# Water hardness

Faucet with jet regulator : The hardness of the tap water is visible here. The lime has settled on the dripping tap .

Water hardness is a system of terms in applied chemistry that has developed from the needs of using natural water with its dissolved ingredients.

In water chemistry , water hardness denotes the substance concentration of the alkaline earth metal ions dissolved in the water , and in special contexts also their anionic partner. These “hardness builders” essentially include calcium and magnesium ions as well as strontium and barium ions, which are normally only contained in traces . These dissolved hardness builders can form poorly soluble compounds, especially carbonates and lime soaps . This tendency towards the formation of insoluble compounds is the reason for paying attention to the dissolved alkaline earths, which has led to the development of the concept and theory system around water hardness.

## History and derivation of the term

Already in the ancient Hippocratic writings a distinction was made between soft (μαλακός) and hard (σκληρός) water. The hard water can be found among other things “near rock springs, warm earth or mineral-rich springs.” Galenos explains: “Because Hippocrates calls 'hard water' the rough water that bites the tongue when drinking and the body when washing. The soft water is the opposite of that. "

To this day, rainwater, which is considered soft, is preferably used for washing laundry by hand. Spring or well water, which serve as reliable hard, be shunned, however, because the high concentration of dissolved minerals in hard water soap increases to water-insoluble calcium soap can coagulate. The part of the soap bound in this way loses its cleaning effect. At the same time, the resulting lime soap turns the washed items gray and makes them hard and stiff after drying on the clothesline. By using soft water for washing, these undesirable effects can largely be avoided.

## Effects

Rainwater is " distilled water " and by nature does not contain any minerals, but only air components and air pollutants that were washed out of the air or condensed themselves when falling to the surface of the earth . That is why rainwater is soft water . In regions with crystalline rocks in the ground, such as granite , gneiss and basalt , the rainwater can only dissolve a few easily soluble minerals, the groundwater is soft water. Also surface water without much contact with rocks is considered soft.

Soft water is cheaper for all applications,

The disadvantage, however, can be strong foaming in detergents and the poor removability of soap z. B. when washing hands.

In contact with calcareous rocks (such as limestone , marble or dolomite ), rainwater can dissolve more minerals and it becomes hard water .

## Emergence

The water hardness is formed during passage of water through carbonatgesteinhaltige soil and rocks and / or aquifer (aquifers) by dissolution of carbonates by using carbon dioxide under formation of soluble bicarbonates (HCO 3 - ).

All dissolved alkaline earth metals (which are then present as carbonates , sulphates , chlorides , nitrites , nitrates and phosphates ) are referred to as total hardness , the parts only bound to carbonic acid as carbonate hardness (also carbonate hardness or temporary hardness or temporary hardness ) and the difference between them as non-carbonate hardness ( permanent hardness or permanent hardness ).

The major part of the water hardness arises as carbonate hardness and is therefore of special importance for the water hardness. It arises from the dissolution of rocks containing carbonate, i.e. lime (CaCO 3 ) or dolomite (Ca-Mg mixed carbonate) according to the following formulas

${\ displaystyle \ mathrm {CaCO_ {3} + CO_ {2} + H_ {2} O \ rightleftharpoons Ca ^ {2 +} + 2 \ HCO_ {3} ^ {-}}}$
${\ displaystyle \ mathrm {MgCO_ {3} + CO_ {2} + H_ {2} O \ rightleftharpoons Mg ^ {2 +} + 2 \ HCO_ {3} ^ {-}}}$

The same reactions and equilibria take place with all alkaline earth and mixed carbonates: SrCO 3  , BaCO 3  , ... The carbonate hardness corresponds to the concentration of the anion hydrogen carbonate (HCO 3 - ).

Magnesium and calcium ions can also get into the water through other dissolution processes, for example through the dissolution of gypsum minerals (CaSO 4 × 2 H 2 O). In extreme cases, groundwater from layers containing gypsum can reach the saturation concentration for gypsum, which corresponds to a hardness of 78.5 ° fH or 44 ° dH. (For the units of measurement ° dH and ° fH, see section Units and conversion below .)

The acids contained in acid precipitation , which have become known through the term acid rain , lead to an increase in total hardness after carbonate rocks have been dissolved. Mainly involved are sulfuric acid (H 2 SO 4 ), which is formed via sulfur dioxide and the formation of sulphurous acid when burning fuels containing sulfur, and nitric acid (HNO 3 ), which is formed via the intermediate stage of nitrogen oxides in particularly hot burns. By measures to air pollution (. Eg flue gas desulphurisation and automotive catalysts and DeNOx systems in power plants), these strains have been drastically reduced in recent decades.

When plant matter decays (dead roots, fallen leaves, plowed stalks) in the soil or when agricultural fertilizer is applied, the nitrogen contained therein is initially released as ammonium (NH 4 + ). This is followed by a bacterial oxidation process , the so-called nitrification . The ammonium is first oxidized to nitric acid (HNO 2 ) and finally to nitric acid (HNO 3 ) (and could also be further denitrified to N 2 ). This nitric acid dissolves hardness from lime - and in the absence of lime from clay minerals - which is then no longer available to the plants. Therefore, agriculturally used lime-poor soils threaten to acidify. In these cases, lime fertilization is necessary. The carbonates that are then abundant again can be partly responsible for an increase in hardness in groundwater.

In groundwaters that are influenced by agricultural activities, the hardness can rise to over 30 ° fH or 17 ° dH, in individual cases even to over 40 ° fH or 23 ° dH. This is due to both increased carbon dioxide formation and increased nitrification.

Rainwater can only absorb hardness builders in exceptional cases when the atmosphere contains calcareous dust particles. This is why the hardness of rainwater is usually close to zero. Even drinking water reservoirs and mountain lakes often contain water of low hardness, even in limestone-rich areas, if their catchment area covers a small geographical area and the rainwater mainly flows in superficially.

### Lime-carbonic acid balance

The water hardness is characterized by a system of different coupled chemical equilibrium reactions and depends on it. In addition to the reaction equilibria , this also includes the solubility equilibria between the various alkaline earth ions and the associated carbonate and sulfate precipitation products ( calcite , dolomite , barite , gypsum, etc.). The solution and dissociation equilibrium of the carbon dioxide-carbonic acid-carbonate system is also coupled.

Meets CO 2 at (rain) water are dissolved more than 99% of the carbon dioxide only physically and less than 1% responding depending on temperature in an equilibrium reaction ( there explained) with the water molecules chemically to form carbonic acid (H 2 CO 3 ), the aqueous solution reacts therefore weakly sour.

${\ displaystyle \ mathrm {\ CO_ {2} + H_ {2} O \ \ rightleftharpoons \ H_ {2} CO_ {3}}}$.

The carbonic acid H 2 CO 3 formed is in an equilibrium reaction with hydrogen carbonate (HCO 3 - ) ions and oxonium (H 3 O + ) ions.

${\ displaystyle \ mathrm {\ H_ {2} CO_ {3} + H_ {2} O \ \ rightleftharpoons \ HCO_ {3} ^ {\, -} + H_ {3} O ^ {+}}}$

The hydrogen carbonate ion HCO 3 - dissociates further in water to form the carbonate ion CO 3 2-

${\ displaystyle \ mathrm {\ HCO_ {3} ^ {\, -} + H_ {2} O \ \ rightleftharpoons \ CO_ {3} ^ {\, 2 -} + H_ {3} O ^ {+}}}$

In water these equilibrium reactions are predominantly on the side of the carbon dioxide (which is then predominantly physically dissolved in the water) and hydrogen carbonate ions are only formed to a small extent.

In water chemistry , dissolved CO 2 is usually combined with the actual acid H 2 CO 3 as free carbonic acid , its equilibrium reaction products, the sum of carbonate and hydrogen carbonate, as bound carbonic acid .

The formula

${\ displaystyle \ mathrm {Ca ^ {2 +} \ + \ 2 \ HCO_ {3} ^ {-} \ rightleftharpoons CaCO_ {3} \ + \ H_ {2} CO_ {3}}}$

describes the lime-carbonic acid balance .

Calcium carbonate itself is hardly soluble in pure water. The solubility is just 14 milligrams per liter, with the carbonate ion going into solution as hydrogen carbonate ion. In the presence of dissolved carbon dioxide, however, the solubility increases more than a hundredfold, with the easily soluble (dissociated) calcium hydrogen carbonate being formed

If one of these equilibrium reactants is added to water (or removed from the water), there is an "overhang" (or "deficiency") on one side in equilibrium, whereupon the chemical reactions run in the other (from the other) direction of the reaction equations an equilibrium has been re- established with the reaction products ( principle of the smallest constraint ). The same applies if the pH value or the water temperature are changed. Water evaporation or dilution also cause a change in the CO 2 concentration and thus a shift in equilibrium. Rainwater absorbs CO 2 from the atmosphere depending on the temperature . An overhang (or deficiency) on one side of the equilibrium reaction leads to an increase (or decrease) on the other side as well, for example ...

From the Hägg diagram for the lime-carbonic acid balance, the concentrations of CO 2  , HCO 3 - and CO 3 2- can be read off depending on the respective pH value or a change in pH
value can be predicted depending on the concentrations. At low pH values ​​mainly carbonic acid and carbon dioxide are present in the water, at pH 8 almost exclusively hydrogen carbonate ions, and at high pH values ​​carbonate ions predominate.

The water is in this lime-carbonic acid equilibrium when the lime precipitation is the same as the lime solution, i.e. it contains just enough carbon dioxide that it does not separate any lime, but it cannot dissolve any lime either . If carbon dioxide is removed from such water, poorly soluble compounds such as calcite and dolomite are formed as particularly poorly soluble (!) Mixed carbonate. The water is then oversaturated with these mixed carbonates , whereupon they precipitate .

### Separation and precipitation of lime and mixed minerals

In addition, there is the carbonate-silicate cycle of rocks, in which silicate rocks are dissolved and then deposited again. The carbonated rainwater erodes silicate rocks by dissolving calcium silicate minerals (compounds of calcium , silicon and oxygen ), whereby the released calcium and hydrogen carbonate ions get into the groundwater. The equation for the conversion of the feldspar anorthite by carbonic acid to form kaolinite serves as an example :

${\ displaystyle \ mathrm {2 \; H_ {2} CO_ {3} \ + Ca [Al_ {2} Si_ {2} O_ {8}] \ + H_ {2} O \ \ rightarrow \ Al_ {2} [ (OH) _ {4} | Si_ {2} O_ {5}] \ + Ca ^ {2 +} \ +2 \; HCO_ {3} ^ {-} \}}$

Mineral biominerals are excreted from the water through further biomineralization . For example, chandelier algae or cyanobacteria capable of photosynthesis ("blue-green algae") also precipitate calcium carbonate, the latter forming mat-shaped stromatolites in microbe mats . The microorganisms in the biofilms are then inactive at the base and die and continue to grow on the film surface. Diatoms ( diatoms ) precipitating silica from the water and form therefrom at normal temperature and normal pressure hydrous amorphous silicon dioxide .

Alleged "limescale deposits" in swimming ponds or swimming pools therefore usually consist of homogeneous mixtures of calcium carbonate, mixed carbonates, apatite, silicon dioxide and silicates and are therefore difficult to dissolve even with acids. Sedimentation of these mineral layers takes place in bodies of water , see Mudde - sludge deposits ( Mulm ) and sea ​​chalk - lime deposits.

Due to the temperature dependence of the entire equilibrium system, scale deposits also form, for example, in hot water systems, coffee machines or cooking pots when hot water is prepared from calcareous water.

### Buffer capacity of such waters

High lime contents and thus high hydrogen carbonate contents in the water also act as chemical buffers (with carbonic acid as acid and hydrogen carbonate ion as base ).

${\ displaystyle {\ rm {CO_ {2} + 2H_ {2} O \ rightleftarrows H_ {2} CO_ {3} + H_ {2} O \ rightleftarrows HCO_ {3} ^ {-} + H {_ {3} } O ^ {+}}}}$

When acid is added, the chemical equilibrium is established by releasing CO 2 . A lot of acid has to be added so that the pH value changes significantly, which at the same time represents a protective function for the biocenosis against strong pH value fluctuations.

To determine the buffer capacities , the acid capacity is determined on the one hand , how much acid is needed to be added to a water sample until the pH value 4.3 is reached, so that the bound carbonic acid is determined. The excess carbon dioxide is the base capacity determines how much the addition of alkali to a water sample is necessary until it reaches pH 8.2.

### "Aggressive Water"

If tap water in houses is decalcified using ion exchange systems , this water can absorb more carbon dioxide. In water chemistry, water that can still dissolve minerals (see also solubility and solubility product ) is called aggressive water .

