CHAPTER 3
Water
Water constitutes 85 to 90 percent of the volume of any beer, and therefore the mineral content of the brewing water has a marked effect on the flavor and appearance of the finished beer — and on the brewing process. Certain beer styles are suited to waters of very specific mineral composition, and an otherwise well-brewed example will always be diminished by the use of totally inappropriate brewing water.
By looking closely at the geology of any given area, it is often possible to find wells or springs that will perfectly suit a given brew. Local, regional, and state water departments and services can be very helpful in locating such sources. For most brewers, however, mineral treatment of the local supply is the more reasonable alternative.
Only brackish, polluted water and sea water are entirely unsuitable for brewing. Most potable fresh waters, whether too “hard,” too carbonate, too “soft,” or iron contaminated, may be boiled, aerated, sedimented, filtered, or treated with an appropriate mineral salt or acid to be made suitable for brewing almost any type of beer. Practically speaking, however, brewing water should be clear, bright, unpolluted, and have agreeable taste and reasonably uniform composition from day to day. It should not be corrosive, have a detectable odor, or throw an appreciable amount of sediment upon resting or boiling.
All naturally occurring waters are dilute solutions of minerals in which small quantities of gases and organic matter may be dissolved. Rainwater should be the purest natural source of water, but because it assimilates atmospheric gases and organic mineral particles wherever the air is the least bit polluted, most rainwater is absolutely unsuitable for use in brewing. Precipitation in areas far removed from large fossil-fuel burning plants more often than not is still polluted by highly corrosive sulfuric acid (H2SO4). Free hydrogen carbonates (HCO3, usually referred to as bicarbonates) are also common in rainwater. They rob the calcium from the mash, wort, and ferment by forming bicarbonate salts that are precipitated from solution during boiling.
Surface waters, besides having the dissolved materials and gases found in rainwater, usually contain large amounts of organic matter, vegetable coloring, soil, silica, clay, and microflora. Especially in marshy terrains and industrial areas, surface water is likely to be heavily contaminated with organic acids and nitrates and is completely unsuitable for use in brewing.
Surface waters suitable for brewing are generally limited to clear-running, spring-fed brooks and streams that flow over gravel, sand, or rocky beds, and deep reservoirs with carefully protected watersheds. Water from old soft-bottomed streams, which flow sluggishly or carry topsoil and vegetation in suspension, and rivers, ponds, and lakes needs to be filtered, at the very least, before use, and is seldom a good choice for brewing. If the water tastes like it comes from a pond, so will the beer.
Municipal water supplies are usually gathered from several deep wells and reservoirs, and sometimes from rivers. They are invariably filtered and treated (most commonly with up to 0.5 parts per million chlorine) to inhibit microbial contamination. Such tap water is often perfectly suitable for use in brewing, after being filtered, rested, or boiled and aerated to drive off free chlorine and induce sedimentation of carbonates, silicates, and incrustants. Chlorine and organics may be removed by activated carbon filtration. Water departments often vary the inflows from several different sources, and water composition varies accordingly.
The quantity and composition of underground water at any given location and depth are contingent upon subsurface geological formation. Although the elementary minerals and metals dispersed in the earth’s crust are relatively few (rock is largely silica with aluminum, iron, calcium, potassium, magnesium, manganese, zinc, and copper), the soluble components (mineral salts) they form yield water of varying composition from place to place. The mosaic geological structure of the earth’s crust causes water to rise to different levels at different locations, and its level may be fairly constant or may fluctuate with changing patterns of precipitation, freezing, and thawing.
The value of a spring or well as a brewing source should first be judged by the seasonal consistency of its flow. There is less fluctuation in the composition of any ground water that has a reasonably constant flow year-round. Shallow wells and seasonal springs do not usually yield acceptable brewing water. Their composition varies widely, and they often carry soil and other surface contamination in suspension. They commonly yield unacceptable levels of bacterial contamination.
Deep wells originating in large subterranean aquifers, and mountain springs percolating up through fissures in bedrock formations without leaching through soil or disintegrated rock are usually of very stable composition and free from surface contamination. Subterranean springs (with the notable exception of mineral springs) emerging from inorganic rock complexes commonly yield water with less dissolved materials than does deep-well water. Deep wells usually tap water sources that have traveled farther than the water of springs, and having contacted more mineral-bearing substrata consequently have more minerals in solution than does spring water. Which minerals occur in ground water is dependent on the nature of the strata the water has contacted. Igneous and metamorphic rocks (granite, basalt, marble, gneiss, and quartz) are both very dense and compositionally very stable; they yield the fewest mineral ions to hydrolysis of any rock formations. Because they are hard, they do not filter the water passing over them as do most sedimentary formations. They are not likely to bear much water unless they are extensively fissured or enclose voids left by the dissolution of softer strata. Granite is the most common of the igneous rocks, composing the greatest part of the continental plates. It commonly yields very “soft” water of less than 100 ppm hardness (as CaCO3), and water of less than 50 ppm hardness is not at all uncommon.
These rocks are eroded by freezing water, scaling, hydrolysis, or friction, and the particles are carried away by wind and water and ultimately deposited in topographical depressions. This alluvia forms unconsolidated sediments, often far from the parent rock.
Sand, which is largely silica dioxide split off from quartz, passes water freely while filtering out most suspended solids. Where it lies above impervious bedrock, it yields excellent water. Clay, on the other hand, is impervious and yields very little water, but water pools above it wherever it makes an unbroken sediment. Clay underlies many excellent aquifers. Clay is mostly hydrous silicates of aluminum, often colored red by iron oxide or dark by carbon-based impurities. Gravel (pebbles of ⅛ to 2 ½ inches in diameter) is usually found with sand or clay, either on river terraces or as the residue of glacial retreat. It often serves to make clay subsoil pervious to penetration by some surface runoff.
Consolidated sediments are deposits of sand, clay, or gravel that have hardened under pressure or have been cemented by lime. They are typically deposits left on the floors of prehistoric seas and lakes. Sandstone is very common; it is porous and generally rather coarse. It filters water very effectively but also gives up its mineral ions to hydrolysis very readily. Sandstone formations yield predominantly “permanent” hardness, but the mineral composition and “total” hardness vary widely from place to place. (For an explanation of water hardness, see pages 61 through 65.) Water drawn from older sandstone may be soft but is usually moderately hard, averaging 50 to 300 ppm hardness. New red sandstone usually bears very hard water (150 to 400 ppm). Although the hardness is largely sulfate, water drawn from new red sandstone is often objectionably inundated by iron.
