Water Hardness Is Not That Hard To Understand
(Water Chemistry Without Pain)

by Jeff Renner

Water Hardness is the cause of a lot of confusion to craftbrewers who haven't done a lot of chemistry. Even those who did high school chemistry still struggle with the concept. I did as well when I first started brewing and trying to figure this all out, hard and soft water, permanent and temporary hardness. I've been turning over in my mind a narrative explanation of these matters.

I have wanted to write it up as "Water Chemistry without Pain" or some such title for a while. Chemical symbols and formulas can scare some folks off right at the start. Well here finally is an article that I hope it is a useful explanation.

Hard and Soft Water

First, hard and soft water - what are they? Historically, before anyone knew anything about chemistry, people discovered that some water made lovely, soft, creamy lather with soap. Rain water is an example of this. Before the advent of mains water, many people had rain water tanks or barrels to collect water for washing laundry and hair, separate from water from other sources, such as bores and wells. Other water, especially bore water, made sticky, un-slippery, un-lathery "soap curd", left a bathtub ring of this material at the high water mark, and was generally poor for cleaning, unless a lot of soap were used. Even today in rural Australia and many rural communities in other countries, rainwater is still collected separately from bore water for this reason. Naturally, the term "soft water" came to be applied to the former, and "hard water" to the latter. Hard water was hard to clean with, hard to get soap to lather, and the staining was hard to get rid of. Soft water was on the other hand easy to use, soft on the skin and on clothing when you washed it.

Sometime in the 19th century, science found the explanation for this phenomenon. Hard water has dissolved calcium and/or magnesium ions from dissolved underground minerals, in other words high in dissolved minerals. Why do these make soap work poorly? To understand this, we need to know what soap is. It is a sodium salt of a fatty acid. It is a fat molecule with a sodium ion at the end. In fact soap is actually made from animal and/or vegetable fat, where all you do is change it to a salt by the addition of sodium (from caustic soda and ordinary table salt).

This end of the soap molecule is attracted to water (sodium <-> fatty acid) this end is soluble in fats and oils.

Soap

Soap allows the mixing of two things that don't like to mix, oil (or fats) and water. The fatty end mixes with fatty dirt, and the ionic sodium end is attracted to water. Voila, emulsified dirt. Other common examples of emulsions where fats are in solutions are milk and mayonnaise.

However, the fatty part of soap would rather have a calcium or magnesium ion on its end, so in the presence of these ions in hard water, they quickly replace the sodium ion making a calcium and/or magnesium soap. Now both calcium and magnesium is not strongly attracted to water like sodium, and soap made from calcium and magnesium are not soluble in water, so it precipitates out as soap curd. If you use enough soap, it uses up all of the calcium and magnesium ions in the water, and then you have some sodium soap left for cleaning. At about the same time as this was explained, it was also discovered why certain kinds of waters made better beers of certain styles.

Permanent and Temporary Hardness

The is the formal definition of hard water is the presence of calcium and magnesium ions. It has nothing to do with the presence of bicarbonate, carbonate, chloride, sulfate, etc. This confused me no end when I first began brewing and read of temporarily and permanently hard water, and carbonate and sulfate waters.

I then learned that temporarily hard water has calcium (and magnesium, usually in smaller amounts) in the presence of bicarbonate ions, which, when boiled, lose carbon dioxide and become carbonate ions. They are insoluble in the presence of calcium or magnesium ions, and precipitate out as calcium (or magnesium) carbonate, which are virtually insoluble.

Permanently hard water has calcium and magnesium in solution, with or without bicarbonate. The ions in this case are typically sulfate and chloride. Boiling has no effect on these ions, so the calcium and magnesium remain in solution to produce soap curd or, happily, to react with phytase to produce phytic acid in our mash and bring the pH to a proper level. This is why we need a minimum level of calcium in our mash water.

