Hard water

Hard water is water that has a high mineral content. 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. 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. 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 formed 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 varying degrees of hardness of water 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 frequently enter water supplies by leaching from minerals within aquifers.
Common calcium-containing minerals are calcite and gypsum. A common magnesium mineral is dolomite. 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 :
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 forming stalactites and stalagmites.
Calcium and magnesium ions can sometimes be removed by water softeners.
Permanent hardness 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 minerals. When dissolved, these types of minerals yield calcium and magnesium cations and carbonate and bicarbonate anions. 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 through the process of lime softening. Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.

Effects

With hard water, soap solutions form a white precipitate instead of producing lather, because the 2+ ions destroy the surfactant properties of the soap by forming a solid precipitate. 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 of precipitation of soap as a characteristic property of water that prevents the lathering of soap. Synthetic detergents do not form such scums.
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.
Hard water also forms deposits that clog plumbing. These deposits, called "scale", are composed mainly of calcium carbonate, magnesium hydroxide, and calcium sulfate. Calcium and magnesium carbonates tend to be deposited as off-white solids on the inside surfaces of pipes and heat exchangers.
This precipitation is principally caused by thermal decomposition of bicarbonate ions but also happens in cases where the carbonate ion is at saturation concentration. 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. 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, 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 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. 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 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 and magnesium 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 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, iron, aluminium, and manganese are also present at elevated levels in some locations.
The presence of iron characteristically confers a brownish colour to the calcification, instead of white.
Water hardness is often not expressed as a molar concentration, but rather in various units, such as degrees of general hardness, German degrees, parts per million, grains per gallon, English degrees, or French degrees. The table below shows conversion factors between the various units.
The various alternative units represent an equivalent mass of calcium oxide or calcium carbonate 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 is usually defined as 1 mg/L CaCO₃. It is equivalent to mg/L without chemical compound specified, and to American degree.
Grain per gallon is defined as 1 grain of calcium carbonate per U.S. gallon, 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 is defined as 10 mg/L CaO or 17.848 ppm.
A Clark degree or English degree is defined as one grain of CaCO₃ per Imperial gallon of water, equivalent to 14.254 ppm.
A French degree'' is defined as 10 mg/L CaCO₃, equivalent to 10 ppm.