Formed from loose sediment, limestone (calcium carbonate) is not an extremely dense rock, but it is impervious and yields water only if it has been extensively fissured by seismic activity or eons of hydrolysis. The passage of water along joints and bedding plates has carved out huge subterranean caverns and labyrinths in some limestone formations, where abundant water may be tapped. Carbonic acid from the atmosphere readily dissolves calcium carbonate, forming very soluble calcium hydrogen-carbonate (bicarbonate) salts. Consequently aquifers and springs from limestone tend to yield water high in “temporary” hardness, with other minerals in solution. Some limestones are largely made up of magnesia and may yield considerable sulfate as well. Hardness of 150 to 350 ppm is common.
Chalk formations are very soft, fine-grained limestone from mudlike sea-bottom sediment. They yield from 150 to 375 ppm hardness, largely as calcium carbonate. Soapstone (talc) is insoluble hydrous magnesium silicate, often found with magnetite (iron oxide) and chlorite. Formed from consolidated clay, shale is impervious but may pass water along its bedding plates. Like other sedimentary rock, it commonly contains many minerals.
Marl is consolidated clay with sand, and most often calcium carbonate, potash, or phosphorus. It is commonly found stratified with sandstone or limestone. Marl generally yields little water, and of an undesirable mineral composition. Conglomerates are impervious, hard rocks from lime-cemented gravel and sand; they vary widely in composition but usually have little effect on the water coming in contact with them.
Using the Water Analysis
Any potential source of brewing water should be analyzed for organic and inorganic composition and biological purity. Because the composition of any water supply is likely to vary seasonally, and even within each season, it is advisable to make or obtain an analysis of it before brewing, or when changes in the brew may be due to changes in the water. Where periodic testing is not possible or practical, a simple pH test demonstrates changes in the mineral content and can indicate what the changes in the mineral distribution may be.
The standard water analysis identifies the amounts of the mineral ions present in the water and indicates the presence of organic pollution as well. A bacterial analysis may be included or made separately. Analyses for community water supplies are available upon request from local water departments; other sources are analyzed by private labs for a fee. Do-it-yourself kits are also available for identifying the pH, hardness, and alkalinity of water, and the presence of various mineral ions (calcium, magnesium, iron, chlorine, sulfate, nitrate, nitrite); they are easy to use, handy, and inexpensive over the long run if analyses would otherwise have to be made by a private lab.
| Water: Typical Analysis U.S. Public Health Service (U.S.P.H.S.) Units | ||||
|---|---|---|---|---|
|
Turbidity |
Color |
pH |
Sediment |
Odor |
|
Specific Conductivity (micromhos/cm) |
||||
|
Hardness – Total (CaCO3, in ppm [mg/L]) |
||||
|
Alkalinity – Total (CaCO3, in ppm) |
||||
| Major Constituents 1–1000 ppm |
|---|
|
Calcium (Ca) |
|
Magnesium (Mg) |
|
Sodium (Na) |
|
Silica (SiO2) |
|
Sulfate (SO4) |
|
Chloride (Cl) |
| Secondary Constituents .01–10 ppm |
|---|
|
Potassium (K) |
|
Nitrogen (Ammonia, NH4) |
|
Nitrogen (Nitrite, NO2) |
|
Iron (Fe) |
|
Nitrogen (Nitrate, NO3) |
| Minor Constituents .001–.1 ppm |
|---|
|
Manganese (Mn) |
|
Copper (Cu) |
| Trace Constituents less than .001 ppm |
|---|
|
Zinc (Zn) |
|
Coliform bacteria (colonies per 100 mL) |
If testing by a lab is necessary, certain procedures for obtaining the specimen are advisable. Rinse a clean quart jar several times with the water to be tested; fill the jar and immerse it nearly to its neck in a kettle of water. Heat and boil for twenty minutes. Decant the water from the jar and invert the jar on a clean paper towel to drain. When cool, fill it with the water to be tested. If it is tap water, allow the cold water to run for a minute before taking the sample to flush clear any mineral deposits jarred loose by the initial release of water from the tap. Cap the jar tightly and rush it to the lab where you have made arrangements to take it. The longer the sample sits, the less accurate the analysis will be. Where pollution is suspected, the water must be tested within twelve hours of collection, and in all other cases, within seventy-two hours.
When the brewing water is to be treated with mineral salts or boiled to precipitate carbonates, an analysis after treatment will pinpoint the resultant mineral distribution, but it is hardly necessary. The carefully weighed addition of salts and a simple pH test after treatment usually adequately indicates the subsequent mineral distribution.
Brewing with an untreated natural water supply is possible in almost all cases by manipulating the brewing procedure. A water analysis should be made before brewing to determine the formulation and procedure that best suits the particular water composition.
Turbidity, Sediment, Color, and Odor
Turbidity and sediment may be caused by suspended clay and other inorganic soil, organic topsoil or waste, colloidal ferrous and aluminum oxides, or manganese and silicon dioxide.
Color is usually due to colloidal vegetable pigments, although a yellow-to-brown hue may be from suspended clay or silt. This sediments out upon resting; particles in colloidal suspension may be eliminated only by filtering. Odor may be from dissolved gases or organic decay.
Overall character may be improved by activated-carbon filtration. Where this is not satisfactory, a clearer water should be found to brew with.
U.S. Public Health Service Drinking Water Standards suggest as limits that should not be exceeded: turbidity — five units; color — fifteen units; odor — threshold number three. For brewing, it is recommended that these all be less than one.
pH
The pH indicates acid to alkalinity ratios and the mineral composition of water as well, both of utmost importance to the brewer. Appropriate acidity is a prerequisite of a successful brewing cycle. Enzyme activity, kettle break, and yeast performance rely upon conducive acidity in the mash, wort, and beer. The acidity of the brewing water source is therefore of concern to the brewer.
pH is the measure of the acidity or alkalinity of a solution. Acid solutions taste sour; alkaline solutions taste bitter and flat. In other terms, acidity is expressed as a greater number of positively charged particles in solution than negatively charged ones; conversely, alkalinity marks an excess of dissociated negative particles.
