Hard water

Hard water is water that has high mineral content (in contrast with "soft water"). Hard water is formed when water percolates through deposits of limestone, chalk or gypsum[1] which are largely made up of calcium and magnesium carbonates, bicarbonates and sulfates.

Hard drinking water may have moderate health benefits, but can pose critical problems in industrial settings, where water hardness is monitored to avoid costly breakdowns in boilers, cooling towers, and other equipment that handles water. In domestic settings, hard water is often indicated by a lack of foam formation when soap is agitated in water, and by the formation of limescale in kettles and water heaters.[2] Wherever water hardness is a concern, water softening is commonly used to reduce hard water's adverse effects.

Hard Water Calcification
A bathtub faucet with built-up calcification from hard water in Southern Arizona.

Sources of hardness

Water's hardness is determined by the concentration of multivalent cations in the water. Multivalent cations are positively charged metal complexes with a charge greater than 1+. Usually, the cations have the charge of 2+. Common cations found in hard water include Ca2+ and Mg2+. These ions enter a water supply by leaching from minerals within an aquifer. Common calcium-containing minerals are calcite and gypsum. A common magnesium mineral is dolomite (which also contains calcium). Rainwater and distilled water are soft, because they contain few ions.[3]

The following equilibrium reaction describes the dissolving and formation of calcium carbonate and calcium bicarbonate (on the right):

CaCO3 (s) + CO2 (aq) + H2O (l) ⇋ Ca2+ (aq) + 2HCO3 (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 atmosphere, sometimes forming stalactites and stalagmites.

Calcium and magnesium ions can sometimes be removed by water softeners.[4]

Temporary hardness

Temporary hardness is a type of water hardness caused by the presence of dissolved bicarbonate minerals (calcium bicarbonate and magnesium bicarbonate). When dissolved, these minerals yield calcium and magnesium cations (Ca2+, Mg2+) and carbonate and bicarbonate anions (CO32−, HCO3). The presence of the metal cations makes the water hard. However, unlike the permanent hardness caused by sulfate and chloride compounds, this "temporary" hardness can be reduced either by boiling the water, or by the addition of lime (calcium hydroxide) through the process of lime softening.[5] Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.

Permanent hardness

Permanent hardness (mineral content) are generally difficult to remove by boiling.[6] If this occurs, it is usually caused by the presence of calcium sulfate/calcium chloride and/or magnesium sulfate/magnesium chloride in the water, which do not precipitate out as the temperature increases. Ions causing permanent hardness of water can be removed using a water softener, or ion exchange column.

Permanent Hardness = Permanent Calcium Hardness + Permanent Magnesium Hardness.


With hard water, soap solutions form a white precipitate (soap scum) instead of producing lather, because the 2+ ions destroy the surfactant properties of the soap by forming a solid precipitate (the soap scum). A major component of such scum is calcium stearate, which arises from sodium stearate, the main component of soap:

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. Synthetic detergents do not form such scums.

A portion of the ancient Roman Eifel aqueduct in Germany. After being in service for about 180 years, the aqueduct had mineral deposits of up to 20 cm thick along the walls.

Hard water also forms deposits that clog plumbing. These deposits, called "scale", are composed mainly of calcium carbonate (CaCO3), magnesium hydroxide (Mg(OH)2), and calcium sulfate (CaSO4).[3] Calcium and magnesium carbonates tend to be deposited as off-white solids on the inside surfaces of pipes and heat 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.[7] 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 failure of the boiler.[8] The damage caused by calcium carbonate deposits varies on the crystalline form, for example, calcite or aragonite.[9]

The presence of ions in an electrolyte, in this case, hard water, can also lead to galvanic corrosion, in which one metal will preferentially corrode when 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 its corrosivity per se. Similarly, where lead plumbing is in use, softened water does not substantially increase plumbo-solvency.[10]

In swimming pools, hard water is manifested by a turbid, 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 the insoluble carbonates, giving rise to the turbidity. This often results from the pH being 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 of hydrochloric acid, the optimum value being in the range of 7.2 to 7.6.


It is often 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 so as to prevent or delay inefficiencies and damage due to scale formation in water heaters. A common method for water softening involves the use of ion exchange resins, which replace ions like Ca2+ by twice the number of monocations such as sodium or potassium ions.

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.