Making Hard Water Softer

So how do make hard water softer. Well you add slaked lime. How does adding slaked lime, Ca(OH)2, reduce bicarbonate? Well, again I'd like to offer the full explanation (my kids say I can never give the short answer). How does bicarbonate get into water in the first place (increasing the temporary hardness of the water)? Rain water falling through the atmosphere dissolves carbon dioxide, producing carbonic acid, naturally acid rain. (Rain falling through atmosphere polluted with smokestack sulfates from coal burning produce sulfuric acid, unnaturally acid rain).

This slightly acid rain percolates through the ground, where it encounters limestone, or calcium carbonate, which is the remnant of the shells of prehistoric sea life. Calcium carbonate is virtually insoluble in water, which is a good thing for the sea life. A clam would be in real trouble if its shell dissolved! But it is soluble in acid, even a very weak acid. So, this acidic rainwater which is now ground water, dissolves the limestone. The addition of carbonic acid to the carbonate produces bicarbonate, a terrible name,
HCO3-. The Germans call it hydrogen-carbonate, which is a much better name. Bicarbonate in the presence of calcium is very soluble.

It is also fairly alkaline, which causes troubles when mashing pale grains, because the enzymes work better if the water is not alkaline. However, this alkalinity is useful when mashing dark grains, which are acidic. Brewers discovered this empirically centuries ago when they found that certain areas such as Munich, London and Dublin, with their carbonaceous (really, bicarbonaceous) water made good dark beer but lousy pale beers. Burton-on-Trent, with its sulfate water, or, more to the point, non-carbonaceous water, made good pale beer and Pilsen, whose water is soft, can make good pale beer, not because it is soft per se, but because it has no bicarbonate.

Slaked lime, which is called that because it is lime, or calcium oxide (CaO), mixed with water, just as we slake our thirst with water. Lime is also called burnt lime, because the heating of limestone (CaCO3), which drives off carbon dioxide (CO2) and leaves calcium oxide (CaO), produces it. When this burnt lime is added to water (H2O), it becomes slaked lime, (Ca(OH)2), calcium hydroxide. This is the calcium equivalent of caustic soda, and is strongly alkaline. If you remember, calcium bicarbonate is the result of acid dissolving calcium carbonate, which is soluble only in an acid. If we add enough slaked lime, a strong alkali, to the calcium bicarbonate solution to raise the pH beyond pH 10.8, it can no longer hold the calcium bicarbonate in solution. The bicarbonate gives up an H+ ion (which combines with an OH- ion from the calcium hydroxide and becomes a water molecule) becomes carbonate, which combines with the calcium present, and precipitates out as calcium carbonate. It seems counterintuitive that adding calcium would reduce it, but it does, by greatly increasing the pH, or alkalinity, of the water.

The addition of slaked lime, Ca(OH)2, reduces the bicarbonate, but also the calcium, in a 2:1 ratio. This means that by itself, it presents problems for preparing brewing water, since we want calcium in the water, too. As a matter of fact, adding lime is a common, inexpensive way of softening (reducing calcium and magnesium) municipal water supplies. Such municipal water has a high residual pH, typically over pH 9, but it is not strongly buffered, so it typically presents no problem, as long as there is sufficient calcium remaining.

Water Softeners

Home water softeners accomplish the same end result, reduction of calcium and magnesium ions, by ion exchange, in other words swapping them. Sodium ions are exchanged in the water for the calcium and magnesium ions. Soap now works fine, but the bicarbonate remains. No problem for soap, but this water is doubly unsuitable for brewing, since it not only still has the pH raising bicarbonate, but none of the pH lowering calcium ions. You will still need to remove the bicarbonates and add some calcium ions to have a trouble free mash.

I hope that this has made water hardness a little easier to understand.

Jeff

Editors Note:
Jeff Renner has published a number of articles in leading brewing magazines such as Zymurgy and Brewing Techniques. He has kindly given us permission to publish this articles but retains all rights for future publication.