Ions
All elements are reducible to single atoms. Atoms are made up of an equal number of positive and negative charges, respectively termed protons and electrons. The atoms of each element are distinguished from those of every other element by the number of protons in their nuclei and the number and arrangement of electrons they have in orbit. Atoms are chemically inactive because the charges of their protons and electrons neutralize each other. They may, however, be unstable. Only atoms having two electrons in their first orbit and eight electrons in their outermost orbit are stable; only these elements occur in their true atomic form.
Atoms
An unstable atom either gives electrons to or receives electrons from another unstable atom, each thereby forming an ion of the element that has a stable electron configuration.
Ions
Unstable Elements Sodium and Chloride
Electron Transfer
Ionic Compound, Sodium Chloride
Because this results in an imbalance in the number of positively charged protons and negatively charged electrons, all ions have an electromagnetic charge. Ions that have a positive attraction are cations; negatively charged ions are anions.
Cations and anions combine to form ionic compounds; the strength of their bond and the charge of the compound itself are dictated by the relative electromagnetic attraction of the ions involved. The greater the difference in their charges, the stronger the ionic bond, and the more acid or alkaline their compound.
Ionic compounds are formed by the exchange of electrons. Other compounds are formed by sharing electrons; these are termed covalent. Most organic compounds (compounds containing carbon) are covalent.
The water molecule, H2O, is covalent. Only a single electron orbits the nucleus of a hydrogen atom; the outer orbit of an oxygen atom contains only six electrons. Both atoms are unstable, hydrogen needing to give up an electron and oxygen needing to receive two. In the covalent compound of the water molecule, the hydrogen atoms achieve stable single orbits by sharing their electrons with the oxygen atom, which then has a stable outer orbit of eight. Since no electrons are exchanged, the bond is nonionic, and the molecule is electromagnetically neutral. The negative valence (-2) of the oxygen atom (0--, or lacking two electrons) is precisely neutralized by the combined positive valence (+2) of the two hydrogen atoms (H2++).
Although neutral, the water molecule retains a strongly polar character, because the eight protons of its oxygen molecule have slightly more attraction for the negatively charged electrons of the shared orbits than do the single protons of the two hydrogen atoms. The oxygen side of the water molecule is thus slightly negatively charged and the hydrogen side equally positively charged; it is even polar enough to disrupt the ionic bonds of many inorganic compounds, causing them to dissolve into their component ions. When all of the water present is absorbed in the reaction with an ionic compound, that compound is said to be hydrated by water of crystallization. For example, gypsum (CaSO4⋅2H2O) is a hydrated salt. Like other compounds, it can be dissolved into its component ions by the introduction of more water. In fact, water is the most universal solvent known, and it is able to react chemically with most inorganic acids and bases, and to dissolve many salts.
Although pure water is characterized as being covalent, a certain number of its molecules react in pairs to form hydroxide ions (OH-) and hydronium ions (H3O+, most often expressed as simply H+). This is an equilibrium reaction, and the combined number of hydroxide and hydronium ions always remains constant; in pure water at 77 degrees F (25 degrees C), the concentration of electropositive hydronium ions and electronegative hydroxide ions are each .000,000,1 moles per liter. When mineral compounds dissolve in water, their positive mineral ions bond to the oxygen side of the water molecule, freeing hydronium (H+) ions, and their negative ions react with the electropositive hydrogen side, releasing hydroxide (OH-) ions.
Water Molecule
When an excess of either hydroxide or hydronium ions is released, the ionic equilibrium of pure water is disturbed. Because the combined value of hydroxide and hydronium ions always remains constant, an increase in either always results in a proportional decrease in the other.
Measuring pH
The pH scale represents the relative molar concentration of hydrogen and hydroxide ions present in any solution by measuring the activity of the free hydronium ions in solution. If the hydronium ion concentration exceeds .000,000,1 moles per liter, the solution is acidic; if it is less than .000,000,1 moles per liter, the solution is alkaline and will neutralize acids and liberate CO2.
The reaction that the hydrogen ion concentration represents causes a color change in litmus paper and other pH indicators. The degree of color change is gauged against a standard scale to identify the pH of any solution. For a more detailed analysis, electromagnetic equipment is employed.
The pH, however, is not expressed as the molar concentration of hydrogen and hydroxide ions. The pH scale uses the exponent of 10 in the logarithm (see table 5) to identify the acidity of a solution. Because it is a logarithmic scale, a solution at pH 4 is 10 times more acidic than a solution at pH 5, 100 times more acidic than a solution at pH 6, and 1,000 times more acidic than a solution at pH 7.
| Part A. | ||||
|---|---|---|---|---|
|
Moles Per Liter, H+ |
pH |
Moles Per Liter, OH- |
||
|
.1 |
1⋅10-1 |
1 |
1⋅10-13 |
.000 000 000 000 1 |
|
.000 1 |
1⋅10-4 |
4 |
1⋅10-10 |
.000 000 000 1 |
|
.000 01 |
1⋅10-5 |
5 |
1⋅10-9 |
.000 000 001 |
|
.000 000 1 |
1⋅10-7 |
7 |
1⋅10-7 |
.000 000 1 |
|
.000 000 001 |
1⋅10-9 |
9 |
1.10-5 |
.000 01 |
| Part B. Hydrogen Ion Concentration in the pH Scale | ||||
|---|---|---|---|---|
|
pH |
4 |
5 |
6 |
7 |
|
H+, moles/L |
.000 1 |
.000 01 |
.000 001 |
.000 000 1 |
pH Adjustment
All changes in the pH of the brewing liquor, the mash, the wort, or the beer, whether induced or consequential, are due to the formation, addition, or precipitation of mineral ions or organic acids. The pH of the brewing water may be adjusted by precipitating out alkaline carbonate salts or by adding organic acids or mineral salts. Salts are added when additional mineral character is also desired. As an alternative to water treatment, the mash may be made more acidic by the metabolism of certain bacteria, which causes the formation of organic acids.
The pH of the mash affects the level of enzyme activity within it, and the acidity of the wort and beer. The pH significantly affects hop extraction and protein precipitation in the kettle, and yeast performance and clarification in the ferment. Because ideal pH levels often cannot be attained, or because the pH at one stage conflicts with the pH optimum of a more critical reaction, concessions are sometimes made; shortcomings can usually be overcome by time and temperature manipulation. For example, in the mash the enzymatic reduction of proteins to soluble nitrogen is most efficient at pH 5 or below, a level of acidity that conflicts with starch reduction optimums during mashing. Therefore, it is necessary to make a longer rest at temperatures conducive to protein degradation, especially when the mash begins at above pH 5.5.
At above pH 6, any mash suffers from sluggish enzyme activity, and in the lauter-tun, troublesome tannins and silicates are leached into the extract. With “soft” water and pale malt, the few acid ions cannot overcome (“invert”) buffers leached from the malt, which results in a strong resistance to further acidulation of the mash.
pH 5.2 to 5.5 should be the target acidity of the saccharification rest for all mashes. This is the range at which enzyme activity, filtering, color, and clarity are best. A mash at pH 5.2 to 5.5 can be expected to yield a sweet wort of 5.5 or slightly above, which best serves hop extraction and flocculation of protein in the kettle.
Total Dissolved Solids/Specific Conductivity
The total dissolved solids are the mineral ions in solution. Total dissolved solids are measured by passing water through a filter fine enough to screen out all sediments and colloids, then weighing the residue left after the filtered water has been evaporated.
The specific conductivity of water measures the mineral ions in solution by gauging the solution’s ability to conduct an electric current at 77 degrees F (25 degrees C). Most inorganic acids, bases, and salts are good conductors of an electric current because ions are by definition electrovalent, or charged, particles, whereas organic compounds conduct current very poorly, if at all. Specific conductance, therefore, is a measure of the total dissolved solids in any water.
The specific conductance in micromhos per centimeter (mho is the opposite of the basic unit of electrical resistance, the ohm) is roughly equatable to the total dissolved solids in solution in parts per million (ppm is the same measurement as milligrams per liter [mg/L]) by multiplying the specific conductance by a factor that ranges from .55 (for water of pH less than 7.2 or more than 8) up to .90 (for saline water). For most natural waters, multiplying the specific conductance by a factor between .55 and .70 gives a reasonable assessment of the total dissolved solids.
Hardness/Alkalinity
The hardness of water is gauged by measuring the dissolved cations of the alkaline-earth elements, most significantly calcium (Ca++) and magnesium (Mg++). These common minerals inhibit the sudsing of sodium-based soap in “hard” water and are precipitated as an insoluble, furry residue.
Alkaline-earth ions are weakly electropositive and give water slight acidity. The polar bonds they form with water molecules are weak and readily broken by strongly basic anions, causing them to precipitate as insoluble salts. Calcium exhibits a fragile solubility. Magnesium is more electropositive than calcium but is not correspondingly less soluble; it is actually more stable in solution. The cations of the larger, less electropositive alkali metals sodium (Na+) and potassium (K+) are even more weakly acidic. They are much more stable in water, give an essentially neutral pH reaction, and do not contribute to water hardness.
Calcium is the most widely occurring “metal” found in water, followed by magnesium, sodium, and then iron. Potassium and manganese are much less common; when considering their effects, potassium is often grouped with sodium, and manganese with iron.
The metal cations all occur out of solution bonded to acid anions in crystalline-structured ionic compounds called mineral salts. The bicarbonate, sulfate, chloride, nitrate, borate, and phosphate ions are all classified as acids because they are derived from carbonic, sulfuric, hydrochloric, nitric, boric, and phosphoric acids; their effects, however, are decidedly alkaline.
The several salts that may be formed by the acid anions with any given metal vary in solubility and acid or alkaline reaction according to the electronegative valence of the particular anion involved. A weak metal and a weakly alkaline anion in solution have very little attraction for each other and stay in solution. The same metal with a moderately alkaline anion may be precipitated out of solution under certain conditions. With a strong base, the metal may even be insoluble. Thus, calcium chloride is freely soluble, calcium sulfate is of limited solubility, and calcium carbonate is nearly insoluble. Sodium and potassium, because they are only slightly electropositive at best, are freely soluble not only with both chloride and sulfate ions but also with the carbonate ion.
Similarly, the pH of a solution is determined by the relative electrovalence of the several cations and anions dissolved together. Since soluble metals are all only weakly acidic and acid anions range from very weakly alkaline to extremely alkaline, their solution in water may be slightly acidic (calcium and the sulfate ion), neutral (sodium and chloride ions), or very alkaline (calcium and carbonate ions). Mild alkalinity is usually indicative of solutions containing more than one anion (calcium with the sulfate and carbonate ions, for instance).
The anions found in water are almost exclusively sulfates (SO4--), chlorides (Cl-), and bicarbonates (HCO3-). The sulfate ion is weakly basic, and the chloride ion only slightly more so, whereas the unstable bicarbonate ion is a very strong buffer, because in its formation from the carbonate ion (CO3--), it pulls a hydronium ion off a water molecule, freeing hydroxide ions into solution.
Far less common than the sulfates, chlorides, and bicarbonates are the weakly basic nitrate, borate, phosphate, and silicate ions; few of their salts are even soluble. In fact, only six salts commonly dissolve in ground water: calcium bicarbonate, magnesium bicarbonate, calcium sulfate, magnesium sulfate, sodium sulfate, and sodium chloride. Only three other salts are occasionally present in significant amounts: calcium chloride, magnesium chloride, and sodium bicarbonate. Potassium compounds are rarely present in any quantity.
Calcium and magnesium significantly affect the brewing process. When calcium and magnesium occur primarily with the bicarbonate ion, calcium precipitates out of solution during boiling, potentially robbing the yeast of a necessary element of its composition and causing high mash pHs that may react sluggishly to acidification. On the other hand, calcium in solution with the sulfate ion provides a very stable vehicle for the transmission of calcium into the ferment, and aids rather than retards mash acidulation.
The hardness of water expresses the calcium and magnesium in solution. The hardness of any water is determined by titration, as with EDTA, after addition of a dye, to an end point. It is expressed as “hardness as CaCO3,” although it represents all of the calcium and magnesium ions in solution, arbitrarily combined with carbonate ions.
Bicarbonates, and to a lesser extent, carbonates, constitute most of the alkalinity of natural waters. Combined with calcium, they are expressed as the temporary or carbonate hardness of water, or that part of the hardness that will precipitate out of solution by boiling or with the addition of lime.
After boiling, calcium and magnesium ions remain in solution with noncarbonate ions. These are expressed as the permanent or noncarbonate hardness. In fact, some carbonate ions remain, bonded to water molecules; calcium carbonate is soluble to 20 ppm, and with magnesium in solution, magnesium carbonate may remain at up to 300 ppm.
The alkalinity of water measures the buffering capacity of dissolved anions, especially the bicarbonate (HCO3; the carbonate ion CO3 is only a significant factor in waters of pH 8.3 and above). By titrating the alkalinity, the temporary hardness can be assessed more readily than by boiling the water and can be used to precisely indicate treatment. An indicator dye, such as Bromcresol Green-Methyl Red, whose color in neutral and alkaline solutions is known, is added to a measured volume of water. The number of drops of a strong mineral-acid solution, such as .035N sulfuric acid, it takes to neutralize the water (overcome its alkalinity) is indicated by the color change of the dye, from blue-green to methyl-orange end point. When multiplied by a factor based upon the sample size, the number of drops it takes to reach end point expresses the bicarbonate alkalinity as parts per million of calcium carbonate. For more alkaline waters with a pH above 8.3, phenolphthalein is used as an indicator dye, and the carbonate alkalinity is titrated before the bicarbonate is measured.
Hardness actually measures calcium and magnesium in solution, and alkalinity measures all the alkaline ions. By expressing hardness and alkalinity in the same terms, as CaCO3 (calcium carbonate), the two values are readily compared; “as CaCO3” is the accepted standard because its cation is the primary mineral of hardness and its anion is the principal cause of alkalinity in most waters. This convention also simplifies water treatment.
Hardness and alkalinity nicely define the permanent and temporary hardness of water. When the alkalinity as CaCO3 exceeds the hardness, then the hardness is largely temporary. When the hardness value exceeds the alkalinity, the difference is indicative of permanent sulfate hardness. Especially where hardness greatly exceeds alkalinity, the water is eminently suitable for brewing and responds well to acidulation during mashing.
Where alkalinity as CaCO3 is unknown, the hardness before and after boiling must be measured to define permanent and temporary hardness. This method is at least as satisfactory an indicator as is the alkalinity reading of a water analysis, but a great deal more difficult to make.
Temporary hardness is always strongly alkaline; permanent hardness is usually only slightly acidic, so only soft waters or waters where hardness well exceeds alkalinity yield a proper mash pH with pale malt.
Most water supplies are slightly alkaline, due to the buffering of any calcium and magnesium in solution by the strongly basic reaction of even a small amount of bicarbonate. At over 50 ppm alkalinity as CaCO3, water reacts sluggishly to acidulation in the mash and kettle. This water becomes weakly acidic upon precipitation of its carbonate salts. The carbonates can be sedimented out by boiling or the addition of slaked lime, or overcome by the addition or formation of organic acids, and to some extent by adding calcium or magnesium as sulfate or chloride salts to the mash or the mash liquor.
The larger of the two readings, hardness or alkalinity, can also indicate how much of the total dissolved solids are sodium, potassium, and chloride, which contribute to neither the hardness nor the alkalinity, by subtracting it from total dissolved solids.
Where calcium and magnesium measurements are not given in an analysis, dividing the hardness reading by 1.25 and by 5.0 roughly indicates the calcium and magnesium in solution (assuming a four-to-one Ca/Mg ratio).
Molarity, Equivalence, and Normality
The chemist’s tools molarity, normality, and equivalence, as described by Dr. George Fix in Principles of Brewing Science (Brewers Publications, 1989), are of significance to brewers, especially for understanding liquor acidification.
Molarity is a method for quantitative analysis of a substance in solution. Molarity employs the mole (mol, gram-molecular weight, gmw, gram mole, combining weight) as its unit of measure. Molecular weight is the sum of the atomic weights of the elements that compose any given substance; a mole is the molecular weight in grams. A molar solution (M) of any given substance is equal to one mole in a liter of solution.
The atomic weights of the elements sulfur and oxygen are 32.066 and 15.9994. The sulphate molecule, SO4, is composed of one atom of sulfur and four atoms of oxygen. The molecular weight of the sulphate ion is 96.0636, because 32.066 + (4 x 15.9994) = 96.0636. A mole of the sulphate ion, then, is 96.0636 grams, and a molar solution of sulphate is equal to 96.0636 grams in one liter of solution.
A molar solution of sulfuric acid (H2SO4) is 98.08 grams, of phosphoric acid (H3PO4) 97.995 grams, citric acid (C6H8O7) 192.13 grams, and lactic acid (C3H6O3) 90.08 grams in one liter of solution.
Molarity is generally given as a decimal percentage; .01M sulfuric acid means that the solution contains .01 moles of lactic acid per liter, or .9808 grams per liter.
The concentration in parts per million (milligrams/liter) of a molar solution can be derived by the formula ppm = molarity x mole x 1000. So for sulfuric acid, .01M: .01 x 98.08 x 1000 = 9.808 ppm.
Equivalence measures the number of moles of hydrogen or hydroxyl ion that a substance can liberate. The moles and equivalency of some common ions are:
|
Mole weight |
Equivalent Weight |
|
|
H+ |
1.00794 |
1.00794 |
|
Ca++ |
40.078 |
20.039 |
|
Mg++ |
24.3050 |
12.1525 |
|
Na+ |
22.989768 |
22.989768 |
|
K+ |
39.0983 |
39.0983 |
|
SO4- |
96.0636 |
96.0636 |
|
CO3-- |
60.0092 |
30.0046 |
|
HCO3- |
61.01714 |
61.01714 |
|
Cl- |
35.4527 |
35.4527 |
The strength of an acid is measured by its ability to release H+ ions, lowering pH. Sulfuric acid (H2SO4) disassociates when it is added to water, releasing two hydrogen ions into solution: 2H++ + SO4--. It has an equivalency, then, of 2. In a solution containing the carbonate ion (CO3--), the two hydrogen ions released by sulfuric acid exert a strong enough attraction on the unstable atoms of the carbonate ion to pull it apart:
H2SO4 + CaCO3 → H2CO3 + Ca++ + SO4-- → H2O + CO2 + Ca++ + SO4--
The alkaline carbonate ion is thus eliminated, because one mole of sulfuric acid neutralizes one mole of carbonate.
Phosphoric acid (H3PO4) is another mineral acid. Each molecule of phosphoric acid contains three hydrogen ions, but still only has an equivalence of 2, because it only partially disassociates in water and releases only two of its three hydrogen ions:
H3PO4 + CaCO3 → H2CO3 + CaHPO4, precipitated
Two organic acids commonly used by brewers are citric and DL-lactic acid. Citric acid (2C6H8O7) partially disassociates and releases three hydrogen ions in solution, giving an equivalence of 3. Lactic acid (2C3H6O3) has an equivalence of 1, because it releases only one hydrogen ion. Only two-thirds of a mole of citric acid is needed to neutralize a mole of carbonate, but two moles of lactic acid are needed to neutralize a mole of carbonate:
2C6H8O7 + 3CaCO3 → 3H2CO3 + 3Ca++ + 2C6H5O7---
2C3H6O3 + CaCO3 → H2CO3 + Ca++ + 2C3H5O3-
Normality (N) is how the strength of an acid is expressed. It may be given as a decimal or fraction: .02N or N/50. One mole of a 1N compound releases one mole of hydrogen (H+) or hydroxyl (OH-) ion. A .02N acid solution releases .02 or one-fiftieth of a mole of hydrogen, and a 5N acid solution releases five moles of hydrogen. Normality is equal to the molarity of the acid times the equivalence of the acid. So the normality of one mole of the four acids are:
Lactic acid, with an equivalence of 1: 1 x 1M = 1N
Sulfuric acid, with an equivalence of 2: 2 x 1M = 2N
Phosphoric acid, with an equivalence of 2: 2 x 1M =2N
Citric acid, with an equivalence of 3: 3 x 1M = 3N
Reversing the equation, 1N solutions of these acids give:
Lactic acid, 1N: gram mole 90.08/equivalence 1= 90.08 grams C3H6O3/liter of solution
Phosphoric acid, 1N: gram mole 97.995/equivalence 2 = 48.998 grams H3PO4/liter of solution
Sulfuric acid, 1N: gram mole 98.08/equivalence 2 = 49.04 grams H2SO4/liter of solution
Citric acid, 1N: gram mole 192.13/equivalence 3 = 64.043 grams C6H8O7/liter of solution
These values allow for ease of calculations for alkalinity adjustments, since they represent milligrams per milliliter as well as grams per liter, and multiplied by 1,000 they give milligrams per liter, or ppm, the measurement by which alkalinity as CaCO3 is expressed. Alkalinity reductions require acid additions that are equivalent to the molarity of the alkaline cations, and the equivalence of acids is measured by normality.
Mineral Ions Common in Water
Cations — Earths
Calcium (Ca++, atomic weight 40.08). Calcium is the principal mineral of hardness. It comes from the water’s passage over limestone, dolomite, gypsum, or calcified gypsiferous shale. Calcium increases mash acidity and inverts malt phosphate to precipitated alkaline phosphate by the following reaction:
In appropriate amounts, calcium is advantageous to the brew. Calcium stimulates enzyme activity and improves protein digestion, stabilizes the alpha-amylase, helps gelatinize starch, and improves lauter runoff. It also extracts fine bittering principles of the hop and reduces wort color. A calcium precipitate formed with potassium phosphate improves hot-break flocculation. It is also an essential part of yeast-cell composition. Small amounts of calcium neutralize substances toxic to yeast, such as peptone and lecithin. It improves clarification during aging, as well as the stability and flavor of the finished beer.
Precipitation of Calcium Carbonate
Boiling
Adding Slaked Lime
In excess, however, calcium precipitation with organic phosphates interferes with runoff filtering and robs the wort of phosphate, a necessary yeast nutrient. Calcium levels are usually 5 to 200 ppm; its solubility is greatly affected by anions in solution with it.
Magnesium (Mg++, atomic weight 24.32). Magnesium is the secondary mineral of hardness. It is essential as a cofactor for some enzymes, and as a yeast nutrient. In concentrations of 10 to 30 ppm, magnesium accentuates the beer’s flavor, but it imparts an astringent bitterness when it is present in excess. At levels higher than 125 ppm it is cathartic and diuretic. Usually found at levels of 2 to 50 ppm, its solubility is less affected by carbonate anions in solution than is calcium.
Cations — Metals
Sodium (Na+, atomic weight 22.991). The sour, salty taste of sodium can accentuate beer’s flavor when it is found in reasonable concentrations, but it is harsh and unpleasant in excess. It is poisonous to yeast, and brewers generally avoid water that contains sodium in excess of 50 ppm, especially where softness is characteristic of the beer flavor. Usually found at levels of 2 to 100 ppm, it is very soluble.
Potassium (K+, atomic weight 39.1). Potassium imparts a salty taste. In excess of 10 ppm, it inhibits enzyme activity and acts as a laxative. It is difficult to measure and is usually grouped with sodium. Levels seldom exceed 20 ppm, although potassium is very soluble.
Iron (Fe++, atomic weight 55.85, Fe+++). Common in ground water, iron gives an unpleasant, inky taste detectable at levels as low as .05 ppm. Above 1 ppm, iron weakens yeast and increases haze and oxidation of tannins. It blackens porcelain and spots fabrics at .02 ppm, causes white turbidity in water, and corrodes metal. Levels should be less than .3 ppm. Reduce iron content to .1 ppm by aerating and filtering the water through sand.
Manganese (Mn++, atomic weight 54.94). Trace amounts of manganese are found in most ground and surface waters. It imparts an unpleasant taste and streaks porcelain at .05 ppm. The manganese level should be less than 2 ppm and optimally below .05 ppm. It can be reduced to .02 ppm by aeration.
Ammonia (NH4+, atomic weight 18.04). Ammonia is a corrosive ion from microbial organic decomposition. Most volatile of the nitrates, it is reduced by oxidation to the corrosive alkaline gases NH3 and NH2. Always indicative of pollution, ammonia is never present in unpolluted water. Levels of ammonia are normally .00 to .03 ppm and should never exceed .05 ppm.
Copper (Cu++, atomic weight 63.54). Elevated levels of copper cause yeast mutation and haze formation. Copper in a water supply is evidenced by blue-green stains on porcelain. Levels of copper should be less than 1 ppm.
Zinc (Zn++, atomic weight 65.38). Zinc is a yeast nutrient when it is found at .1 to .2 ppm, but toxic to yeast and inhibiting to enzymes above 1 ppm.
Anions
Carbonate (CO3--, atomic weight 60.0092). Carbonate is a strongly alkaline buffer formed by the reaction of atmospheric carbon dioxide with hydroxides of alkaline-earth and alkali metals. Carbonates go into solution as hydrogen carbonates (HCO3-, “bicarbonates”), which are strong buffers. Bicarbonates form by the reaction of a carbonate ion with a molecule each of carbon dioxide and water.
Bicarbonate resists increases in the mash acidity by neutralizing acids as they are formed. It also hinders gelatinization of starch by alpha-amylase, impedes trub flocculation during the cold break, and increases risk of contamination in the ferment. It contributes a harsh, bitter flavor that is overwhelming in delicate lagers. Carbonate in excess of 200 ppm is tolerable only when dark-roasted malts are employed to buffer its excessive alkalinity. Carbonates in the brewing liquor should be less than 50 ppm if the mash is from only pale malts and no liquor acidulation is employed. Where carbonates exceed 50 ppm, water treatment is generally in order.
Sulfate (SO4--, atomic weight 96.0576). Sulfate is weakly basic, and its alkalinity is overcome by most acids. It is fairly soluble. It gives beer a dry, fuller flavor, although the taste can be objectionably sharp. With sodium and magnesium it is cathartic. Above 500 ppm it is strongly bitter, and levels are generally kept at less than 150 ppm unless the beer is very highly hopped. With intensely bitter beers, sulfate at 150 to 350 ppm gives a cleaner, more piquant bitterness.
Chloride (Cl-, atomic weight 35.453). Chloride is very weakly basic, and readily neutralized. It accentuates bitterness, but also increases mellowness; it increases the stability of any solution and improves clarity. The “salt” taste of chloride generally enhances beer flavor and palate fullness, but the salt flavor is reduced by the presence of calcium and magnesium. Usually found at levels of 1 to 100 ppm, chloride levels in the brewing liquor may be as high as 250 ppm for British mild ales.
Silica (SiO2, atomic weight 60.0843), silicon dioxide. Originating from sand or quartz, silica is insoluble. As a colloid, it interferes with the filtering of the mash. Under certain conditions, silica forms silicate (HSiO3-), which causes hazes, precipitating out of the boil as scale with calcium and magnesium. Silica levels are usually less than 10 ppm but may be as high as 60 ppm.
Nitrate (NO3-, atomic weight 62.0049). Nitrate is the most highly oxidized naturally occurring form of nitrogen. It may be from geological strata or originate from contact with sewage or oxidized organic matter. Above 10 ppm, it is indicative of pollution by sewage. It is alkaline, and during fermentation in the presence of chlorides, it forms nitrites, which are more strongly alkaline yet.
Nitrite (NO2-, atomic weight 46.0055). Strongly basic, nitrite originates from nitrates during decomposition of organic matter by coliform bacteria. It rarely exceeds .1 ppm, and is always indicative of pollution. Nitrite is toxic to yeast in minute concentrations; as little as .1 ppm may retard or terminate yeast growth.
Table 6: Water Composition Indicators
| Hardness and Mineral Content of Water | |
|---|---|
|
Nature |
Hardness as CaCO3, ppm |
|
Very soft |
Less than 50 |
|
Soft |
50–100 |
|
Slightly hard |
100–150 |
|
Moderately hard |
150–250 |
|
Hard |
250–350 |
|
Very hard |
350 and above |
|
Multiplying the ions below by the corresponding factors yields hardness as CaCO3. |
|
|
Ca–2.497 Mg–4.116 Fe–1.792 Mn–1.822 Zn–1.531 |
|
| Total Dissolved Solids/Specific Conductivity | ||
|---|---|---|
|
Total Dissolved Solids |
Specific Conductivity |
|
|
Water low in ionized matter |
Below 50 ppm |
Below 90 micromhos/cm |
|
Range of average water supplies |
30–275 ppm |
50–500 micromhos/cm |
|
Very highly mineralized water |
Above 275 ppm |
Above 500 micromhos/cm |
| pH as an Alkalinity and Treatment Indicator | |||
|---|---|---|---|
|
pH |
Alkalinity, as CaCO3 |
||
|
% HCO3 |
% CO3 |
% H2CO3 |
|
|
10 |
68 |
32 |
0 |
|
9 |
95 |
5 |
0 |
|
8 |
97 |
0 |
3 |
|
7 |
81 |
0 |
19 |
|
6 |
30 |
0 |
70 |
|
5 |
4 |
0 |
96 |
|
From table A–1, Principles of Brewing Science, by George Fix. |
|||
Other
Coliform bacteria. Measures the amount of any fecal bacteria, such as Escherichia coli, Streptococcus faecalis, pathogenic Salmonella strains, Shigella dysenteriae, and Vibrio cholerae. The U.S.P.H.S. standard for drinking water is that there should be less than 2.2 colonies per 100 milliliters. For brewing water, it is recommended that this be zero.
Where bacterial population is not given in an analysis, nitrate, nitrite, and ammonia values suffice to indicate water pollution.
Water Treatment
Brewing water sources should be chosen first for their purity and second for their mineral composition. In fact, treatment is only necessary when the mineral distribution of any water is unsatisfactory, or when accentuation of bitterness or saltiness is desired.
The most common correction of brewing water is the reduction of bicarbonate to yield a satisfactory mash acidity/pH. The bicarbonate is alkaline, and if it is reduced, so is the alkalinity.
Where liquor of less than 50 ppm of alkalinity is mashed with pale malt, proper mash acidity is usually realized without liquor treatment, because phosphates dissolved from the malt react with calcium bicarbonate, precipitating calcium phosphate and releasing CO2. With more alkaline water, an excess of carbonate remains in solution, and the mash pH will be too high. Decomposition of the carbonates may be accomplished in several ways.
The most basic is by bringing water to a boil and aerating it thoroughly to decompose bicarbonates to carbonates, which are precipitated as calcium or magnesium carbonate salts, and to decompose carbonic acid to CO2, which is driven off. After a reasonable rest to allow the carbonates to sediment the water should be decanted off the sediment, so that gradual dissolution of atmospheric CO2 back into the water does not result in bicarbonates re-forming from the precipitate.
The use of naturally acidic toasted malt, or a portion of dark-roasted malt, is a time-honored manner of water treatment. The acidity released by intensive kilning can overcome the alkalinity of even moderately alkaline waters.
