Everything You Ever Wanted To Know About Brewing Water

by: Tom M.

1.0 INTRODUCTION

Although it makes up 90-95% of beer ingredients by weight, the majority of brewers rarely consider what type of water to use. How often have you seen a recipe in Zymurgy that lists what water was used to brew it? Probably never. Most brewers rarely give it much thought, and water is the least appreciated of beer ingredients. At the other extreme, many brewers spend hard earned dollars on spring water because they don’t fully understand how water affects flavor.

This article will cover water chemistry from both a theoretical and a practical 'what does this stuff have to do with beer?' standpoint and hopefully benefit advanced all-grain brewers by providing insight into water chemistry, and how it can affect beer flavor. At the very least, you can print this out and save it for when and if your interest grows and you actually give a hoot about water chemistry. Extract brewers are strongly encouraged to relax, not worry, have a homebrew, and skip to the extract brewing tips section.

Some of the background information here can be found in a home brewing how-to book, such as Dave Miller's 'The Complete Handbook of Homebrewing', Randy Mosher's 'The Homebrewing Companion', or John Palmer's 'How to Brew', my personal favorite.

This basic type information is usually scattered, and for me never gave a complete picture to understand exactly what was going on. This article is my attempt to bring together information from a large number of sources into one stand-alone water chemistry crash course that also pre-calculates what you need to know for brewing with Nashville water.

2.0 IONS IN BREWING

First, what is an ion? Basically it is a mineral that has been dissolved in water into separate parts, called ions. For example, salt is a crystalline solid of the chemical form NaCl until you put it in water, and when dissolved, the Na and Cl separate and attach themselves to the electrically polar water molecules. Dissolved salt is then two separate ions: Na+, and Cl^-. The plus sign refers to sodium's lack of an electron and the negative sign means the chloride has an extra electron.

Water soluble minerals are naturally found in soil deposits. As very soft rainwater makes its way down through the soil, through underground caves, over rocks into streams or aquifers it picks up some of the following dissolved minerals (ions) along the way: calcium (Ca⁺⁺), carbonate (CO3^--), chloride (Cl^-), bicarbonate (HCO3^-), magnesium (Mg⁺⁺), sodium (Na+), sulfate (SO4^-), and trace metals such as iron (Fe) and copper (Cu).

I would recommend contacting your local water authority to obtain a free water report if you believe your water may be questionable for brewing or if you want to check some of the things talked about here. For private wells or springs, you may obtain a full water report from www.wardlab.com. Order Household Mineral Test W-6 ($15), which tests for all brewing related ions and parameters.

Side Note:
The symbol Cl can cause confusion. In water chemistry chloride (Cl^-) is from a salt, and is not the same thing as elemental chlorine (Cl₂), which does not exist in tap water, even when bleach is added. This is another confusion - that bleach puts chlorine in water. Bleach or sodium hypochlorite (NaOCl) added to water creates hypochlorous acid (HOCl). The HOCl partially breaks up into H+ and OCl- in a chemical equilibrium, as all acids do. The percentage of which it breaks up is dependent on pH. HOCl is a strong sanitizer, and OCl- is a strong oxidizer (bleaching agent). OCl- is unstable and can break up into Cl₂ readily. HOCl is also broken down rapidly at low pH levels into H₂O and Cl₂gas, which is why you should never mix acid and bleach. The reason is acids give up lots of H+ ions, which react with HOCl and to yield water and chlorine gas.

3.0 CALCIUM CARBONATE (CaCO₃)

The two biggest players in all-grain brewing water are the ions calcium (Ca⁺⁺) and bicarbonate (HCO₃^-). Top surface runoff dissolves CaCO₃ out of limestone, but can only dissolve enough CaCO₃ to result in about 58 ppm of bicarbonate (HCO₃^-), because this is the equilibrium saturation level for HCO₃^- in water exposed to air. The maximum amount of HCO₃^-possible in water at equilibrium is directly related to the partial pressure of CO₂, or more simply, the % of CO₂ in the air we breathe.

