Hard water is water with a high mineral content (in contrast to "soft water"). Hard water is formed when water percolates through deposits of limestone, chalk or gypsum, which are largely made up of calcium and magnesium carbonates, bicarbonates and sulfates.
Drinking hard water may have moderate health benefits due to the mineral content. It can pose critical problems in industrial settings, where water hardness is monitored to prevent limescaling in boilers, cooling towers, and other water-handling equipment.
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. Wherever water hardness is a concern, water softening is commonly used to reduce hard water's adverse effects.
Origins 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 composed of hard, impervious, and calcium-poor rocks, only very low concentrations of divalent cations are found, and the water is termed soft water. Examples include Snowdonia in Wales and the Western Highlands in Scotland.
Areas with complex geology can produce water with varying degrees of hardness over short distances.
Types
Permanent hardness 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 often enter water supplies through leaching from minerals in aquifers. They are often accompanied by traces of Fe2+ ions, which stain the poorly soluble carbonate deposits rust-orange after being oxidised to less soluble Fe3+ ions.
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.
The following equilibrium reaction describes the dissolving and formation of calcium carbonate and calcium bicarbonate (on the right):
The reaction can go in either direction. Rain containing dissolved carbon dioxide can react with calcium carbonate, carrying calcium ions with it. The calcium carbonate may be redeposited as calcite as carbon dioxide is released into the atmosphere, sometimes forming stalactites and stalagmites in caves.
Calcium and magnesium ions can sometimes be removed by water softeners.
Permanent hardness (mineral content) is generally difficult to remove by boiling. 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 the permanent hardness of water can be removed using a water softener, or ion-exchange column.
Temporary hardness
Temporary hardness is caused by the presence of dissolved bicarbonate salts (calcium bicarbonate and magnesium bicarbonate). When dissolved, carbonate minerals yield calcium and magnesium cations (Ca2+, Mg2+), and carbonate and bicarbonate anions ( and ). The presence of these divalent alkaline earth 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 to degas or by the addition of lime (calcium hydroxide) through the process of lime softening. Boiling promotes the formation of carbonate from bicarbonate, precipitating calcium carbonate from solution and leaving water softer upon cooling.
Effects 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:
Hardness can thus be defined as the soap-consuming capacity of a water sample, or the capacity to precipitate soap, a characteristic property of water that prevents lathering. Synthetic detergents do not form such scums.
Because soft water contains fewer calcium ions, there is no inhibition of the lathering action of soaps, and no soap scum forms during normal washing. Similarly, soft water produces no calcium deposits in water heating systems.
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). 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 primarily caused by the thermal decomposition of bicarbonate ions, but it also occurs when the carbonate ion is at saturation concentration. The resulting build-up of scale restricts the flow of water in pipes. In boilers, deposits impair heat transfer to the water, reducing heating efficiency and allowing metal boiler components to overheat.
In a pressurized system, this overheating can lead to the failure of the boiler. The damage caused by calcium carbonate deposits varies according to the crystalline form, for example, calcite or aragonite.
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.
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 insoluble carbonates, and giving rise to 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 is in the range of 7.2 to 7.6.
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 of ion-exchange resins, which replace ions like Ca2+ by twice the number of mono cations 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.
Water that has been treated by a water softening may be termed softened water. In these cases, the water may also contain elevated levels of sodium or potassium and bicarbonate or chloride ions.
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". In fact, the United States National Research Council has found that hard water serves as a dietary supplement for calcium and magnesium.
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 recommend a level of hardness.
Recommendations have been made for the minimum and maximum 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.
Other studies have shown weak correlations between cardiovascular health and water hardness.
The prevalence of atopic dermatitis (eczema) in children may be increased by hard drinking water. 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, using water softeners at home does not reduce the severity of the symptoms.
Measurement
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 are 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; lowercase f is used to prevent confusion with degrees Fahrenheit). The table below shows conversion factors between the various units. {| class=" wikitable" |+ 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). It is equivalent to mg/L without chemical compound specified, and to American degree. Grain 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. 1 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, )) is defined as 10 mg/L CaO or 17.848 ppm. A Clark degree (°Clark) or English degree (°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 As it is the precise mixture of minerals dissolved in the water, together with water's pH and temperature, that determine the behaviour of the hardness, a single-number scale does not adequately describe hardness. However, the United States Geological Survey uses the following classification for hard and soft water: