Tweet Hard water is water that has high mineral content. Hard drinking water is generally not harmful to one’s health, but can pose serious 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 [...]
Hard drinking water is generally not harmful to one’s health, but can pose serious 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 suds formation when soap is agitated in water. Wherever water hardness is a concern, water softening is commonly used to reduce hard water’s adverse effects.
Sources of hardness
Water’s hardness is determined by the concentration of multivalent cations in the water. Multivalent cations are cations (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 also contain few ions.
The following equilibrium reaction describes the dissolving/formation of calcium carbonate scales:
CaCO3 + CO2 + H2O ⇋ Ca2+ + 2HCO3−
Calcium carbonate scales formed in water-heating systems are called limescale.
Calcium and magnesium ions can sometimes be removed by water softeners.
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 softening 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.
Permanent hardness is hardness (mineral content) that cannot be removed by boiling. When this is the case, it is usually caused by the presence of calcium and magnesium sulfates and/or chlorides in the water, which become more soluble as the temperature increases. Despite the name, the hardness of the water can be easily removed using a water softener, or ion exchange column.
Effects of hard water
With hard water, soap solutions form a white precipitate (soap scum) instead of producing lather. This effect arises 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- + Ca2+ → (C17H35COO)2Ca
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.
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 surfaces of pipes and the surfaces of heat exchangers. This precipitation (formation of an insoluble solid) is principally caused by thermal decomposition of bi-carbonate ions but also happens to some extent even in the absence of such ions. 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. The damage caused by calcium carbonate deposits varies depending on 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 the insoluble carbonates, giving rise to the turbidity. This often results from the alkalinity (the hydroxide concentration) being excesively high (pH > 7.6). Hence, a common solution to the problem is to, while maintaining the chlorine concentration at the proper level, raise the acidity (lower the pH) by the addition of hydrochloric acid, the optimum value being in the range of 7.2 to 7.6.
For the reasons discussed above, 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 practiced, 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.
The World Health Organization says that “there does not appear to be any convincing evidence that water hardness causes adverse health effects in humans”.
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 were inadequate to allow for a recommendation for a level of hardness.
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.
Other studies have shown weak correlations between cardiovascular health and water hardness.
Some studies correlate domestic hard water usage with increased eczema in children.
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.
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 (°f). The table below shows conversion factors between the various units.
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.
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 (°F or f) is defined as 10 mg/L CaCO3, equivalent to 10 ppm. The lowercase f is often used to prevent confusion with degrees Fahrenheit.
Because it is the precise mixture of minerals dissolved in the water, together with the water’s pH and temperature, that determines the behavior of the hardness, a single-number scale does not adequately describe hardness.
Langelier Saturation Index (LSI)
The Langelier Saturation Index (sometimes Langelier Stability Index) is a calculated number used to predict the calcium carbonate stability of water. 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). The LSI is expressed as the difference between the actual system pH and the saturation pH:
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.
It is also worth noting that 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.
Hard water in the United States
More than 85% of American homes have hard water. 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, and southern California.
Measuring Hardness and LSI
The Myron L Ultrameter III 9PTK measures water hardness and LSI, as well as 7 other water quality parameters.
Measures 9 Parameters: Conductivity, Resistivity, TDS, Alkalinity, Hardness, LSI, pH, ORP/Free Chlorine, Temperature
LSI Calculator for hypothetical water balance calculations
Wireless data transfer capability with bluDock option
Auto-ranging delivers increased resolution across diverse applications
Adjustable Temperature Compensation and Cond/TDS conversion ratios for user-defined solutions
Nonvolatile memory of up to 100 readings for stored data protection
Date & time stamp makes record-keeping easy
pH calibration prompts alert you when maintenance is required
Auto-off minimizes energy consumption
Low battery indicator
(Includes instrument with case and solutions)
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