Health considerations

The World Health Organization says that "there does not appear to be any convincing evidence that water hardness causes adverse health effects in humans".[2] In fact, the United States National Research Council has found that hard water actually serves as a dietary supplement for calcium and magnesium.[11]

Some studies have shown a weak inverse relationship between water hardness and cardiovascular disease in 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 allow for a recommendation for a level of hardness.[2]

Recommendations have been made for the maximum and minimum levels of calcium (40–80 ppm) 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.[12]

Other studies have shown weak correlations between cardiovascular health and water hardness.[13][14][15]

Some studies correlate domestic hard water usage with increased eczema in children.[16][17][18][19]

The Softened-Water Eczema Trial (SWET), a multicenter randomized controlled trial of ion-exchange softeners for treating childhood eczema, was undertaken in 2008. However, no meaningful difference in symptom relief was found between children with access to a home water softener and those without.[20]


Hardness can be quantified by instrumental analysis. The total water hardness is the sum of the molar concentrations of 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 prevalent divalent metal ions), iron, aluminium, and manganese can also be 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 color 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; lowercase f is used to prevent confusion with degrees Fahrenheit). 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).[21] It is equivalent to mg/L without chemical compound specified, and to American degree.
  • Grains per Gallon (gpg) is defined as 1 grain (64.8 mg) of calcium carbonate per U.S. gallon (3.79 litres), or 17.118 ppm.
  • a mmol/L is equivalent to 100.09 mg/L CaCO3 or 40.08 mg/L Ca2+.
  • A degree of General Hardness (dGH or 'German degree (°dH, deutsche Härte))' is defined as 10 mg/L CaO or 17.848 ppm.
  • A Clark degree (°Clark) or English degrees (°e or e) is defined as one grain (64.8 mg) of CaCO3 per Imperial gallon (4.55 litres) of water, equivalent to 14.254 ppm.
  • A French degree (°fH or °f) is defined as 10 mg/L CaCO3, equivalent to 10 ppm.

Hard/soft classification

Because it is the precise mixture of minerals dissolved in the water, together with the water's pH and temperature, that determine the behavior of the hardness, a single-number scale does not adequately describe hardness. However, the United States Geological Survey uses the following classification into hard and soft water,[22]

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

Seawater is considered to be very hard due to various dissolved salts. Typically seawater's hardness is in the range of 6630 ppm. In contrast, freshwater has hardness in the range of 15 - 375 ppm.[23]


Several indices are used to describe the behaviour of calcium carbonate in water, oil, or gas mixtures.[24]

Langelier saturation index (LSI)

The Langelier saturation index[25] (sometimes Langelier stability index) is a calculated number used to predict the calcium carbonate stability of water.[26] 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).[27] The LSI is expressed as the difference between the actual system pH and the saturation pH:[28]

LSI = pH (measured) − pHs
  • For LSI > 0, water is super saturated and tends to precipitate a scale layer of CaCO3.
  • For LSI = 0, water is saturated (in equilibrium) with CaCO3. A scale layer of CaCO3 is 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 has a tendency 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 hot 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 x 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 CaCO3 is also called Calcium Hardness and is calculated as=2.5(Ca2+))
D = Log10[alkalinity as CaCO3] = 1.53

Ryznar Stability Index (RSI)

The Ryznar stability index (RSI)[25]:525 uses a database of scale thickness measurements in municipal water systems to predict the effect of water chemistry.[26]:72[29]

Ryznar saturation index (RSI) was developed from empirical observations of corrosion rates and film formation in steel mains. It is defined as:[30]

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 under saturated and, therefore, would tend to dissolve any existing solid CaCO3
  • For RSI < 6,5 water tends to be scale forming

Puckorius Scaling Index (PSI)

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

Other indices include the Larson-Skold Index,[31] the Stiff-Davis Index,[32] and the Oddo-Tomson Index.[33]

Regional information

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

Analysis of water hardness in major Australian cities by the Australian Water Association shows a range from very soft (Melbourne) to hard (Adelaide). Total Hardness levels of calcium carbonate in ppm are: Canberra: 40;[34] Melbourne: 10–26;[35] Sydney: 39.4–60.1;[36] Perth: 29–226;[37] Brisbane: 100;[38] Adelaide: 134–148;[39] Hobart: 5.8–34.4;[40] Darwin: 31.[41]

In Canada

Prairie provinces (mainly Saskatchewan and Manitoba) contain high quantities of calcium and magnesium, often as dolomite, which are readily soluble in the groundwater that contains high concentrations of trapped carbon dioxide from the last glaciation. In these parts of Canada, the total hardness in ppm of calcium carbonate equivalent frequently exceed 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: Montreal 116 ppm,[42] Calgary 165 ppm, Regina 496 ppm,[43] Saskatoon 160-180 ppm,[44] Winnipeg 77 ppm,[45] Toronto 121 ppm,[46] Vancouver < 3 ppm,[47] Charlottetown, PEI 140–150 ppm,[48] Waterloo Region 400 ppm, Guelph 460 ppm,[49] Saint John (West) 160-200 ppm.[50]

