Wikipedia

Hard water

(Redirected fromSoft water)
Not to be confused withHeavy waterorIce.

Hard wateriswaterthat has a highmineralcontent (in contrast with "soft water" ). Hard water is formed when waterpercolatesthrough deposits oflimestone,chalkorgypsum,[1]which are largely made up ofcalciumandmagnesiumcarbonates,bicarbonatesandsulfates.

Abathtubfaucetwith built-up calcification from hard water in Southern Arizona.

Drinking hard water may have moderate health benefits. It can pose critical problems in industrial settings, where water hardness is monitored to avoid costly breakdowns inboilers,cooling towers,and other equipment that handles water. In domestic settings, hard water is often indicated by a lack offoamformation whensoapis agitated in water, and by the formation oflimescalein kettles and water heaters.[2]Wherever water hardness is a concern,water softeningis commonly used to reduce hard water's adverse effects.

Origins

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Natural rainwater, snow and other forms ofprecipitationtypically have low concentrations ofdivalentcationssuch as calcium and magnesium. They may have small concentrations of ions such assodium,chlorideandsulfatederived from wind action over the sea. Where precipitation falls in drainage basins formed of hard, impervious and calcium-poor rocks, only very low concentrations of divalent cations are found and the water is termedsoft water.[3]Examples includeSnowdoniain Wales and the Western Highlands in Scotland.

Areas with complex geology can produce varying degrees of hardness of water over short distances.[4][5]

Types

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Permanent hardness

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The permanent hardness of water is determined by the water'sconcentrationofcationswith charges greater than or equal to 2+. Usually, the cations have a charge of 2+, i.e., they aredivalent.Common cations found in hard water include Ca2+and Mg2+,which frequently enter water supplies by leaching from minerals withinaquifers.Commoncalcium-containing minerals arecalciteandgypsum.A commonmagnesiummineral isdolomite(which also contains calcium).Rainwateranddistilledwater aresoft,because they contain few of theseions.[3]

The followingequilibrium reactiondescribes thedissolvingand formation ofcalcium carbonateandcalcium bicarbonate(on the right):

CaCO3(s) + CO2(aq) + H2O (l) ⇌ Ca2+(aq) + 2HCO−
3(aq)

The reaction can go in either direction. Rain containing dissolved carbon dioxide can react with calcium carbonate and carry calcium ions away with it. The calcium carbonate may be re-deposited as calcite as the carbon dioxide is lost to the atmosphere, sometimes formingstalactitesandstalagmites.

Calcium and magnesium ions can sometimes be removed by water softeners.[6]

Permanent hardness (mineral content) is generally difficult to remove byboiling.[7]If this occurs, it is usually caused by the presence ofcalcium sulfate/calcium chlorideand/ormagnesium sulfate/magnesium chloridein the water, which do not precipitate out as thetemperatureincreases. Ions causing the permanent hardness of water can be removed using a water softener, orion-exchangecolumn.

Temporary hardness

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See also:Carbonate hardness

Temporary hardness is caused by the presence ofdissolvedbicarbonateminerals(calcium bicarbonateandmagnesium bicarbonate). When dissolved, these types of minerals yield calcium and magnesiumcations(Ca2+,Mg2+) and carbonate andbicarbonateanions(CO2−
3andHCO−
3). The presence of the metal cations makes the water hard. However, unlike the permanent hardness caused bysulfateandchloridecompounds,this "temporary" hardness can be reduced either by boiling the water or by the addition oflime(calcium hydroxide) through the process oflime softening.[8]Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.

Effects

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With hard water, soap solutions form a white precipitate (soap scum) instead of producinglather,because the 2+ ions destroy thesurfactantproperties of the soap by forming a solid precipitate (the soap scum). A major component of such scum iscalcium stearate,which arises fromsodium stearate,the main component ofsoap:

2 C17H35COO−(aq) + Ca2+(aq) → (C17H35COO)2Ca (s)

Hardness can thus be defined as the soap-consuming capacity of a water sample, or the capacity of precipitation of soap as a characteristic property of water that prevents the lathering of soap. Syntheticdetergentsdo not form such scums.