Calcareous materials (concrete, asbestos cement, etc.) are also corroded if there is an excess of CO 2 in the water . The excess carbonic acid has an influence on the likelihood of corrosion of materials, and this particularly applies to materials made of iron, if so-called aggressive carbonic acid is present. Equilibrium waters with a minimum hardness of 1.5 mmol / l tend to form a lime-rust protection layer. See also Tillmans' formula .

Aggressive water with excess carbonic acid has the consequence that limescale deposits in old water pipes are removed again and rust can occur in damaged areas of the inner zinc layer . To prevent this, after decalcification, phosphates are often added to the tap water as phosphating , which again create a coating as corrosion protection and prevent limescale detachment.

From dolomite (CaMg (CO 3 ) 2 ) is at a lower Kalkbrenn temperature- calcined dolomite (CaCO 3 · MgO; also Magno called) were prepared. This has proven itself in drinking water treatment as silica-free (SiO 2 -free) filter granules for deacidification to bind excess CO 2 . The material is also used in fish and swimming ponds to increase and stabilize the pH value and for more effective liming. The MgO component reacts preferentially with water. More under Magno (chemical) .

## Importance of total hardness

The total hardness indicates the sum of the concentrations of the cations of alkaline earth metals in water. These cations have a great, positive physiological importance , but interfere with some uses of the water.

### Physiological importance

Magnesium and calcium are essential for the organism. The human body contains 0.47 g / kg magnesium and 15 g / kg calcium. However, drinking water plays a subordinate role in supplying the body with these elements. Strontium, like calcium, is contained in the bones, but has no special physiological significance.

Barium is poisonous in dissolved form. In waters containing sulphate, toxicologically questionable concentrations are not reached because barium sulphate, which is extremely difficult to dissolve , is formed. Barium sulfate is the main component of orally administered medical X-ray contrast media .

## Non-carbonate hardness

The permanent hardness is not bound to hydrogen carbonate or carbonate and can therefore in principle not be removed from the water as calcium or magnesium carbonate. This non-removable part is due to anions such. B. Chlorides , nitrates and sulfates balanced ("bound"). The exact different concentrations in which these anions are present does not play a role in relation to the water hardness, but provides information about the origin of these components. In fact, however, this permanent hardness has a decisive influence on the precipitation behavior of the carbonate hardness components, because the resulting increased concentrations of calcium and magnesium are included in the calculation of the ion products with the carbonate and thus the threshold values ​​e.g. B. increase the "associated carbonic acid" for the occurrence of the hardness-typical precipitation reactions.

The concentrations of magnesium and calcium ions are also often determined separately and then referred to as “magnesium hardness” or “calcium hardness”. Their sum corresponds to a good approximation of the total water hardness.