In underground situations, respiring bacteria can increase the CO₂ level to an equilibrium partial pressure of 60 times as much as surface water, and this allows a much higher amount of bicarbonate (HCO₃^-) to be dissolved. For example, the famous brewing water of Burton on Trent is from underground sources, and one source is reported as having over 300 ppm of HCO₃^-.

Equilibrium is analogous to the most sugar that will dissolve in a liquid at a given temperature. The un-dissolved sugar at the bottom is in equilibrium with a saturated sugar solution above it. No more is dissolving, no more is dropping out. Now, if you cool that liquid, it becomes supersaturated with sugar, which means there is more dissolved than should be, so if crystals have a place to form they will, at a very slow rate, but faster when cooled greatly. With carbo (bicarbonate, carbonic, and carbonate) the opposite is true, it is one of the few solutions whose solubility (amount that can be dissolved) goes down as the water temp increases. This is why boilers, coffee makers, and irons get calcium carbonate deposits on their hot surfaces. We will use that to our advantage in Part 2 under the pre-boiling section.

Nashville Metro Water has a HCO₃^- level of 61 ppm, so from this we figure Metro water consists mainly of surface runoff water, which coming from a river it certainly does. Since it is over 58 ppm, it is slightly supersaturated with bicarbonate (HCO₃^-).

4.0 HARDNESS

Hardness is the amount of Ca⁺⁺ and/or Mg⁺⁺ ions, and is usually given in ppm (same as mg/l) or in terms of ppm as CaCO₃.

If a water report gives only total hardness and you want to estimate the calcium and magnesium hardness, assume that 4/5ths of the total hardness is calcium hardness and 1/5th is magnesium hardness. This ratio varies by water source, but Metro water stays fairly constant at this ratio.

There is another type of hardness called temporary hardness, which is simply the amount of hardness you can get rid of by boiling. Boiling causes Ca to react with CO₃ in the water and precipitate out (see section on pre-boiling). Permanent hardness cannot be removed by boiling, hence the name.

Hardness is beneficial to brewing water. Calcium helps to protect mash enzymes, and if depleted, mashes could take much longer, and may not have the proper mix of alpha and beta amylase activity. Many sugar refineries that use amylase and glucosidase enzymes add calcium chloride to soft tap water to obtain over 35 ppm of calcium, which prolongs enzyme life. Calcium also aids in getting a good hot break (coagulation of proteins and polyphenols in the boil). Magnesium hardness is needed in trace amounts for yeast health.

Side Note: Water Softeners and Filters

If you have a household water softener to reduce water hardness, you should bypass it when drawing out water used for brewing. These contain an ion exchange bed with sodium (Na+) ions. The Ca⁺⁺ and Mg⁺⁺ ions are almost all replaced by Na+, which can lead to very high levels of sodium if the water is hard. Also, don't use Brita(tm) filters for brewing water, these have both activated charcoal and ion exchange resin. The only filter you need for brewing water is activated charcoal to remove chlorine compounds. See the section on Removal of Chlorine Compounds in this article

Although the hardness level of water is not detrimental, it is always accompanied by high alkalinity. It is the alkalinity (bicarbonates et. al.) which makes hard water undesirable for brewing. As will be discussed later, water can be extremely hard and still be capable of producing a relatively light colored beer.

5.0 ALKALINITY

Alkalinity of tap water is defined as the amount of acid required to bring a sample of water from its initial pH down to a pH of 4.3. The acid is added slowly, or titrated, and when the solution hits a pH of 4.3 the amount of acid is recorded and converted into the appropriate units of Alkalinity, which are ppm as CaCO₃.

This can be estimated mathematically for tap water assuming all alkalinity comes from CaCO3, and not from nitrates or other anions. The formula to use is:

Alkalinity = 2 x carbonate (CO₃^--) + bicarbonate (HCO₃^-) + OH^- - H⁺

where they are all in the same units. Alkalinity is almost always given in water reports in terms of ppm as CaCO₃. The OH^- and H⁺ terms are due to the dissociation of water itself, and are extremely small, so we can ignore them.