A portion of the ancient RomanEifel Aqueductin Germany. After being in service for about 180 years, the aqueduct had mineral deposits of up to 20 cm (8 in) thick along the walls.

Because soft water has few calcium ions, there is no inhibition of the lathering action of soaps and nosoap scumis formed in normal washing. Similarly, soft water produces no calcium deposits inwater heatingsystems.

Hard water also forms deposits that clog plumbing. These deposits, called "scale",are composed mainly ofcalcium carbonate(CaCO3),magnesium hydroxide(Mg(OH)2), andcalcium sulfate(CaSO4).[3]Calcium and magnesium carbonates tend to be deposited as off-white solids on the inside surfaces of pipes andheat exchangers.This precipitation (formation of an insoluble solid) is principally caused by thermal decomposition of bicarbonate ions but also happens in cases where the carbonate ion is at saturation concentration.[9]The resulting build-up of scale restricts the flow of water in pipes. In boilers, the deposits impair the flow of heat into water, reducing the heating efficiency and allowing the metal boiler components to overheat. In a pressurized system, this overheating can lead to the failure of the boiler.[10]The damage caused by calcium carbonate deposits varies according to the crystalline form, for example,calciteoraragonite.[11]

The presence ofionsin anelectrolyte,in this case, hard water, can also lead togalvanic corrosion,in which one metal will preferentiallycorrodewhen in contact with another type of metal when both are in contact with an electrolyte. The softening of hard water by ion exchange does not increase itscorrosivityper se.Similarly, where lead plumbing is in use, softened water does not substantially increaseplumbo-solvency.[12]

In swimming pools, hard water is manifested by aturbid,or cloudy (milky), appearance to the water. Calcium and magnesium hydroxides are both soluble in water. The solubility of the hydroxides of the alkaline-earth metals to which calcium and magnesium belong (group 2 of the periodic table) increases moving down the column. Aqueous solutions of these metal hydroxides absorb carbon dioxide from the air, forming insoluble carbonates, and giving rise to turbidity. This often results from thepHbeing excessively high (pH > 7.6). Hence, a common solution to the problem is, while maintaining the chlorine concentration at the proper level, to lower the pH by the addition ofhydrochloric acid,the optimum value is in the range of 7.2 to 7.6.

Softening

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Main article:Water softening

In some cases it is desirable to soften hard water. Most detergents contain ingredients that counteract the effects of hard water on the surfactants. For this reason, water softening is often unnecessary. Where softening is practised, it is often recommended to soften only the water sent to domestic hot water systems to prevent or delay inefficiencies and damage due to scale formation in water heaters. A common method for water softening involves the use ofion-exchange resins,which replace ions like Ca2+by twice the number of mono cations such assodiumorpotassiumions.

Washing soda (sodium carbonate,Na2CO3) is easily obtained and has long been used as a water softener for domestic laundry, in conjunction with the usual soap or detergent.

Water that has been treated by awater softeningmay be termedsoftened water.In these cases, the water may also contain elevated levels ofsodiumorpotassiumandbicarbonateorchlorideions.

Health considerations

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TheWorld Health Organizationsays that "there does not appear to be any convincing evidence that water hardness causes adverse health effects in humans".[2]In fact, theUnited States National Research Councilhas found that hard water serves as a dietary supplement for calcium and magnesium.[13]

Some studies have shown a weakinverse relationshipbetween water hardness andcardiovascular diseasein men, up to a level of 170 mg calcium carbonate per litre of water. The World Health Organization has reviewed the evidence and concluded the data was inadequate to recommend a level of hardness.[2]

Recommendations have been made for the minimum and maximum levels of calcium (40–80ppm) and magnesium (20–30 ppm) in drinking water, and a total hardness expressed as the sum of the calcium and magnesium concentrations of 2–4 mmol/L.[14]