## Methods of determination

• The best-known practicable method of determining the total hardness is the complexometric titration with an aqueous solution of the disodium salt of ethylenediaminetetraacetic acid (EDTA, trade name: Titriplex III) with a known concentration . EDTA forms with all alkaline earth metal hardness components Ca 2+ , Mg 2+ , Ba 2+ , Sr 2+ , ... soluble, stable chelate . 100 ml of the water sample to be examined are mixed with 2 ml of 25% ammonia solution, a pH 11 buffer ( ammonia - ammonium acetate ) and the indicator eriochrome black T. The indicator is usually available together with the buffer as so-called “indicator buffer tablets”. The indicator, when masked with a yellow dye, forms a red colored complex with the Ca 2+ and Mg 2+ . If all alkaline earth ions are bound by the EDTA at the end of the titration, the Eriochrome Black-T is free and colored green. The unmasked indicator changes color from magenta to blue. The total hardness is calculated from the volume of EDTA solution used. For a water sample of 100 ml, 1 ml of used EDTA solution (c = 0.1 mol / l) corresponds to 5.6 ° dH (German degrees of hardness), which corresponds to 1 mmol / l of alkaline earth metal ions. In order to determine the calcium and magnesium concentration individually, titration is carried out against Ca 2+ with EDTA at a lower pH of approx. 8 , because at this pH the Mg-EDTA complex is not yet stable. At the transition point for calcium, the pH is then adjusted to 11 and titrated with EDTA against Mg 2+ .
• A somewhat older method is hydrolytic precipitation titration with alcoholic potassium palmitate solution , in which palmitations react with calcium and magnesium ions (or all alkaline earth ions that are also included) to form the corresponding insoluble salts ( lime soap ) of palmitic acid. When the equivalence point is exceeded, palmitate ions react hydrolytically to form hydroxide ions , which are detected using phenolphthalein as an indicator . 1 ml of a potassium palmitate solution with a concentration of 0.1 mol / l corresponds to a total hardness of 1 meq / l.
${\ displaystyle \ mathrm {2 \ RCOO ^ {-} + Ca ^ {2 +} \ \ longrightarrow \ (RCOO) _ {2} Ca}}$
${\ displaystyle \ mathrm {\ RCOO ^ {-} + H_ {2} O \ \ rightleftharpoons \ RCOOH + OH ^ {-}}}$
• The carbonate hardness is determined by the hydrochloric acid binding capacity (SBV), the carbonate hardness corresponds to the acid capacity (see also buffer capacity ). For this purpose, 100 ml of the water are titrated with hydrochloric acid (c = 0.1 mol / l) up to pH 4.3 ( pH meter or change of methyl orange indicator). Here (almost) all carbonate and hydrogen carbonate is converted to "free carbonic acid". The acid consumption in ml therefore corresponds to the hydrogen carbonate concentration in meq / l. Multiplication by 2.8 results in German degrees of hardness (° dH), provided the result of the calculation does not exceed the total hardness. The determination of free carbonic acid is determined by determining the base capacity .
If the analysis of a natural water results in a higher value for the carbonate hardness than for the total hardness, then this water also contains sodium hydrogen carbonate. In this case, the carbonate hardness is identical to the total hardness, as this cannot be greater than the total hardness.
• In analytical laboratories the alkaline earth ions as well as the anions of the acid residues can be determined with the help of ion chromatography or capillary electrophoresis . Calcium can also be determined spectroscopically with the aid of flame atomic emission spectrometry (F-AES).

## Units and conversion

According to the SI system of measurement , the content of alkaline earth ions , i.e. the total hardness, is given in moles per liter or, given the low concentrations, in millimoles per liter (mmol / l).

In Germany and Austria, water hardness was previously given in degrees of German hardness (° dH). 1 ° dH was formally defined as 10 mg CaO per liter of water. The other hardness constituents such as magnesium were defined as an amount equivalent to this (7.19 mg MgO per liter). Later, the indication of water hardness in the practice-oriented substance quantity equivalent unit millival per liter (mEq / l) was used. Today the above-mentioned molar information is required by law, regardless of the practical requirements.

In Switzerland, the French degrees of hardness ° fH are decisive.

Other units of measurement were or are in use in other countries, but these are only comparable to a limited extent. They become comparable if one assumes a standard ion ratio. This is possible because most natural waters have a relatively similar cation distribution, regardless of the total salt content. The following table can only be used for conversion if this is the case:

Conversion for the units of water hardness
° dH ° e (° Clark) ° fH ° rH ppm (° aH) mval / l mmol / l
German degree 1 ° dH = 1 1.253 1.78 7.118 17.8 0.357 0.1783
English degree
(degree Clark)
1 ° e = 0.798 1 1.43 5.695 14.3 0.285 0.142
French degree 1 ° fH = 0.560 0.702 1 3.986 10 0.2 0.1
Russian degree 1 ° rH = 0.140 0.176 0.251 1 0.146 0.050 0.025
ppm CaCO 3
1 ppm = 0.056 0.07 0.1 6.834 1 0.02 0.01
mval / l alkaline earth ions 1 meq / l = 2.8 3.51 5.00 20.040 50 1 0.50
mmol / l alkaline earth ions 1 mmol / l = 5.6 7.02 10.00 40.080 100.0 2.00 1

The unit 1 ppm is used here, contrary to the actual sense of the word, in the sense of 1 mg CaCO 3 per liter of water, i.e. H. in the sense of around 1 mg per kilogram.