In England and Wales

Hardness water level of major cities in the UK
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]

Information from the British Drinking Water Inspectorate[57] shows that drinking water in England is generally considered to be 'very hard', with most areas of England, particularly east of a line between the Severn and Tees estuaries, exhibiting above 200 ppm for the calcium carbonate equivalent. Water in London, for example, is mostly obtained from the River Thames and River Lea both of which derive significant proportion of their dry weather flow from springs in limestone and chalk aquifers. Wales, Devon, Cornwall and parts of North-West England are softer water areas, and range from 0 to 200 ppm.[58] In the brewing industry in England and Wales, water is often deliberately hardened with gypsum in the process of Burtonisation.

Generally water is mostly hard in urban areas of England where soft water sources are unavailable. A number of cities built water supply sources in the 18th century as the industrial revolution and urban population burgeoned. Manchester was a notable such city in North West England and its wealthy corporation built a number of reservoirs at Thirlmere and Haweswater in the Lake District to the north. There is no exposure to limestone or chalk in their headwaters and consequently the water in Manchester is rated as 'very soft'.[52] Similarly, tap water in Birmingham is also soft as it is sourced from the Elan Valley Reservoirs in Wales.

In Ireland

The EPA has published a standards handbook for the interpretation of water quality in Ireland in which definitions of water hardness are given.[59] In this section, reference to original EU documentation is given, which sets out no limit for hardness. In turn, the handbook also gives no "Recommended or Mandatory Limit Values" for Hardness. The handbooks 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/l CaCO3."

In the United States

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".[22] The other half were classified as soft or moderately hard. More than 85% of American homes have hard water.[60] The softest waters occur in parts of the New England, South Atlantic-Gulf, Pacific Northwest, and Hawaii regions. Moderately hard waters are common in many of the rivers of the Tennessee, Great Lakes, and Alaska regions. 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