Other studies have shown weak correlations between cardiovascular health and water hardness.[15][16][17]

The prevalence ofatopic dermatitis(eczema) in children may be increased by hard drinking water.[18][19]Living in areas with hard water may also play a part in the development of AD in early life. However, when AD is already established, usingwater softenersat home does not reduce the severity of the symptoms.[19]

Measurement

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Hardness can be quantified byinstrumental analysis.The total water hardness is the sum of themolar concentrationsof Ca2+and Mg2+,in mol/L or mmol/L units. Although water hardness usually measures only the total concentrations of calcium and magnesium (the two most prevalentdivalentmetal ions),iron,aluminium,andmanganeseare also present at elevated levels in some locations. The presence of iron characteristically confers a brownish (rust-like) colour to the calcification, instead of white (the colour of most of the other compounds).

Water hardness is often not expressed as a molar concentration, but rather in various units, such as degrees of general hardness (dGH), German degrees (°dH), parts per million (ppm, mg/L, or American degrees), grains per gallon (gpg), English degrees (°e, e, or°Clark), or French degrees (°fH, °f or °HF; lowercasefis used to prevent confusion with degreesFahrenheit). The table below shows conversion factors between the various units.

Hardness unit conversion.
1 mmol/L 1 ppm, mg/L 1 dGH, °dH 1 gpg 1 °e, °Clark 1 °fH
mmol/L 1 0.009991 0.1783 0.171 0.1424 0.09991
ppm, mg/L 100.1 1 17.85 17.12 14.25 10
dGH, °dH 5.608 0.05603 1 0.9591 0.7986 0.5603
gpg 5.847 0.05842 1.043 1 0.8327 0.5842
°e, °Clark 7.022 0.07016 1.252 1.201 1 0.7016
°fH 10.01 0.1 1.785 1.712 1.425 1

The various alternative units represent an equivalent mass of calcium oxide (CaO) or calcium carbonate (CaCO3) that, when dissolved in a unit volume of pure water, would result in the same total molar concentration of Mg2+and Ca2+.The different conversion factors arise from the fact that equivalent masses of calcium oxide and calcium carbonates differ and that different mass and volume units are used. The units are as follows:

  • Parts per million (ppm)is usually defined as 1 mg/L CaCO3(the definition used below).[20]It is equivalent tomg/Lwithout chemical compound specified, and toAmerican degree.
  • Grain per gallon(gpg)is defined as 1grain(64.8 mg) of calcium carbonate perU.S. gallon(3.79 litres), or 17.118 ppm.
  • ammol/Lis equivalent to 100.09 mg/L CaCO3or 40.08 mg/L Ca2+.
  • Adegree of General Hardness (dGHor 'German degree (°dH,deutsche Härte))' is defined as 10 mg/L CaO or 17.848 ppm.
  • AClark degree (°Clark)orEnglish degrees (°e or e)is defined as onegrain(64.8 mg) of CaCO3perImperial gallon(4.55 litres) of water, equivalent to 14.254 ppm.
  • AFrench degree (°fH or °f)is defined as 10 mg/L CaCO3,equivalent to 10 ppm.

Hard/soft classification

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As it is the precise mixture of minerals dissolved in the water, together with water'spHand temperature, that determine the behaviour of the hardness, a single-number scale does not adequately describe hardness. However, theUnited States Geological Surveyuses the following classification for hard and soft water:[5]

Classification mg-CaCO3/L (ppm) mmol/L dGH/°dH gpg
Soft 0–60 0–0.60 0–3.37 0–3.50
Moderately hard 61–120 0.61–1.20 3.38–6.74 3.56–7.01
Hard 121–180 1.21–1.80 6.75–10.11 7.06–10.51
Very hard ≥ 181 ≥ 1.81 ≥ 10.12 ≥ 10.57