If the values ​​for magnesium (Mg) and calcium (Ca) are known, the hardness of the water (e.g. for mineral water) can be calculated as follows:

Degree of hardness of the water in ${\ displaystyle \ mathrm {mmol / l} \ approx [{\ text {Ca value in mg / l}}] / 40 + [{\ text {Mg value in mg / l}}] / 24 {,} 3}$
or in ${\ displaystyle {\ text {° dH}} \ approx 0 {,} 14 \ cdot [{\ text {Ca value in mg / l}}] + 0 {,} 23 \ cdot [{\ text {Mg- Value in mg / l}}]}$

## Hardness ranges

### Germany

#### Hardness ranges for the dosage of detergents

In accordance with Section 7 (1), sentence 1, no.5 of the Washing and Cleaning Agents Act (WRMG), since 1988 graduated dosage recommendations in milliliters for hardness ranges 1 to 4 have had to be given on the packaging of detergents and cleaning agents that contain phosphates or other hardness-binding substances . The information regarding millimoles of total hardness per liter was prescribed by law . The following hardness ranges have been defined:

Hardness range Millimoles of total hardness per liter ° dH
1 (soft) to 1.3 to 7.3
2 (medium) 1.3 to 2.5 7.3 to 14
3 (hard) 2.5 to 3.8 14 to 21.3
4 (very hard) over 3.8 over 21.3

#### New regulation of the hardship areas

On February 1, 2007, the German Bundestag passed the new version of the Washing and Cleaning Agents Act (WRMG), which came into force on May 5, 2007. In it were u. a. the hardness ranges are adapted to European standards and the specification millimoles total hardness per liter is replaced by the specification millimoles calcium carbonate per liter (which is nonsensical from a chemical point of view) . Water supply companies will probably continue to publish the total hardness, but this is not required by law. According to statements by the BMU to the DVGW , millimoles of calcium carbonate per liter should be interpreted as millimoles of total hardness per liter . The new hardness areas hardly differ from the previous ones, only areas 3 and 4 are merged to form the hardness area "hard" and the numbers 1, 2, 3 and 4 are replaced by the - already common - descriptions "soft", "medium" and " hard “replaced. The new hardness ranges are defined as follows:

Hardness range Millimoles of calcium carbonate per liter ° dH
soft less than 1.5 less than 8.4 ° dH
medium 1.5 to 2.5 8.4 to 14 ° dH
hard more than 2.5 more than 14 ° dH

According to Section 8 Paragraph 1 Clause 1 WRMG, recommended quantities and / or dosage instructions in milliliters or grams for a normal washing machine filling with soft, medium and hard water hardness levels and taking into account one or two wash cycles must be indicated on the packaging of detergents . In order to save detergent, you have to know the local water hardness and then read off the corresponding amount of detergent on the packaging. With harder drinking water (from hardness range 3 - "hard"), a separate, phosphate-free softener should be added at temperatures of 60 ° C and above. The water supply companies inform the customer of the local water hardness or send stickers, which are expediently stuck on the washing machine.

For drinking water there are regulations regarding water hardness, see there.

### Switzerland

According to the food law, water in Switzerland is divided into six degrees of hardness, which are given in millimoles per liter (number of calcium and magnesium particles per liter of water) or in French degrees of hardness ºfH.

Hardness in ° fH mmol / l designation
0 to 7 0 to 0.7 very soft
7 to 15 0.7 to 1.5 soft
15 to 25 1.5 to 2.5 medium hard
25 to 32 2.5 to 3.2 pretty hard
32 to 42 3.2 to 4.2 hard
greater than 42 greater than 4.2 very hard

While the water in the foothills of the Alps , in the Alps and on the southern side of the Alps is usually very soft or soft, it is medium-hard in the Jura and hard to very hard in the Central Plateau.

## Softening methods

Decarbonisation : This measure only reduces the carbonate hardness. Calcium hydroxide is added to the water as "lime water", which triggers the following reaction:

${\ displaystyle \ underbrace {\ mathrm {Ca ^ {2 +} + 2 \ HCO_ {3} ^ {-}}} _ {\ text {Carbonate hardness}} \, \ mathrm {+ \, Ca (OH) _ { 2}} \ \ mathrm {\ longrightarrow 2 \ CaCO_ {3} \ downarrow \ +2 \ H_ {2} O}}$

In some German waterworks, decarbonization is carried out on very hard water.

Softening through ion exchange: ion exchangers that are regenerated with common salt are able to exchange calcium and magnesium ions for sodium ions. This principle is z. B. used in dishwashers to protect the heating elements and to avoid "limescale" on the dishes. Water softening systems for non-professional use for softening drinking water use this principle. Occasionally it is also used to treat small amounts of water, for example for watering flowers or making tea.