You will notice that the carbonate is multiplied by 2 in the definition of alkalinity. If you are thinking that is because carbonate has two little minuses beside it, you are right. In very simple terms, alkalinity is a measurement of the anions in the water, or put even simpler, adding up all the negatives.

Recall alkalinity is measured by adding acid. Well, the acid is continually added, contributing H+ ions until the pH is 4.3 (the ‘zero’ point). At that point the amount of plus’s added (amount of acid) is figured out and this yields the alkalinity, or number of minuses. That’s all there is to it, in algebraic terms.

One thing of note in regards to alkalinity is that bicarbonate (HCO₃^-), carbonic acid (H₂CO₃), and carbonate (CO₃^--) all exist together in a 3-way relationship that is determined solely by pH of the water they are dissolved in. In normal water of about 7 pH, there is 81% bicarbonate (HCO₃^-), 19% carbonic (H₂CO₃), and 0% carbonate (CO₃^--). For any given pH, the % of bicarbonate, carbonic, and carbonate can be easily determined.

This chart shows the fractional amount (percentage) of each that exist for any given pH

Chart 5.0 - Carbonate System in tap water

Notice from this table that carbonate (CO₃^--) does not exist below a pH of 8.3. As a result Bicarbonate (HCO₃^-) is the source of all alkalinity in tap water of less than 8.3 pH.

Side Note:
It is important to note that, as shown in the above chart of bicarbonate, carbonic, and carbonate, if the pH of the water is lowered (say with acid addition), the alkalinity is decreased as the ‘carbo’ shifts from bicarbonate (HCO₃^-) to carbonic acid (H₂CO₃), which does not contribute to alkalinity. See part 2 for acid additions.

Because the ratio of carbonate and bicarbonate is related to pH, if your water report gives carbonate/ bicarbonate levels and pH, but doesn’t list alkalinity, you can roughly calculate the alkalinity as outlined in section 5.1 below.

5.1 CALCULATING ALKALINITY

Alright so we know what alkalinity is. Remember that solar powered calculator you got free a few years ago? Let’s pull that out of the catch-all drawer and blow the dust off it.

Side Note:
The unit conversions for all minerals shown here are given in an easy to use table form - see Attachment 2

If your water report gives: bicarbonate (HCO₃^-) in ppm (aka mg/l), then you can calculate alkalinity assuming the pH is <8.3 and no carbonate exists. Note that sometimes water profiles erroneously list CO₃^-- when they are actually referring to HCO₃^-. As shown in the table above, CO₃^-- is not present at normal tap water pH. The formula to calculate alkalinity given this information is:

Alkalinity (in ppm as CaCO₃) = (mg HCO₃^-/l ) x (1 mEq/61 mg HCO₃^-) x (50 mg CaCO₃/1 mEq)

or without units shown: Alkalinity = HCO₃^- x 50 / 61

If your water report gives Alkalinity as ‘hardness’, ‘ppm CaCO₃’, or no units, then your calculator can go back in the drawer, you already have your alkalinity in ppm as CaCO₃ (sorry for the let down). If you need or want to, you can calculate HCO₃^- using:

mg/l of HCO₃^- = Alkalinity [in mg CaCO₃/liter] x (1 mEq /50 mg CaCO₃) x
(61 mg HCO₃^-/mEq)

or without units shown: HCO₃^- = Alkalinity x 61 / 50 (this assumes pH of <8.3)

5.2 HOW ALKALINITY AFFECTS BREWING

To give some history, before brewing chemistry was widely understood highly alkaline waters were generally unsuitable for brewing light styles. The water in London was favorable to brewing ESB and Porters, highly alkaline Dublin water was used to brew stouts, and only the very soft water of Pilsen was capable of properly brewing with extremely light kilned malts. The reason dark beers are traditionally brewed with alkaline waters is because roasted grains add acidity to the mash, buffering the alkalinity of water and bringing mash pH down to an acceptable range.

