Hard water

Not to be confused with heavy water.

Hard water caused calcification on this faucet.

Hard water is water that has high mineral content (mainly calcium and magnesium ions) (in contrast with soft water). Hard water minerals primarily consist of calcium (Ca²⁺), and magnesium (Mg²⁺) metal cations, and sometimes other dissolved compounds such as bicarbonates and sulfates. Calcium usually enters the water as either calcium carbonate (CaCO₃), in the form of limestone and chalk, or calcium sulfate (CaSO₄), in the form of other mineral deposits. The predominant source of magnesium is dolomite (CaMg(CO₃)₂). Hard water is generally not harmful to one's health.

The simplest way to determine the hardness of water is the lather/froth test: soap or toothpaste, when agitated, lathers easily in soft water but not in hard water. More exact measurements of hardness can be obtained through a wet titration. The total water 'hardness' (including both Ca²⁺ and Mg²⁺ ions) is read as parts per million (ppm) or weight/volume (mg/L) of calcium carbonate (CaCO₃) in the water. Although water hardness usually only measures the total concentrations of calcium and magnesium (the two most prevalent, divalent metal ions), iron, aluminium, and manganese may also be present at elevated levels in some geographical locations.

Hardness

Hardness in water is defined as the presence of multivalent cations. Hardness in water can cause water to form scales and a resistance to soap. It can also be defined as water that doesn’t produce lather with soap solutions, but produces white precipitate (scum). For example, sodium stearate reacts with calcium:

2C₁₇H₃₅COONa + Ca²⁺ → (C₁₇H₃₅COO)₂Ca + 2Na⁺

Types of hard water

In the 1960s, scientist Chris Gilby Jnr discovered that hard water can be categorized by the ions found in the water. A distinction is also made between 'temporary' and 'permanent' hard water.

Temporary hardness

Temporary hardness is caused by a combination of calcium ions and bicarbonate ions in the water. It can be removed by boiling the water or by the addition of lime (calcium hydroxide). Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.

The following is the equilibrium reaction when calcium carbonate (CaCO₃) is dissolved in water:

CaCO₃(s) + H₂CO₃(aq) ⇋ Ca²⁺(aq) + 2HCO₃^-(aq)

Upon heating, less CO₂ is able to dissolve into the water (see Solubility). Since there is not enough CO₂ around, the reaction cannot proceed from left to right, and therefore the CaCO₃ will not dissolve as rapidly. Instead, the reaction is forced to the left (i.e. products to reactants) to re-establish equilibrium, and solid CaCO₃ is formed. Boiling the water will remove hardness as long as the solid CaCO₃ that precipitates out is removed. After cooling, if enough time passes the water will pick up CO₂ from the air and the reaction will again proceed from left to right, allowing the CaCO₃ to "re-dissolve" into the water.

For more information on the solubility of calcium carbonate in water and how it is affected by atmospheric carbon dioxide, see calcium carbonate.

Permanent hardness

Permanent hardness is hardness (mineral content) that cannot be removed by boiling. 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 rises. Despite the name, permanent hardness can be removed using a water softener or ion exchange column, where the calcium and magnesium ions are exchanged with the sodium ions in the column.

Hard water causes scaling, which is the left over mineral deposits that are formed after the hard water had evaporated. This is also known as limescale. The scale can clog pipes, ruin water heaters, coat the insides of tea and coffee pots, and decrease the life of toilet flushing units.

Similarly, insoluble salt residues that remain in hair after shampooing with hard water tend to leave hair rougher and harder to untangle.

In industrial settings, water hardness must be constantly monitored to avoid costly breakdowns in boilers, cooling towers, and other equipment that comes in contact with water. Hardness is controlled by the addition of chemicals and by large-scale softening with zeolite and ion exchange resins.

Measurement

It is possible to measure the level of hard water by obtaining a free water testing kit. These are supplied by most water softening companies. There are several different scales used to describe the hardness of water in different contexts.