Full desalination: Full desalination not only removes hardness builders, but also all ions. It is achieved through a combination of cation and anion exchangers. Fully demineralized water is used wherever water is needed in its pure form. The largest amounts are used as boiler feed water . Reverse osmosis and distillation , which also remove non-ionic dissolved solids, achieve a similar result .

Other methods: The complex formation with polyphosphates reduces the hardness, but leads to over-fertilization of surface waters. Detergents often contain small amounts of complexing agents, but the softening is nowadays mainly by cation exchangers such as zeolite A . This prevents the formation of lime soaps , increases the stability of the emulsion required for the wash cycle and protects the heating elements of the washing machine.

When steam locomotives also is inside feed water treatment used.

Devices with electric or magnetic fields do not remove the hardness, and their effect is controversial. At best, it is conceivable that during the crystallization of the excess calcium carbonate under the influence of these fields, the unstable aragonite form is formed, which consists of fine needle-shaped crystals and remains suspended . The normal crystallization to the more stable calcite, on the other hand, forms the well-known incrustations (scale). The effect of this type of water treatment is limited in time and is therefore lost again after a certain flow distance behind the device. A prerequisite for the described effect seems to be that alternating fields are used or that water is swirled in a static field. Therefore z. B. magnetic shoes placed on the water pipe without any effect.

Alternatively prevent sacrificial anode systems on zinc-based crystallization of calcite on surfaces. Released zinc ions react with the carbonate ions dissolved in the water and form crystallization nuclei, from which only weakly crystalline and non-adherent minerals arise directly in the water phase. These are then transported away with the flow of water and their abrasive effect reduces existing deposits. It is not really about removing the hardness of the water, but it is about avoiding limescale deposits.

## literature

• Walter Kölle: Water analyzes - judged correctly. Basics, parameters, water types, ingredients, limit values ​​according to the Drinking Water Ordinance and the EU Drinking Water Directive. 2nd updated and expanded edition. WILEY-VCH, Weinheim 2003, ISBN 3-527-30661-7 .
• Hanns-J. Krause: aquarium water. Diagnosis, therapy, processing. 2nd improved edition, new edition. bede-Verlag, Kollnburg 1993, ISBN 3-927997-00-5

## Individual evidence

1. Anne Liewert: The meteorological Medicine of Corpus Hippocraticum. De Gruyter, Berlin / Munich / Boston 2015, ISBN 978-3-11-041699-2 , p. 98, Google Books view
2. Anne Liewert: The meteorological Medicine of Corpus Hippocraticum. De Gruyter, Berlin / Munich / Boston 2015, ISBN 978-3-11-041699-2 , p. 98, footnote 109, Google Books view
3. Wilhelmine Buchholz: water and soap, or, general laundry book. ; Hamburg and Leipzig; 1866 ( limited preview in Google Book search).
4. ^ A b c Karl Höll, Helmut Peter, Dietrich Lüdemann: Water . ISBN 3-11-125936-6 ( limited preview in Google Book Search).
5. Applied chemistry and environmental technology for engineers. P. 340 ( limited preview in Google Book search).
6. Wolfgang F. Tegethoff: Calcium Carbonate From the Cretaceous Period into the 21st Century . Springer-Verlag, 2013, ISBN 978-3-0348-8259-0 , pp. 3 ( limited preview in Google Book Search).
7. ^ Harry H. Binder: Lexicon of the chemical elements. S. Hirzel Verlag, Stuttgart 1999, ISBN 3-7776-0736-3 .
8. Jander / Jahr / Knoll: Maßanalyse, Göschen Collection Volume 221, de Gruyter Berlin 1966, pp. 209 ff.
9. Table based on: Hanns-J. Krause: aquarium water. Diagnosis, therapy, processing. 2nd improved edition, new edition. bede-Verlag, Kollnburg 1993, ISBN 3-927997-00-5 , page 35.
10. from 1987, Federal Law Gazette I p. 875
11. i. Verb. M. of Regulation (EC) No. 648/2004 (PDF) (last amended by EC Regulation No. 907/2006 ) Art. 11 Paragraph 4 and in accordance with Annex VII Section B thereof.
12. Swiss Gas and Water Association: Water hardness: What must be considered? (PDF; 333 kB), accessed on May 8, 2017.