If a pilsener grist were mashed with untreated Dublin water, the result would be a high mash pH, resulting in poor enzyme activity, long mash times, and possible extraction of tannins from the husk. Without modern day knowledge of water chemistry, brewing cities gravitated towards styles which mashed easily and whose flavors blended well with the local water.

6.0 RESIDUAL ALKALINITY

Now that we have calculated and understand basic water chemistry and alkalinity, we can discuss residual alkalinity. If you recall from the discussion on calcium and carbonates, we said that Burton-on-Trent water was very high in bicarbonate (300 ppm). Consequently, it is also high in hardness. Here is the water profile:

Water Profile: Burton on Trent*

Ca⁺⁺	HCO₃^--	Cl^-	Mg⁺⁺	Na⁺	SO₄^-- ^--
295	300	25	45	55	725

*all values listed are in mg/l (same thing as ppm)

We have also discussed that generally, water high in alkalinity (which HCO₃^-contributes to) is only suitable for brewing darker beers. However, in the case of Burton on Trent, brewers have traditionally been able to brew good pale ales and IPAs. The reason for this lies in a property of brewing water called Residual Alkalinity.

Looking at the Burton profile you can see that calcium is a high 295 ppm, and yet is only balanced by 300 ppm of carbonate. Where did all that calcium come from then? Probably CaSO₄, (gypsum) dissolved in the soil. The SO₄^-- levels are sky high at 725 ppm. Remember that when we defined alkalinity, we said it was HCO₃^-+ CO₃^-- (no SO₄^--term), so although gypsum (CaSO₄) adds calcium it does not contribute to alkalinity, but as we will discover later, it does reduce ‘Residual’ Alkalinity.

Residual Alkalinity was studied by a German brewing scientist named Kolbach, who observed that calcium and magnesium reacted with phytin present in the mash to produce hydrogen ions, which effectively lowers the mash pH.

This means that the higher the calcium and magnesium for a given level of alkalinity, the lower the mash pH will be.

6.1 CALCULATING RESIDUAL ALKALINITY

Residual alkalinity can be determined using the known characteristics of the water and a simple formula. For water where the report gives alkalinity, calcium, and magnesium all in ‘hardness’ (ppm as CaCO₃), the formula is

Residual Alkalinity (RA) = Alkalinity - Ca / 3.5 - Mg / 7, where all are in ppm as CaCO₃

The formula for calculating RA where Ca and Mg are given in ppm (not as hardness) is:

RA = Alkalinity (ppm as CaCO₃) - 0.714 x Ca (ppm) - 0.585 x Mg (ppm)

For water where the report gives alkalinity and total hardness (TH) both in hardness (ppm as CaCO₃), the residual alkalinity can be estimated by assuming ~80% of hardness is Ca, and ~20% is Mg. The formula is then:

RA = Alkalinity - 0.80 x TH / 3.5 - 0.20 x TH / 7

6.2 CALCULATING pH SHIFT

Now for the most important part of this entire article! The pH shift! Kolbach observed that water with a residual alkalinity of 10 degrees German Hardness caused the pH of a pale malt mash to be 0.3 pH units above the pH realized if distilled water was used.

This observation quantifies the relationship between residual alkalinity and its resultant effect on mash pH. Kolbach did his work based on German degrees of hardness, but if we convert his numbers to appropriate US units, we find that the pH shift of the mash can be calculated from Residual Alkalinity as follows:

pH Shift = 0.00168 x RA

Where residual alkalinity (RA) is in ppm as CaCO₃ (as calculated above)

Assuming a 100% pale/pils malt with distilled water results in a mash pH of 5.8, this gives us the estimated pH of the mash:

pH = 5.8 + 0.00168 x RA

6.3 EXAMPLE PROBLEM

Example 1:

Burton on Trent water, HCO₃^- = 300 ppm, Ca = 295 ppm, Mg = 45 ppm

First we need to calculate alkalinity in terms of ppm as CaCO₃. This is done by assuming the water pH is less than 8.3 (which is likely) and recognizing that 100% of alkalinity is from HCO₃^-; since the pH is less than 8.3 (see carbo chart in alkalinity section) CO₃^-- is zero, and remember H₂CO₃ does not contribute to alkalinity. Convert to the proper units as follows, remembering that 300 ppm = 300 mg/l

300 mg HCO₃^-/liter x (mMol/61 mg HCO₃^-) x (50 mg CaCO₃/mMol)= 246 mg/l as CaCO₃

Plugging into the appropriate formulas gives:

RA = 246 – 295 x 0.714 – 45 x 0.585, RA = +9.05 ppm as CaCO₃.
and

Mash pH = 5.8 + 0.00168 x (+9.05), pH = 5.815

Discussion:

Therefore a 100% pale malt mash with this water is predicted to be 5.815 pH. This is not a significant amount of increase over 5.8, but consider that the optimum range of a mash pH is 5.3 to 5.5 pH, and 5.8 pH is the upper bound for a proper mash. However, a British Pale Ale often consists of up to 10% crystal or other higher kilned specialty malts. All higher kilned malts contribute to the acidity of the mash, and would lower the pH below the 5.815 pH predicted for a 100% pale malt mash. With a grainbill of 20% crystal or roasted malt, the pH will be shifted down as much as 0.5 pH.

7.0 WATER FOR EXTRACT BREWING:

The water used for extract brewing is not nearly as critical. Since your malt has already been mashed, you don't have to worry about the mash pH, and are only concerned with flavor impact ions like Na, Cl, Mg, and SO4. However, you don’t want water that is so hard it tastes bad. Metro Nashville water is relatively soft, and is perfectly fine for extract brewing all beer styles.

You may want to consider adding a small amount of CaCl to aide clarity and hotbreak. The chloride is also good in malty styles. Some gypsum would accentuate the hops in an English Bitter. Do some research, and take a look at the recommended mineral table in the attachment to this article.

It is worth noting however, that extract contains all of the minerals (calcium, carbonate, sulfate, chloride etc.) of the extract Manufacturer’s water supply. The mash uses their water supply, then the wort is run off and evaporated, concentrating all the minerals into the extract (and some of the manufacturers have hard water supplies). Now, when you use Nashville tap water you are adding in even more minerals.

For this reason, if you feel you are getting some harsh carbonate or sulfate bitterness in your extract beer, I recommend switching to 100% distilled water. If that doesn’t clear it up then you know it is not a water issue.

8.0 REMOVAL OF CHLORINE COMPOUNDS

If you don’t already know, chlorine is great for keeping pools clean and sanitizing things, but it is very bad for brewing. Failure to remove chlorine compounds prior to adding steeping grains, mashing in, or adding extract will result in the polyphenols (aka tannins) in the grain/extract reacting with the chlorine compounds and creating funky medicinal off flavors called chlorophenolics, and you don’t want that in your beer.

Metro does seem to have a large amount of chlorine odor and the water report shows the yearly average level of trihalomethanes (THM) that is half the acceptable EPA limit of the maximum allowed. THMs are a cancerous byproduct of chlorine based sanitizer reacting with organic material in the water, and at high levels is indicative of a poor separation of organics from the water that is treated.

Anyway, if you use Metro water, I would recommend you use a carbon filter tube – the inline kind made for ice maker tubing, or the undersink kind. The carbon will filter out all chlorine in Metro water. It will not remove chloramines, so if you are not on Metro water, you might want to call and ask if your supplier adds ammonia to the water to stabilize the chlorine and create chloramines.

If you can’t filter your tap water, do one of the following:

· Use all store bought spring water. It is ozonated, not chlorinated, and all excess ozone is destroyed before bottling.

· Boil your tapwater then let it cool back down before adding steeping grains (works for chlorine only)

· Add ½ a crushed campden tablet (or 1/16 teaspoon of metabisulfite) per
5 gallons cool water, stir in and let stand 5 minutes for chlorine to be driven
off. Not advisable for people who can’t tolerate sulfites (wine has tons of them). Note: This is the only known method to adequately remove chloramines short of using an RO water filter.