  1. ^ "Hard water". National Groundwater Association. Retrieved 28 June 2019.
  2. ^ a b c World Health Organization Hardness in Drinking-Water, 2003
  3. ^ a b Hermann Weingärtner, "Water" in Ullmann's Encyclopedia of Industrial Chemistry, 2006[December], Wiley–VCH, Weinheim. doi:10.1002/14356007.a28_001
  4. ^ Christian Nitsch, Hans-Joachim Heitland, Horst Marsen, Hans-Joachim Schlüussler, "Cleansing Agents" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley–VCH, Weinheim. doi:10.1002/14356007.a07_137
  5. ^ "Lime Softening". Retrieved 4 November 2011.
  6. ^ Sengupta, Pallav (August 2013). "Potential Health Impacts of Hard Water". International Journal of Preventive Medicine. 4 (8): 866–875. ISSN 2008-7802. PMC 3775162. PMID 24049611 – via NCBI.
  7. ^ Wisconisin DNR - Carbonate chemistry
  8. ^ Stephen Lower (July 2007). "Hard water and water softening". Retrieved 2007-10-08.
  9. ^ PP Coetzee (1998). "Scale reduction and scale modification effects induced by Zn" (PDF). Retrieved 2010-03-29.
  10. ^ Sorg, Thomas J.; Schock, Michael R.; Lytle, Darren A. (August 1999). "Ion Exchange Softening: Effects on Metal Concentrations". Journal AWWA. 91 (8): 85–97. doi:10.1002/j.1551-8833.1999.tb08685.x. ISSN 1551-8833
  11. ^ "Drinking Water Hardwater Hardness Calcium Magnesium Scale Stained Laundry". Water-research.net. Retrieved 2013-01-26.
  12. ^ František Kožíšek Health significance of drinking water calcium and magnesium, February 2003
  13. ^ Pocock SJ, Shaper AG, Packham RF (April 1981). "Studies of water quality and cardiovascular disease in the United Kingdom". Sci. Total Environ. 18: 25–34. Bibcode:1981ScTEn..18...25P. doi:10.1016/S0048-9697(81)80047-2. PMID 7233165.
  14. ^ Marque S, Jacqmin-Gadda H, Dartigues JF, Commenges D (2003). "Cardiovascular mortality and calcium and magnesium in drinking water: an ecological study in elderly people" (PDF). Eur. J. Epidemiol. 18 (4): 305–9. doi:10.1023/A:1023618728056. PMID 12803370.
  15. ^ Rubenowitz E, Axelsson G, Rylander R (January 1999). "Magnesium and calcium in drinking water and death from acute myocardial infarction in women". Epidemiology. 10 (1): 31–6. doi:10.1097/00001648-199901000-00007. PMID 9888277.
  16. ^ McNally NJ, Williams HC, Phillips DR, Smallman-Raynor M, Lewis S, Venn A, Britton J (1998). "Atopic eczema and domestic water hardness". The Lancet. 352 (9127): 527–531. doi:10.1016/S0140-6736(98)01402-0. PMID 9716057.
  17. ^ Miyake Y, Yokoyama T, Yura A, Iki M, Shimizu T (Jan 2004). "Ecological association of water hardness with prevalence of childhood atopic dermatitis in a Japanese urban area". Environ. Res. 94 (1): 33–7. Bibcode:2004ER.....94...33M. doi:10.1016/S0013-9351(03)00068-9. PMID 14643284.
  18. ^ Arnedo-Pena A, Bellido-Blasco J, Puig-Barbera J, Artero-Civera A, Campos-Cruañes JB, Pac-Sa MR, Villamarín-Vázquez JL, Felis-Dauder C (2007). "Domestic water hardness and prevalence of atopic eczema in Castellon (Spain) school children". Salud Pública de México. 49 (4): 295–301. doi:10.1590/S0036-36342007000400009. PMID 17710278.
  19. ^ Perkin MR, Craven J, Logan K, Strachan D, Marrs T, Radulovic S, Campbell LE, MacCallum SF, McLean WH, Lack G, Flohr C, Enquiring About Tolerance Study Team (2016). "Association between domestic water hardness, chlorine, and atopic dermatitis risk in early life: A population-based cross-sectional study". J Allergy Clin Immunol. 138 (2): 509–516. doi:10.1016/j.jaci.2016.03.031. PMID 27241890.
  20. ^ A multicentre randomized controlled trial of ion-exchange water softeners for the treatment of eczema in children: protocol for the Softened Water Eczema Trial (SWET) (ISRCTN: 71423189) http://www.swet-trial.co.uk/
  21. ^ "Water Hardness". thekrib.com.
  22. ^ a b USGS - U.S. Geological Survey Office of Water Quality. "USGS Water-Quality Information: Water Hardness and Alkalinity". usgs.gov.
  23. ^ Total water hardness
  24. ^ Corrosion by water Archived 2007-10-20 at the Wayback Machine
  25. ^ a b McTigue, Nancy E.; Symons, James M., eds. (2011). The Water Dictionary: A Comprehensive Reference of Water Terminology. American Water Works Association. pp. 333–. ISBN 978-1-61300-101-1.