Seawater is considered to be very hard due to various dissolved salts. Typically seawater's hardness is in the area of 6,570; ppm (6.57 grams per litre).[21]In contrast, freshwater has a hardness in the range of 15 to 375 ppm; generally around 600 mg/L.[22]

Indices

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Several indices are used to describe the behaviour of calcium carbonate in water, oil, or gas mixtures.[23]

Langelier saturation index (LSI)

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The Langelier saturation index[24](sometimes Langelier stability index) is a calculated number used to predict the calcium carbonate stability of water.[25]It indicates whether the water will precipitate, dissolve, or be in equilibrium with calcium carbonate. In 1936, Wilfred Langelier developed a method for predicting the pH at which water is saturated in calcium carbonate (called pHs).[26]The LSI is expressed as the difference between the actual system pH and the saturation pH:[27]

LSI = pH (measured) − pHs
  • For LSI > 0, water is supersaturated and tends to precipitate a scale layer of CaCO3.
  • For LSI = 0, water is saturated (in equilibrium) with CaCO3.A scale layer of CaCO3is neither precipitated nor dissolved.
  • For LSI < 0, water is under-saturated and tends to dissolve solid CaCO3.

If the actual pH of the water is below the calculated saturation pH, the LSI is negative and the water has a very limited scaling potential. If the actual pH exceeds pHs, the LSI is positive, and being supersaturated with CaCO3,the water tends to form scale. At increasing positive index values, the scaling potential increases.

In practice, water with an LSI between −0.5 and +0.5 will not display enhanced mineral dissolving or scale-forming properties. Water with an LSI below −0.5 tends to exhibit noticeably increased dissolving abilities while water with an LSI above +0.5 tends to exhibit noticeably increased scale-forming properties.

The LSI is temperature-sensitive. The LSI becomes more positive as the water temperature increases. This has particular implications in situations where well water is used. The temperature of the water when it first exits the well is often significantly lower than the temperature inside the building served by the well or at the laboratory where the LSI measurement is made. This increase in temperature can cause scaling, especially in cases such as water heaters. Conversely, systems that reduce water temperature will have less scaling.

Water analysis:
pH = 7.5
TDS= 320 mg/L
Calcium = 150 mg/L (or ppm) as CaCO3
Alkalinity = 34 mg/L (or ppm) as CaCO3
LSI formula:
LSI = pH − pHs
pHs= (9.3 + A + B) − (C + D) where:
A =⁠log10[TDS] − 1/10⁠= 0.15
B = −13.12 × log10(°C + 273) + 34.55 = 2.09 at 25 °C and 1.09 at 82 °C
C = log10[Ca2+as CaCO3] - 0.4 = 1.78
(Ca2+as CaCO3is also called calcium hardness, and is calculated as 2.5[Ca2+])
D = log10[alkalinity as CaCO3] = 1.53

Ryznar stability index (RSI)

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The Ryznar stability index (RSI)[24]: 525 uses a database of scale thickness measurements in municipal water systems to predict the effect of water chemistry.[25]: 72 [28]It was developed from empirical observations of corrosion rates and film formation in steel mains.

This index is defined as:[29]

RSI = 2 pHs– pH (measured)
  • For 6.5 < RSI < 7 water is considered to be approximately at saturation equilibrium with calcium carbonate
  • For RSI > 8 water is undersaturated and, therefore, would tend to dissolve any existing solid CaCO3
  • For RSI < 6.5 water tends to be scale form

Puckorius scaling index (PSI)

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The Puckorius scaling index (PSI) uses slightly different parameters to quantify the relationship between the saturation state of the water and the amount of limescale deposited.

Other indices

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Other indices include the Larson-Skold Index,[30]the Stiff-Davis Index,[31]and the Oddo-Tomson Index.[32]

Regional information

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The hardness of local water supplies depends on the source of water. Water in streams flowing over volcanic (igneous) rocks will be soft, while water from boreholes drilled into porous rock is normally very hard.