Stayed tuned for the exciting conclusions in Part 2 of this series!!

ATTACHMENT 1:

RECOMMENDED MINERAL RANGES FOR BREWING WATER

The recommended ranges for brewing ions/minerals are shown in the following table, and are adapted from various sources. Also shown are the actual levels found in Nashville Metro water (2001 yearly average).

Recommended mineral ranges:

Name	Symbol	Ideal Range*	Metro Water*	Notes
Bicarbonate	HCO₃^--	0-250 ppm	61 ppm	Ideal range varies by style. High levels in pale-based beers cause high mash pH, decreased enzyme activity, and can increase tannin levels extracted from grain husk, creating tea-like harshness. Increases ‘reddening’ of wort.
Sulfate	SO4^--	<150 ppm	44 ppm	At upper levels can increase perception of hop bitterness and makes it seem drier. 150-300 ppm ok for British style bitters
Chloride	Cl^-	<200 ppm	8 ppm	Rounds out flavor of beer, keep low if paired with sodium.
Sodium	Na⁺	<150 ppm	5.3 ppm	At 70-100 ppm it rounds flavor and accentuates sweetness of malt.
Magnesium	Mg⁺⁺	10-15 ppm	5.3 ppm	Yeast trace nutrient, required at 10-15 ppm level, but some is provided by grains. Higher levels cause unpleasant sourness.
Calcium	Ca⁺⁺	50-100 ppm	29.6 ppm	Required for mash enzyme stabilization and a yeast nutrient. Higher levels increase hot break of boil and clarity of finished beer.

* Note that ppm is equivalent to mg/l for all measurements

Additional parameters:

Parameter	Ideal Range	Metro Water	Notes
pH	6.5-8	7.0-7.2	Generally unimportant to brewing chemistry, but required to calculate alkalinity using only HCO₃^—ppm if pH>8.3
Alkalinity (ppm as CaCO₃)	Varies by style	50	Ideal range varies by style. Same impacts as bicarbonate listed above.

ATTACHMENT 2:

WATER CHEMISTRY UNIT CONVERSION TABLE

Use this table to convert a water chemistry value from one unit to the other. The ‘mEq/l’ designates milliequivalents per liter. Simply put this is the molecular weight divided by the absolute value of the charge. For example, Ca++ is 40 / 2 = 20, since it has a charge of 2. Likewise HCO₃^- is 61 / 1 = 61, since it has a charge of only 1.

To Get	From	Do This
Alkalinity as CaCO₃	HCO₃ (ppm)	Divide by 61 and multiply by 50
Ca (mEq/l)	Ca (ppm)	Divide by 20
Ca (ppm)	Ca (mEq/l)	Multiply by 20
Ca (ppm)	Ca Hardness as CaCO₃	Divide by 50 and multiply by 20
Ca Hardness as CaCO₃	Total Hardness as CaCO₃	Divide by 5 and multiply by 4 (estimated)
Ca Hardness as CaCO₃	Ca (ppm)	Divide by 20 and multiply by 50
CaCO₃(mEq/l)	CaCO₃ (ppm)	Divide by 50
HCO₃ (mEq/l)	HCO₃ (ppm)	Divide by 61
HCO₃ (ppm)	Alkalinity as CaCO₃	Divide by 50 and multiply by 61
Mg (mEq/l)	Mg (ppm)	Divide by 12.1
Mg (ppm)	Mg (mEq/l)	Multiply by 12.1
Mg (ppm)	Mg Hardness as CaCO₃	Divide by 50 and multiply by 12.1
Mg Hardness as CaCO₃	Total Hardness as CaCO₃	Divide by 5 (estimated)
Mg Hardness as CaCO₃	Mg (ppm)	Divide by 12.1 and multiply by 50
Total Hardness as CaCO₃	Ca as CaCO₃ and Mg as CaCO₃	Add them.

Note: ppm is equivalent to mg/l for all measurements