  26. ^ a b Reid, Robert N. (2003). Water Quality Systems: Guide For Facility Managers. CRC Press. pp. 66–. ISBN 978-0-8247-4010-8.
  27. ^ Langelier, W. F. (October 1936). "The Analytical Control of Anti-Corrosion Water Treatment". Journal of the American Water Works Association. 28 (10): 1500–1521. doi:10.1002/j.1551-8833.1936.tb13785.x. JSTOR 41226418.
  28. ^ Aquaprox, ed. (2009). Treatment of cooling water. Springer. pp. 104–. ISBN 978-3-642-01985-2.
  29. ^ Emerson, A. G. D. (2003). Quantitative Forecasting of Problems in Industrial Water Systems. World Scientific. pp. 7–. ISBN 978-981-238-184-2.
  30. ^ Ryznar, John W.; Langelier, W. F. (April 1944). "A New Index for Determining Amount of Calcium Carbonate Scale Formed by a Water". Journal of the American Water Works Association. 36 (4): 472–486. doi:10.1002/j.1551-8833.1944.tb20016.x. JSTOR 23345279.
  31. ^ T.E., Larson and R. V. Skold, Laboratory Studies Relating Mineral Quality of Water to Corrosion of Steel and Cast Iron, 1958 Illinois State Water Survey, Champaign, IL pp. [43] — 46: ill. ISWS C-71
  32. ^ Stiff, Jr., H.A., Davis, L.E., A Method For Predicting The Tendency of Oil Field Water to Deposit Calcium Carbonate, Pet. Trans. AIME 195;213 (1952).
  33. ^ Oddo, J.E., Tomson, M.B., Scale Control, Prediction and Treatment Or How Companies Evaluate A Scaling Problem and What They Do Wrong, CORROSION/92, Paper No. 34, (Houston, TX:NACE INTERNATIONAL 1992). KK
  34. ^ "Dishwasher and water hardness - Canberra water quality - About Us". actewagl.com.au. Archived from the original on 2012-03-26.
  35. ^ "Page not found" (PDF). melbournewater.com.au.
  36. ^ Sydney Typical Drinking Water Analysis
  37. ^ "Water Corporation of WA - 404" (PDF). watercorporation.com.au. Archived from the original (PDF) on 2007-09-04.
  38. ^ Brisbane Drinking Water
  39. ^ Adelaide Water Quality
  40. ^ "Hobart City Council, Tasmania Australia". hobartcity.com.au.
  41. ^ Darwin Water Quality
  42. ^ "Ville de Montréal - L'eau de Montréal". .ville.montreal.qc.ca. 2013-01-22. Retrieved 2013-01-26.
  43. ^ Canadian Water Quality Association. "file:///D|/_Current_Sites/cwqa2006/2011/_faq/water_hardness_canada.inc.html [1/29/2013 3:09:35 AM] Water Hardness/Total Households Canadian Cities" (PDF). Archived from the original (PDF) on 4 October 2013. Retrieved 4 October 2013.
  44. ^ "Frequently Asked Questions". Saskatoon.ca. Retrieved 2013-01-26.
  45. ^ 2006 Winnipeg drinking water quality test results
  46. ^ "Water - Services - Living In Toronto - City of Toronto". toronto.ca.
  47. ^ GVRD Wash Smart – Water Facts
  48. ^ "CITY OF CHARLOTTETOWN WATER & SEWER UTILITY Water Report 2006" (PDF). Aquasafecanada.com. Retrieved 2013-01-26.
  49. ^ "REGION OF WATERLOO Residential Water Softener Performance Study Testing Report #1 April, 2011" (PDF). Regionofwaterloo.ca. Retrieved 2013-01-26.
  50. ^ "Public Information Notice: West Side Water Supply | Saint John". www.saintjohn.ca. Retrieved 2017-10-10.
  51. ^ "Table 2 Drinking Water Hardness". United Utilities. Retrieved 2012-03-03.
  52. ^ a b "Drinking water quality". United Utilities. Retrieved 2012-03-03.
  53. ^ "Severn Trent Water — B1 1DB". Severn Trent Water. Retrieved 2012-03-03.
  54. ^ "Bristol water hardness level". Bristol Water. Archived from the original on 2011-08-01. Retrieved 2012-03-03.
  55. ^ "Southern Water — SO14 area". Southern Water. Retrieved 2012-03-03.
  56. ^ "EC1A 7BE — Water quality in your area". Thames Water. Retrieved 2012-03-03.
  57. ^ dwi.gov.uk
  58. ^ anglianwater.co.uk
  59. ^ Section 36 "Hardness" https://www.epa.ie/pubs/advice/water/quality/Water_Quality.pdf
  60. ^ Wilson, Amber; Parrott, Kathleen; Ross, Blake (June 1999). "Household Water Quality – Water Hardness". Retrieved 2009-04-26.
  61. ^ Briggs, J.C., and Ficke, J.F.; Quality of Rivers of the United States, 1975 Water Year -- Based on the National Stream Quality Accounting Network (NASQAN): U.S. Geological Survey Open-File Report 78-200, 436 p. (1977)
  62. ^ "Got Hard Water? Here's What You Need To Know About It". Modern Home Pulse. Retrieved 2018-09-22.