• Parts per million (ppm)
  Usually defined as one milligram of calcium carbonate (CaCO₃) per litre of water (the definition used below).
• grains/gallon (gpg)
  Defined as 1 grain (64.8 mg) of calcium carbonate per U.S. gallon (3.79 litres), or 17.118 ppm
• mmol/L (millimoles per litre)
  One millimole of calcium (either Ca²⁺ or CaCO₃) per litre of water corresponds to a hardness of 100.09 ppm or 5.608 dGH, since the molar mass of calcium carbonate is 100.09 g/mol.
• Degrees of General Hardness (dGH)
  One degree of General Hardness is defined as 10 milligrams of calcium oxide per litre of water, which is the same as one German degree (17.848 ppm).
• Various obsolete "degrees":
  • Clark degrees (°Clark)/English degrees (°E)
    One degree Clark is defined as one grain (64.8 mg) of calcium carbonate per Imperial gallon (4.55 litres) of water, equivalent to 14.254 ppm.
  • German degrees (Deutsche Härte, °dH)
    One degree German is defined as 10 milligrams of calcium oxide per litre of water. This is equivalent to 17.848 milligrams of calcium carbonate per litre of water, or 17.848 ppm.
  • French degrees (°f) (letter to be written in lowercase to avoid confusion with degree Fahrenheit — not always adhered to)
    One degree French is defined as 10 milligrams of calcium carbonate per litre of water, equivalent to 10 ppm.
  • American degrees
    One degree American is defined as one milligram of calcium carbonate per litre of water, equivalent to 1 ppm.

Although most of the above measures define hardness in terms of concentrations of calcium in water, any combination of calcium and magnesium cations having the same total molarity as a pure calcium solution will yield the same degree of hardness. Consequently, hardness concentrations for naturally occurring waters (which will contain both Ca²⁺ and Mg²⁺ ions), are usually expressed as an equivalent concentration of pure calcium in solution. For example, water that contains 1.5 mmol/L of elemental calcium (Ca²⁺) and 1.0 mmol/L of magnesium (Mg²⁺) is equivalent in hardness to a 2.5 mmol/L solution of calcium alone (250.2 ppm).

Because it is the precise mixture of minerals dissolved in the water, together with the water's pH and temperature, that determines the behaviour of the hardness, a single-number scale does not adequately describe hardness. Descriptions of hardness correspond roughly with ranges of mineral concentrations:

• Very soft: 0-70 ppm, 0-4 dGH
• Soft: 70-140 ppm, 4-8 dGH
• Slightly hard: 140-210 ppm, 8-12 dGH
• Moderately hard: 210-320 ppm, 12-18 dGH
• Hard: 320-530 ppm, 18-30 dGH
• Very hard >530 ppm, >30 dGH

Indices

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

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. 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 - pHs

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.

Langelier saturation index is defined as:

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

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.

Ryznar Stability Index (RSI)

The Ryznar stability index (RSI) uses a database of scale thickness measurements in municipal water systems to predict the effect of water chemistry.

Ryznar saturation index (RSI) was developed from empirical observations of corrosion rates and film formation in steel mains.

Ryznar saturation index is defined as:

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, the Stiff-Davis Index, and the Oddo-Tomson Index.

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."

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.

In a review by František Kožíšek, M.D., Ph.D. National Institute of Public Health, Czech Republic there is a good overview of the topic, and unlike the WHO, sets some recommendations 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.

A UK nationwide study, funded by the Department of Health, is investigating anecdotal evidence that childhood eczema may by correlated with hard water.

Very soft water can corrode the metal pipes in which it is carried and as a result the water may contain elevated levels of cadmium, copper, lead and zinc.

Softening

Main article: water softening

It is often desirable to soften hard water, as it does not readily form lather with soap. Soap is wasted when trying to form lather, and in the process, scum forms. Hard water may be treated to reduce the effects of scaling and to make it more suitable for laundry and bathing.

Process

A water softener, like a fabric softener, works on the principle of cation or ion exchange in which ions of the hardness minerals are exchanged for sodium or potassium ions, effectively reducing the concentration of hardness minerals to tolerable levels and thus making the water softer and giving it a smoother feeling.

