [http://www.dooyoo.de/bier/uerige-alt/296389/ on-line statement] the brewing water for that brewery comes from their own deep well. The Düsseldorf public water however is essentially cleaned and treated water from the river ''Rhein''.

The current water report for Düsseldorf is available on-line: [http://www.swd-ag.de/privatkunden/wasser/trinkwasseranalyse.php Water report Düsseldorf] and upon calculating the residual alkalinity for that water one finds 108 ppm as CaCO<sub>3</sub> or 6.1 dH (dH is German Hardness). Such a residual alkalinity is well suited for brewing a dark amber beer like an Alt and I assume that most Alt breweries use city water without any water treatment.

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= The process step by step =

== Calculating the salt additions==

In order to calculate the addition of dissolved chalk correctlt you need a brewing water spreadsheet that supports dissolved chalk. Most of the available spreadsheets don't since they assume that each chalk molecule contributes only one bicarbonate ion. My water spread sheer, [http://braukaiser.com/documents/Kaiser_water_calculator.xls Kaiser_water_calculator.xls] supports both dissolved and undissolved chalk in brewing water. In general, when dissolving chalk only half the amount of chalk is needed to build waters with the same residual alkalinity.

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===Step 1===

On the "advanced" work sheet enter the starting water profile. I use reverse osmosis water which is very soft water and for which I have a water analysis. If you are using reverse osmosis water and don't have an analysis simply leave the fields for the starting water profile 0. Due to the low ion concentration compared to the final ion concentration of the brewing water the error will be minimal and to small to matter. 

[[Image:Dissolving_chalk_base_water.gif|center]]

===Step 2===

With an eye on the resulting water profile (section below or right most column) enter concentrations of brewing salts. Rather then entering the final amounts of salts needed and considering the volume of water that will be treated salts are entered as ppm (equivalent to mg/l). This makes the "water recipe" independent of the volume of water that needs to be treated. Once the water volume is known the spreadsheet calculates the required amounts of salts and displays them in a later section.

[[Image:Dissolving_chalk_salts.gif|center]]

===Step 3===

Check the resulting water profile. If you compare this profile to the Düsseldorf water report that was referenced earlier you'll find it to be a very close match. This is not always possible since we brewers don't have all the combinations of cations (positive ions) and anions (negative ions), i.e. salts, available to us. In those cases focus on matching the residual alkalinity and calcium, sulfate and chloride content of the water.

[[Image:Dissolving_chalk_resulting_water.gif|center]]

===Step 4===

Now it is time to enter the total water volume and the strike water volume. The spreadsheet will calculate the sparge water volume from these numbers. You should also enter the grist weight which will allow for the the calculation of the mash thickness. The mash thickness is important for making a mash pH estimation from the beer color and the residual alkalinity since mash thickness determines how much pH shift a given residual alkalinity can cause.

Once you enter the beer color in SRM and how much of the specialty malts are roasted malts an estimate of the mash pH can be made. In this case it is 5.5 and well within the acceptable range of 5.2 - 5.7. The lower right hand corner of the mash pH section also shows estimates for the residual alkaliniies that would be necessary to reach a mash pH of 5.2, 5.4 and 5.6 respectively. The large range from ~-100 to ~200 ppm as CaCO<sub>3</sub> shows how wide of a range of brewing waters could be used to brew a 20 SRM beer with a mash thickness of 4 l/kg (~2 qt/lb)

[[Image:Dissolving_chalk_pH_estimate.gif|center]]

===Step 5===

Using the water volumes specified earlier, the amounts of salts needed can now be calculated. Because I don't have the ability to treat all the water at once I have to treat mash and sparge water separately. But for those able to treat all the needed water at once the amounts needed for the total volume are also given.

Of interest for dissolving chalk is the "dissolving chalk" section. Here you enter the volume of water that the chalk should be dissolved in and the spreadsheet calculates the CO<sub>2</sub> pressure that is needed to dissolve the chalk. You'll quickly find out that the relation between chalk concentration and pressure is not linear and that the pressure quickly shoots up once you exceed about 750 ppm. As a result this is generally the practial limit and to dissolve more chalk you need to dissolve it in a larger mount of water. 