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Atopic dermatitis

Atopic dermatitis (AD), also known as atopic eczema, is a type of inflammation of the skin (dermatitis). It results in itchy, red, swollen, and cracked skin. Clear fluid may come from the affected areas, which often thickens over time. While the condition may occur at any age, it typically starts in childhood with changing severity over the years. In children under one year of age much of the body may be affected. As children get older, the back of the knees and front of the elbows are the most common areas affected. In adults the hands and feet are the most commonly affected areas. Scratching worsens symptoms and affected people have an increased risk of skin infections. Many people with atopic dermatitis develop hay fever or asthma.The cause is unknown but believed to involve genetics, immune system dysfunction, environmental exposures, and difficulties with the permeability of the skin. If one identical twin is affected, there is an 85% chance the other also has the condition. Those who live in cities and dry climates are more commonly affected. Exposure to certain chemicals or frequent hand washing makes symptoms worse. While emotional stress may make the symptoms worse, it is not a cause. The disorder is not contagious. The diagnosis is typically based on the signs and symptoms. Other diseases that must be excluded before making a diagnosis include contact dermatitis, psoriasis, and seborrheic dermatitis.Treatment involves avoiding things that make the condition worse, daily bathing with application of a moisturising cream afterwards, applying steroid creams when flares occur, and medications to help with itchiness. Things that commonly make it worse include wool clothing, soaps, perfumes, chlorine, dust, and cigarette smoke. Phototherapy may be useful in some people. Steroid pills or creams based on calcineurin inhibitors may occasionally be used if other measures are not effective. Antibiotics (either by mouth or topically) may be needed if a bacterial infection develops. Dietary changes are only needed if food allergies are suspected.Atopic dermatitis affects about 20% of people at some point in their lives. It is more common in younger children. Males and females are equally affected. Many people outgrow the condition. Atopic dermatitis is sometimes called eczema, a term that also refers to a larger group of skin conditions. Other names include "infantile eczema", "flexural eczema", "prurigo Besnier", "allergic eczema", and "neurodermatitis".

Blechnum wattsii

Blechnum wattsii or the hard water fern is a common terrestrial fern growing in rainforest and open forest. Often seen near creeks in much of south eastern Australia, including Victoria, Tasmania (and King Island), South Australia, New South Wales and Queensland. The Blechnum wattsii was named for Reverend William Walter Watts (1856-1920). Reverend Watts was considered an authority on mosses and ferns and has more than 30 species named for him. Common names by which B. wattsii may be called are hard water fern - from its stiff leathery fronds, leech fern - as forest workers often encounter leaches while working in clusters of these ferns, hard hill fern - from the fern's habit and habitat, and red cabbage fern - from the bronze-pink colour of the young fronds resembling cooked red cabbage.

Calcium carbonate

Calcium carbonate is a chemical compound with the formula CaCO3. It is a common substance found in rocks as the minerals calcite and aragonite (most notably as limestone, which is a type of sedimentary rock consisting mainly of calcite) and is the main component of pearls and the shells of marine organisms, snails, and eggs. Calcium carbonate is the active ingredient in agricultural lime and is created when calcium ions in hard water react with carbonate ions to create limescale. It is medicinally used as a calcium supplement or as an antacid, but excessive consumption can be hazardous.

Calcium hypochlorite

Calcium hypochlorite is an inorganic compound with formula Ca(ClO)2. It is the main active ingredient of commercial products called bleaching powder, chlorine powder, or chlorinated lime, used for water treatment and as a bleaching agents. This compound is relatively stable and has greater available chlorine than sodium hypochlorite (liquid bleach). It is a white solid, although commercial samples appear yellow. It strongly smells of chlorine, owing to its slow decomposition in moist air. It is not highly soluble in hard water, and is more preferably used in soft to medium-hard water. It has two forms: dry (anhydrous); and hydrated (hydrous).

Chara (alga)

Chara is a genus of charophyte green algae in the family Characeae. They are multicellular and superficially resemble land plants because of stem-like and leaf-like structures. They are found in fresh water, particularly in limestone areas throughout the northern temperate zone, where they grow submerged, attached to the muddy bottom. They prefer less oxygenated and hard water and are not found in waters where mosquito larvae are present. They are covered with calcium carbonate deposits and are commonly known as stoneworts. Cyanobacteria have been found growing as epiphytes on the surfaces of Chara, where they may be involved in fixing nitrogen, which is important to plant nutrition.


A detergent is a surfactant or a mixture of surfactants with cleaning properties in dilute solutions. These substances are usually alkylbenzenesulfonates, a family of compounds that are similar to soap but are more soluble in hard water, because the polar sulfonate (of detergents) is less likely than the polar carboxylate (of soap) to bind to calcium and other ions found in hard water.

In most household contexts, the term detergent by itself refers specifically to laundry detergent or dish detergent, as opposed to hand soap or other types of cleaning agents. Detergents are commonly available as powders or concentrated solutions. Detergents, like soaps, work because they are amphiphilic: partly hydrophilic (polar) and partly hydrophobic (non-polar). Their dual nature facilitates the mixture of hydrophobic compounds (like oil and grease) with water. Because air is not hydrophilic, detergents are also foaming agents to varying degrees.


A humidifier is a device, primarily an electrical appliance that increases humidity (moisture) in a single room or an entire building. In the home, point-of-use humidifiers are commonly used to humidify a single room, while whole-house or furnace humidifiers, which connect to a home's HVAC system, provide humidity to the entire house. Medical ventilators often include humidifiers for increased patient comfort. Large humidifiers are used in commercial, institutional, or industrial contexts, often as part of a larger HVAC system.