Where only pale malts are being mashed with soft to moderately alkaline waters (less than 250 ppm alkalinity or 150 ppm HCO3), proper mash acidity is most often achieved by the formation or addition of mild acid. Mixing a portion of sourmalt or lactic-acid mash into the main mash reduces alkalinity and contributes flavor nuances that help round out a beer’s flavor. Sourmalt is made by allowing limited lactic-acid bacterial activity prior to the malt’s being kilned. In the brew house, formation of lactic acid may be accomplished by Lactobacillus delbruckii activity in a partial mash, held closely covered at 95 to 120 degrees F (35 to 50 degrees C) for forty-eight to seventy-two hours, until its pH drops below 4.
Where acids (most commonly lactic, phosphoric, sulphuric, or citric) are used to reduce alkalinity, carbonates are decomposed with the formation of carbonic acid and lactate, phosphate, sulphate, or citrate anions, but the reactions are to some extent reversible. Moreover, if the water is more than moderately alkaline, the excessive amounts of acid required give the beer a noticeable sourness and characteristic taste. Generally, DL-lactic acid is preferred by brewers, but orthophosphoric, monohydrate citric, and sulfuric acids are also commonly used.
Approximate carbonate reduction can be made by gradual acid addition, checking the pH of the liquor after each dose, until it drops below 7. For moderately modified pale malt that won’t undergo acid or protein rests, the pH reduction of the liquor may need to be as low as pH 6.
Carbonate and alkalinity reduction of a water can be more accurately made, where the parts per million of alkalinity as CaCO3 is known, by calculating the treatment beforehand. Some difficulty arises because various dilutions of the acids are offered, and may be expressed as percentage solutions or percentage-of-normality solutions (N). Normality is discussed below. The following quantities of the commonly used acids are equivalent to 1N:
Lactic acid, 1N = 90.08 milligrams C3H6O3 per milliliter of solution
Phosphoric acid, 1N = 48.998 milligrams H3PO4 per milliliter of solution
Sulfuric acid, 1N = 49.04 milligrams H2SO4 per milliliter of solution
Citric acid, 1N = 64.043 milligrams C6H8O7 per milliliter of solution
And as percentage solutions, common dilutions give:
Lactic acid 85 to 90% w/w: 1,020 milligrams per milliliter of solution
Phosphoric acid 85 to 88% w/w: 1,445 milligrams per milliliter of solution
Sulfuric acid 95 to 98% w/w: 1,766 milligrams per milliliter of solution
Citric acid is generally available as the monohydrate, in granular or powder form, and so can be weighed out, each milligram of the monohydrate giving .9143 milligrams of citric acid.
These values allow ease of calculations for alkalinity adjustments. Given a water with alkalinity as CaCO3 of 220 ppm and pH 7, the brewer wants to reduce the alkalinity to below 50 ppm; to, say, 35 ppm: 220 - 35 = 185 ppm. 185 ppm of alkalinity should be removed. Reference to table 6 shows that at pH 7, 81 percent of that alkalinity is bicarbonates, and the remainder is harmless carbonic acid. 185 x .81 = 150 ppm of alkalinity as CaCO3 (at this pH, actually bicarbonate, HCO3) needs to be disassociated. The acid treatment required to accomplish this is predicted by calculating the total alkalinity to be removed; that is, the ppm of alkalinity as CaCO3 times the total volume of liquor.
Where the brewer is using liters as a measure, this is simply done by multiplying the alkalinity as CaCO3 times the number of liters of liquor needed for the brew. Where the brewer is working with gallons, the conversion factor of liters in a gallon (3.7854) needs to be included in the formula:
Alkalinity as CaCO3 x 3.7854 x number of gallons of liquor
Divided by the milligrams per milliliter that the particular acid solution on hand bears, the formula gives the milliliters of acid needed to disassociate the carbonate/bicarbonate alkalinity. For example, for 185 ppm of alkalinity to be disassociated, and 7.5 gallons of water to be treated:
185 x 3.7854 x 7.5 = 5,250 ppm of alkalinity as CaCO3 to be disassociated.
Using 85 percent lactic acid, which bears 1,020 milligrams of lactic acid per milliliter:
5,250/1,020 = 5.2 milliliters of 85 to 90 percent lactic acid will reduce the alkalinity as CaCO3 of 7.5 gallons of water by approximately 185 ppm.
Mineral Salt Treatment
Ion-exchange water softeners should never be used to reduce hardness. They do not remove the carbonate ion from solution, but precipitate calcium and magnesium by exchanging them for more soluble sodium ions, correspondingly increasing the sodium concentration.
| Estimated Characters of the Classic Brewing Waters | |||||||||
|---|---|---|---|---|---|---|---|---|---|
|
Ca |
Mg |
K |
Na |
SO4 |
HCO3 |
Cl |
Hardness |
Total Dissolved Solids |
|
|
Pilsen |
7 |
2 |
2 |
5 |
15 |
5 |
30 |
35 |
|
|
Munich |
75 |
18 |
2 |
10 |
150 |
2 |
250 |
275 |
|
|
Vienna |
200 |
60 |
8 |
125 |
120 |
12 |
750 |
850 |
|
|
Dortmund |
225 |
40 |
60 |
120 |
180 |
60 |
750 |
1000 |
|
|
London |
90 |
5 |
15 |
40 |
125 |
20 |
235 |
300 |
|
|
Dublin |
120 |
5 |
12 |
55 |
125 |
20 |
300 |
350 |
|
|
Yorkshire |
100 |
15 |
25 |
65 |
150 |
30 |
275 |
400 |
|
|
Edinburgh |
120 |
25 |
55 |
140 |
225 |
65 |
350 |
650 |
|
|
Burton |
275 |
40 |
25 |
450 |
260 |
35 |
875 |
1100 |
|
Iron, manganese, and colloids that cause hazes are best removed by aeration, followed by filtration or sedimentation.
Mineral salts may be added to the brewing water when additional hardness or other mineral character is desired, or to precipitate carbonates. All salts should be carefully weighed (on a gram scale for small batches) before they are added to the brewing water. Mineral salts cannot be accurately dispensed by volume. One level teaspoon of finely powdered gypsum might weigh 3.65 grams. Tightly packed, it weighs 5 grams. A teaspoon of more crystalline magnesium sulfate weighs 4.55 grams, finely granular potassium chloride 5.05 grams, and sodium chloride 6.45 grams. For accuracy, salts need to be measured by weight.
It is advisable to first mix salts into a small quantity of boiling water before introducing them to the brewing water. Salts should never be added directly to the mash because uniform dispersal is unlikely.