In Australia

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Analysis of water hardness in major Australian cities by theAustralian Water Associationshows a range from very soft (Melbourne) to hard (Adelaide). Total hardness levels of calcium carbonate in ppm are:

In Canada

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Prairieprovinces (mainlySaskatchewanandManitoba) contain high quantities of calcium and magnesium, often asdolomite,which are readily soluble in the groundwater that contains high concentrations of trappedcarbon dioxidefrom the lastglaciation.In these parts of Canada, the total hardness in ppm of calcium carbonate equivalent frequently exceeds 200 ppm, if groundwater is the only source of potable water. The west coast, by contrast, has unusually soft water, derived mainly from mountain lakes fed by glaciers and snowmelt.

Some typical values are:

In England and Wales

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Hardness water level of major cities in England and Wales
Area Primary source Level[51]
Manchester Lake District(Haweswater,Thirlmere)Pennines(Longdendale Chain) 1.750 °clark / 25 ppm[52]
Birmingham Elan Valley Reservoirs 3 °clark / 42.8 ppm[53]
Bristol Mendip Hills(Bristol Reservoirs) 16 °clark / 228.5 ppm[54]
Southampton Bewl Water 18.76 °clark / 268 ppm[55]
London (EC1A) Lee Valley Reservoir Chain 19.3 °clark / 275 ppm[56]
Wrexham (LL11) Hafren Dyfrdwy 4.77 °clark[57]

Information from the British Drinking Water Inspectorate[58]shows that drinking water inEnglandis generally considered to be 'very hard', with most areas of England, particularly east of a line between theSevernandTeesestuaries, exhibiting above 200 ppm for the calcium carbonate equivalent. Water in London, for example, is mostly obtained from theRiver ThamesandRiver Leaboth of which derive a significant proportion of their dry weather flow from springs in limestone and chalk aquifers.Wales,Devon,Cornwalland parts ofnorthwest Englandare softer water areas and range from 0 to 200 ppm.[59]In thebrewingindustry in England and Wales, water is often deliberately hardened withgypsumin the process ofBurtonisation.

Generally, water is mostly hard in urban areas of England where soft water sources are unavailable. Several cities built water supply sources in the 18th century as theIndustrial Revolutionand urban population burgeoned.Manchesterwas a notable such city in North West England and its wealthy corporation built several reservoirs atThirlmereandHaweswaterin theLake Districtto the north. There is no exposure tolimestoneorchalkin theirheadwatersand consequently the water in Manchester is rated as 'very soft'.[52]Similarly, tap water inBirminghamis also soft as it is sourced from theElan Valley Reservoirsin Wales, even though groundwater in the area is hard.

In Ireland

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The EPA has published a standards handbook for the interpretation of water quality in Ireland in which definitions of water hardness are given.[60] In this section, reference to original EU documentation is given, which sets out no limit for hardness. The handbook also gives no "Recommended or Mandatory Limit Values" for hardness. The handbook does indicate that above the midpoint of the ranges defined as "Moderately Hard", effects are seen increasingly: "The chief disadvantages of hard waters are that they neutralise the lathering power of soap[...] and, more important, that they can cause blockage of pipes and severely reduced boiler efficiency because of scale formation. These effects will increase as the hardness rises to and beyond 200 mg/LCaCO
3."

In the United States

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A collection of data from the United States found that about half the water stations tested had hardness over 120 mg per litre of calcium carbonate equivalent, placing them in the categories "hard" or "very hard".[5]The other half were classified as soft or moderately hard. More than 85% of American homes have hard water.[citation needed]The softest waters occur in parts of theNew England,South Atlantic-Gulf,Pacific Northwest,andHawaiiregions. Moderately hard waters are common in many of the rivers of theTennessee,Great Lakes,andAlaskaregions. Hard and very hard waters are found in some of the streams in most of the regions throughout the country. The hardest waters (greater than 1,000 ppm) are in streams in Texas, New Mexico, Kansas, Arizona, Utah, parts of Colorado, southern Nevada, and southern California.[61][62]

See also

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References

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