Kieserite is the magnesium sulfate mineral (MgSO4·H2O) and is named after Dietrich Georg von Kieser (Jena, Germany 1862). It has a vitreous luster and it is colorless, grayish-white or yellowish. Its hardness is 3.5 and crystallizes in the monoclinic crystal system. Gunningite is the zinc member of the kieserite group of minerals.

Lake Bolluk

Lake Bolluk is a lake in Turkey.

The lake is in Cihanbeyli ilçe (district) of Konya Province at 38°32′25″N 32°56′34″E. It is situated to the east of the highway D.715, which connects Ankara to Silifke and to the west of Lake Tuz. The area of the lake is 11.5 square kilometres (4.4 sq mi). Its elevation with respect to sea level is 940 metres (3,080 ft). The hard water of the lake contains sodium. Rrecently, there are two threats to lake; the underground water level falls as a result of excessive irrigation and the creeks, which feed the lake, are polluted. World Water Forum Turkey conducts a project to protect the lake.

Lake Çöl

Lake Çöl (Turkish: Çöl Gölü, literally "Desert lake") is a hard water lake in Turkey.

Laundry detergent

Laundry detergent, or washing powder, is a type of detergent (cleaning agent) that is added for cleaning laundry. While detergent is still sold in powdered form, liquid detergents have been taking major market shares in many countries since their introduction in the 1950s.

Laundry detergent pods have also been sold in the United States since 2012 when they were introduced by Procter & Gamble as Tide Pods. Earlier instances of laundry detergent pods include Salvo tablets sold in the 1960s and 1970s.


Limescale is a hard, off-white, chalky deposit often found in kettles and hot water boilers and on the inside of hot water pipework. It is also often found as a similar deposit on the inner surfaces of old pipes and other surfaces where "hard water" has evaporated.

In addition to being unsightly and hard to clean, limescale can seriously damage or impair the operation of various plumbing and heating components. Descaling agents are commonly used to remove limescale. Prevention of fouling by scale build-up relies on the technologies of water softening.

Magnetic water treatment

Magnetic water treatment (also known as anti-scale magnetic treatment or AMT) is a method of supposedly reducing the effects of hard water by passing it through a magnetic field as a non-chemical alternative to water softening. Magnetic water treatment is regarded as unproven and unscientific.

There is a lack of peer-reviewed laboratory data, mechanistic explanations, and documented field studies to support its effectiveness. Erroneous conclusions about their efficacy are based on applications with uncontrolled variables. There are, however, some studies which have claimed significant effects and proposed possible mechanisms for the observed decrease in water scale.

Nevada wine

Nevada wine refers to wine made from grapes grown in the U.S. state of Nevada, where wine has been produced since 1990. There are currently no designated American Viticultural Areas in Nevada.

Nevada has five commercial wineries: Basin and Range Cellars in Reno(opening in June 2018), Nevada Sunset Winery also in Reno,Churchill Vineyards in Fallon, Pahrump Valley Winery in Pahrump and Sanders Family Winery (also located in Pahrump).

Locally high boron content of the soil, soil salinity, and hard water provide a few challenges to growing grapes, especially Vitis vinefera. Some environmental elements common to Nevada however are favorable for viticulture; these include the ample sunshine, low humidity (which decreases the risk for rot and mildew and thus the need for fungicides) University of Nevada, Reno professor Dr. Grant Cramer is currently studying the best varietals and techniques at UNR's Valley Road Vineyard. Wine grapes have been grown with success in both the northern and southern part of the state since 1991, and all five wineries have produced wines made from grapes grown in Nevada.

Until recently, state law restricted commercial wineries so that they were illegal in counties with more than 100,000 people (Washoe and Clark Counties). This was due to lobbying by the liquor distributors in the State for fight to control the liquor supply. With the passage of Assembly Bill 4 (AB4)in November 2015 [1] this law was changed. Since then three wineries have filed for licenses and began winemaking in Reno (Basin and Range Cellars, Nevada Sunset Winery and Great Basin Winery. The laws for winemaking in Nevada are still considered to be restrictive when compared to neighboring states that have successful wine industries.

Additionally, instructional wine-making facilities (such as the Valley Road Vineyard) may operate in any county but must meet special license requirements and are restricted to selling or distributing no more than 60 gallons of wine in any 12-month period.

Siren (DC Comics)

Siren is the name of two fictional supervillains, both appearing in books published by DC Comics.