The most economical way to soften household water is with an ion exchange water softener. This unit uses sodium chloride (table salt) to recharge beads made of the ion exchange resins that exchange hardness mineral ions for sodium ions. Artificial or natural zeolites can also be used. As the hard water passes through and around the beads, the hardness mineral ions are preferentially absorbed, displacing the sodium ions. This process is called ion exchange. When the bead or sodium zeolite has a low concentration of sodium ions left, it is exhausted, and can no longer soften water. The resin is recharged by flushing (often back-flushing) with saltwater. The high excess concentration of sodium ions alter the equilibrium between the ions in solution and the ions held on the surface of the resin, resulting in replacement of the hardness mineral ions on the resin or zeolite with sodium ions. The resulting saltwater and mineral ion solution is then rinsed away, and the resin is ready to start the process all over again. This cycle can be repeated many times.

The discharge of brine water during this regeneration process has been banned in some jurisdictions (notably California, USA) due to concerns about the environmental impact of the discharged sodium.

Potassium chloride (softener salt substitute) may also be used to regenerate the resin beads. It exchanges the hardness ions for potassium. It also will exchange naturally occurring sodium for potassium resulting in sodium-free soft water.

Some softening processes in industry use the same method, but on a much larger scale. These methods create an enormous amount of salty water that is costly to treat and dispose of.

Temporary hardness, caused by hydrogen carbonate (or bicarbonate) ions, can be removed by boiling. For example, calcium bicarbonate, often present in temporary hard water, may be boiled in a kettle to remove the hardness. In the process, a scale forms on the inside of the kettle in a process known as "furring". This scale is composed of calcium carbonate.

Ca(HCO₃)₂ → CaCO₃ + CO₂ + H₂O

Hardness can also be reduced with a lime-soda ash treatment. This process, developed by Thomas Clark in 1841, involves the addition of slaked lime (calcium hydroxide — Ca(OH)₂) to a hard water supply to convert the hydrogen carbonate hardness to carbonate, which precipitates and can be removed by filtration:

Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃ + 2H₂O

The addition of sodium carbonate also permanently softens hard water containing calcium sulfate, as the calcium ions form calcium carbonate which precipitates out and sodium sulfate is formed which is soluble. The calcium carbonate that is formed sinks to the bottom. Sodium sulfate has no effect on the hardness of water.

Na₂CO₃ + CaSO₄ → Na₂SO₄ + CaCO₃

Effects on skin

Some confusion may arise after a first experience with soft water. Hard water does not lather well with soap and leaves a "less than clean" feeling. Soft water lathers better than hard water but leaves a "slippery feeling" on the skin after use with soap. A certain water softener manufacturer contends that the "slippery feeling" after showering in soft water is due to "cleaner skin" and the absence of "friction-causing" soap scum.

However, the chemical explanation is that softened water, due to its sodium content, has a much reduced ability to combine with the soap film on your body and therefore, it is much more difficult to rinse off. Solutions are to use less soap or a synthetic liquid body wash.

Regional information

Hard water in Australia

Analysis of water hardness in major Australian cities by the Australian Water Association shows a range from very soft (Melbourne) to very hard (Adelaide). Total Hardness levels of Calcium Carbonate in ppm are: Canberra: 40; Melbourne: 10 - 26; Sydney: 39.4 - 60.1; Perth: 29 - 226; Brisbane: 100; Adelaide: 134 - 148; Hobart: 5.8 - 34.4; Darwin: 31.

Hard water 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, Calgary 165 ppm, Regina 202 ppm, Saskatoon < 140 ppm, Winnipeg 77 ppm, Toronto 121 ppm, Vancouver < 3 ppm, Charlottetown PEI 140 - 150 ppm.

Hard water in England and Wales

Information from the British Drinking Water Inspectorate shows that drinking water in England is generally considered to be 'very hard', with most areas of England, particularly the East, exhibiting above 200 ppm for the calcium carbonate equivalent. Wales, Devon, Cornwall and parts of North-West England are softer water areas, and range from 0 to 200 ppm . In the brewing industry in England and Wales, water is often deliberately hardened with gypsum in the process of Burtonisation.

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, Pacific Northwest, and Alaska regions. Hard and very hard waters are found in some of the streams in most of the regions throughout the country. Hardest waters (greater than 1,000 ppm) are in streams in Texas, New Mexico, Kansas, Arizona, and southern California.

See also

• Water softener
• Water quality
• Water treatment
• Water purification
• Zamzam Well

External links

• Alcoa Chemical, Langelier Saturation Index (LSI) Calculator
• Water hardness unit converter, Converter for hardness of water