Hard water

Hard water is water that has a 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.

Not to be confused with Heavy water or Ice.

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 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.

Origins[edit]

Natural rainwater, snow and other forms of precipitation typically have low concentrations of divalent cations such as calcium and magnesium. They may have small concentrations of ions such as sodium, chloride and sulfate derived 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 termed soft water.^[3] Examples include Snowdonia in 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[edit]

Permanent hardness[edit]

The permanent hardness of water is determined by the water's concentration of cations with charges greater than or equal to 2+. Usually, the cations have a charge of 2+, i.e., they are divalent. Common cations found in hard water include Ca²⁺ and Mg²⁺, which frequently enter water supplies by leaching from minerals within aquifers. 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 of these ions.^[3]

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

Parts per million (ppm) is usually defined as 1 mg/L CaCO₃ (the definition used below). It is equivalent to mg/L without chemical compound specified, and to American degree.

[20]

(gpg) is defined as 1 grain (64.8 mg) of calcium carbonate per U.S. gallon (3.79 litres), or 17.118 ppm.

Grain per gallon

a mmol/L is equivalent to 100.09 mg/L CaCO₃ or 40.08 mg/L Ca²⁺.

A degree of General Hardness ( or 'German degree (°dH, deutsche Härte))' is defined as 10 mg/L CaO or 17.848 ppm.

dGH

A Clark degree (°Clark) or English degrees (°e or e) is defined as one (64.8 mg) of CaCO₃ per Imperial gallon (4.55 litres) of water, equivalent to 14.254 ppm.

grain

A French degree (°fH or °f) is defined as 10 mg/L CaCO₃, equivalent to 10 ppm.

: 40^[33]

Canberra

: 10–26^[34]

Melbourne

: 39.4–60.1^[35]

Sydney

: 29–226^[36]

Perth

: 100^[37]

Brisbane

: 134–148^[38]

Adelaide

: 5.8–34.4^[39]

Hobart

: 31^[40]

Darwin

Fouling

Water purification

Water quality

Water treatment

. Akzo Nobel. Archived from the original on 30 August 2017. Retrieved 29 August 2017.

"Langelier Saturation Index (LSI) Calculato"

. Archived from the original on 3 February 2010. Retrieved 29 August 2017.

"Water hardness unit converter"

. Archived from the original on 2018-01-13. Retrieved 12 January 2018.