If possible I like to dissolve my chalk additions in 2 l soda bottles and in this case it will be possible. Using 1.9 l to dissolve the mash water addition and 1.6 l to dissolve the sparge water addition I need only 1.55 bar and 1.38 bar respectively. Those pressure numbers are absolute which means that they include atmospheric pressure and that I have to use >0.55 bar and >0.38 bar as the setting on the regulator. A CO<sub>2</sub> pressure regulator measures the pressure in excess of the atmospheric pressure. 

[[Image:Dissolving_chalk_salt_amounts.gif|center]]

Another practical means of dissolving chalk are corny kegs. They hold a much larger water amount and as a result will require less CO2 pressure to dissolve chalk. Confirming that the chalk has been dissolved is more difficult, since it requires pulling a sample. 

==Building the brewing water==

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| colspan="2" align="center" | [[Image:Dissolving_chalk_needed_items.jpg|frame|'''Figure 5''' - These are the needed items: 2 liter soda bottles, carbonator cap, funnel, weighed salts and water. I mark the bottles with batch number and if the salt water is for mash or sparge water treatment]]
| [[Image:Dissolving_chalk_add_salts.jpg|frame|'''Figure 6''' - Set the funnel into the bottle and dump the salts into the funnel]]

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| valign="top" | [[Image:Dissolving_chalk_add_water.jpg|frame|'''Figure 7''' - Add water. The water will wash any salts, that got stuck in the funnel, into the bottle. Even if close to 2 l water are needed, don't fill the bottle all the way yet. Leaving a headspace makes carbonating the water easier]]
| valign="top" | [[Image:Dissolving_chalk_remove_air.jpg|frame|'''Figure 8''' - Add the carbonator cap, press it open and squeeze out the air. Removing all the air and replacing it with CO<sub>2</sub> will make carbonation much more effective since the headspace will be completely filled with CO<sub>2</sub>]]
| valign="top" | [[Image:Dissolving_chalk_hook_up_CO2.jpg|frame|'''Figure 9''' - Hook up the CO<sub>2</sub> bottle via a regulator and shake the bottle to carbonate the water. Keep shaking until no more CO<sub>2</sub> is flowing into the bottle. Once carbonated top off with more water if necessary and carbonate again. Repeat the process for the other bottle(s)]]
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| valign="top" | [[Image:Dissolving_chalk_clear_water.jpg|frame|'''Figure 10''' - After standing for a while the water should clear. A light haze may remain but there should not be a large amount of chalk sediment that turns the water cloudy once shaken up. If the water doesn't clear divide the contents evenly between two bottles, add water and carbonate again. Maybe the chalk concentration was to high to be dissolved by the used CO<sub>2</sub> pressure]]
| valign="top" | [[Image:Dissolving_chalk_strike_water.jpg|frame|'''Figure 11''' - The day before or on brewday I add the salt brine to the remaining reverse water that to get the desired volume of strike water and then later the same for the sparge water. While the CO<sub>2</sub> is now able to escape from the water, the chalk has been diluted so far that it will take a very long time for enough CO<sub>2</sub> to escape that chalk will start to precipitate again. You'll see more CO2 escape while heating the water but since you are not boiling it enough CO<sub>2</sub> will remain in solution to keep the chalk dissolved]]
| valign="top" | [[Image:Dissolving_chalk_pH.jpg|frame|'''Figure 12''' - If you test the pH of the water you will find that it is lower than what you expect from fairly alkaline water. This low pH is the result of an excess of carbonic acids but that pH does not affect the alkalinity of the water. When I tested the hardness and alkalinity with a [[At_home_water_testing|GH&KH test kit]] I found a hardness of 15 dH (270 ppm as CaCO<sub>3</sub>) and an alkalinity of 10 dH (180 ppm as CaCO<sub>3</sub>).]]
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An alternative vessel for dissolving chalk are corny kegs. Their advantage is that they can hold larger volumes of water and as a result the chalk concentration can be much lower. With that the needed CO<sub>2</sub> pressure is also much lower. The main disadvantage is that visual inspection of the process is difficult since it requires drawing a sample. But due the lower CO<sub>2</sub> pressure that will be needed it is much more likely that all chalk gets dissolved in the larger water volume in a corny keg compared to the 2 liter limit in a soda bottle.


=References=

:[deLange] A. J. deLange, [http://ajdel.wetnewf.org:81/Brewing_articles/Cerevesia/Final_galley Alkalinity, Hardness, Residual Alkalinity and Malt Phosphate: Factors in the Establishment of Mash pH], 2004


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