Soap scum

Soap scum or lime soap is the white solid composed of calcium stearate, magnesium stearate, and similar alkali metal derivatives of fatty acids. These materials result from the addition of soap and other anionic surfactants to hard water. Hard water contains calcium and magnesium ions, which react with the surfactant anion to give these metallic or lime soaps.

2 C17H35COO−Na+ + Ca2+ → (C17H35COO)2Ca + 2 Na+In this reaction, the sodium cation in soap is replaced by calcium to form calcium stearate.

Lime soaps build deposits on fibres, washing machines, and sinks. Synthetic surfactants are less susceptible to the effects of hard water. Most detergents contain builders that prevent the formation of lime soaps.

Soap scum on vinyl shower curtains may have a rich microbial biofilm, containing potentially pathogenic bacteria.

Soft water

Soft water is surface water that contains low concentrations of ions and in particular is low in ions of calcium and magnesium. Soft water naturally occurs where rainfall and the drainage basin of rivers are formed of hard, impervious and calcium-poor rocks. Examples in the UK (United Kingdom) include Snowdonia in Wales and the Western Highlands in Scotland.

The term may also be used to describe water that has been produced by a water softening process although such water is more correctly termed softened water. In these cases the water may also contain elevated levels of sodium and bicarbonate ions.

Because soft water has few calcium ions, there is no inhibition of the lathering action of soaps and no soap scum is formed in normal washing. Similarly, soft water produces no calcium deposits in water heating systems. Water that isn't soft is referred to as hard water.

In the UK, water is regarded as soft if the hardness is less than 50 mg/l of calcium carbonate. Water containing more than 50 mg/l of calcium carbonate is termed hard water. In the United States soft water is classified as having less than 60 mg/l of calcium carbonate.In the USA, due to ancient sea beds with high limestone (calcium carbonate) concentrations, as much as 85% of water is hard, and many need water softening treatment. Waters in Eastern and Southern Florida, North-western Texas, and the entirety of South Dakota are considered extremely hard; the metropolitan cities with hardest water include Indianapolis, Las Vegas, Minneapolis, Phoenix, San Antonio and Tampa. The only states with soft water are Mississippi and Maine.

Total dissolved solids

Total dissolved solids (TDS) is a measure of the dissolved combined content of all inorganic and organic substances present in a liquid in molecular, ionized, or micro-granular (colloidal sol) suspended form. Generally, the operational definition is that the solids must be small enough to survive filtration through a filter with 2-micrometer (nominal size, or smaller) pores. Total dissolved solids are normally discussed only for freshwater systems, as salinity includes some of the ions constituting the definition of TDS. The principal application of TDS is in the study of water quality for streams, rivers, and lakes. Although TDS is not generally considered a primary pollutant (e.g. it is not deemed to be associated with health effects), it is used as an indication of aesthetic characteristics of drinking water and as an aggregate indicator of the presence of a broad array of chemical contaminants.

Primary sources for TDS in receiving waters are agricultural runoff and residential (urban) runoff, clay-rich mountain waters, leaching of soil contamination, and point source water pollution discharge from industrial or sewage treatment plants. The most common chemical constituents are calcium, phosphates, nitrates, sodium, potassium, and chloride, which are found in nutrient runoff, general stormwater runoff and runoff from snowy climates where road de-icing salts are applied. The chemicals may be cations, anions, molecules or agglomerations on the order of one thousand or fewer molecules, so long as a soluble micro-granule is formed. More exotic and harmful elements of TDS are pesticides arising from surface runoff. Certain naturally occurring total dissolved solids arise from the weathering and dissolution of rocks and soils. The United States has established a secondary water quality standard of 500 mg/l to provide for palatability of drinking water.

Total dissolved solids are differentiated from total suspended solids (TSS), in that the latter cannot pass through a sieve of 2 micrometers and yet are indefinitely suspended in solution. The term settleable solids refers to material of any size that will not remain suspended or dissolved in a holding tank not subject to motion, and excludes both TDS and TSS. Settleable solids may include larger particulate matter or insoluble molecules.

Water softening

Water softening is the removal of calcium, magnesium, and certain other metal cations in hard water. The resulting soft water requires less soap for the same cleaning effort, as soap is not wasted bonding with calcium ions. Soft water also extends the lifetime of plumbing by reducing or eliminating scale build-up in pipes and fittings. Water softening is usually achieved using lime softening or ion-exchange resins but is increasingly being accomplished using nanofiltration or reverse osmosis membranes.


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