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Every household and every factory uses water, and none of it is pure. One class of impurity that is of special interest is "hardness". Hard water is water that has high mineral content (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 sulphates. Calcium usually enters the water as either calcium carbonate (CaCO3), in the form of limestone and chalk, or calcium sulphate (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 measures only 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. Iron in this case is important for, if present, it will be in its tervalent form, causing the calcification to be brownish (the color of rust) instead of white (the color of most of the other compounds).
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 does not produce lather with soap solutions, but produces white precipitate (scum). For example, sodium stearate reacts with calcium:
2C17H35COONa + Ca2+ â†’ (C17H35COO)2Ca + 2Na+
Hardness of water may also be defined as the soap-consuming capacity of water, or the capacity of precipitation of soap as a characteristic property of water that prevents the lathering of soap.
Origin of water "hardness"
Carbon dioxide reacts with water to form carbonic acid (1) which at ordinary environmental pH exists mostly as bicarbonate ion (2). Microscopic marine organisms take this up as carbonate (4) to form calcite skeletons which, over millions of years, have built up extensive limestone deposits. Groundwaters, made slightly acidic by CO2 (both that absorbed from the air and from the respiration of soil bacteria) dissolve the limestone (3), thereby acquiring calcium and bicarbonate ions and becoming "hard". If the HCO3- concentration is sufficiently great, the combination of processes (2) and (4) causes calcium carbonate ("lime scale") to precipitate out on surfaces such as the insides of pipes. (Calcium bicarbonate itself does not form a solid, but always precipitates as CaCO3.)
Indications of Hard Water
Hard water interferes with almost every cleaning task from laundering and dishwashing to bathing and personal grooming. Clothes laundered in hard water may look dingy and feel harsh and scratchy. Dishes and glasses may be spotted when dry. Hard water may cause a film on glass shower doors, shower walls, bathtubs, sinks, faucets, etc. Hair washed in hard water may feel sticky and look dull. Water flow may be reduced by deposits in pipes.
Dealing with hard water problems in the home can be a nuisance. The amount of hardness minerals in water affects the amount of soap and detergent necessary for cleaning. Soap used in hard water combines with the minerals to form a sticky soap curd. Some synthetic detergents are less effective in hard water because the active ingredient is partially inactivated by hardness, even though it stays dissolved. Bathing with soap in hard water leaves a film of sticky soap curd on the skin. The film may prevent removal of soil and bacteria. Soap curd interferes with the return of skin to its normal, slightly acid condition, and may lead to irritation. Soap curd on hair may make it dull, lifeless and difficult to manage.
When doing laundry in hard water, soap curds lodge in fabric during washing to make fabric stiff and rough. Incomplete soil removal from laundry causes graying of white fabric and the loss of brightness in colors. A sour odor can develop in clothes. Continuous laundering in hard water can shorten the life of clothes. In addition, soap curds can deposit on dishes, bathtubs and showers, and all water fixtures.
Hard water also contributes to inefficient and costly operation of water-using appliances. Heated hard water forms a scale of calcium and magnesium minerals that can contribute to the inefficient operation or failure of water-using appliances. Pipes can become clogged with scale that reduces water flow and ultimately requires pipe replacement.
Types of hardness
A distinction is made between 'temporary' and 'permanent' hard water.
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 limewater (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) + CO2(aq) + H2O â‡‹ 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 is hardness (mineral content) that cannot be removed by boiling. It is usually caused by the presence in the water of calcium and magnesium sulfates and/or chlorides 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 (Na2Al2Si2O8.xH2O) and ion exchange resins.
Measurement of hardness
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. Descriptions of hardness correspond roughly with ranges of mineral concentrations:
It is possible to measure the level of total hardness in water by obtaining a total hardness water testing kit. These kits measure the level of calcium and magnesium in the water. Temporary hardness test kits do not normally measure calcium and magnesium levels but normally use an approximation based on some form of alkalinity test. Measuring temporary hardness accurately would involve a series of tests to work out how much bicarbonates and carbonates are present and how much calcium and magnesium is present and what percentage combination there is. In most cases, the temporary hardness kit is a good approximation, but anions such as hydroxides, borates, phosphates can have quite an effect on temporary hardness test kits.
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 per 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 alternative "degrees":
Clark degrees (°Clark)/English degrees (°e or 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 or 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 or f) (letter written in lower-case 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.
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).
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. 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.
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. It 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
Conventional water softening
Most conventional water-softening devices depend on a process known as ion-exchange in which "hardness" ions trade places with sodium and chloride ions that are loosely bound to an ion-exchange resin or a zeolite (many zeolite minerals occur in nature, but specialized ones are often made artificially.)
The illustration depicts a negatively-charged zeolite to which [positive] sodium ions are attached. Calcium or magnesium ions in the water displace sodium ions, which are released into the water. In a similar way, positively-charged zeolites bind negatively-charged chloride ions (Cl-), which get displaced by bicarbonate ions in the water. As the zeolites become converted to their Ca2+ and HCO3- forms they gradually lose their effectiveness and must be regenerated. This is accomplished by passing a concentrated brine solution though them, causing the above reaction to be reversed. Herein lies one of the drawbacks of this process: most of the salt employed in the regeneration process gets flushed out of the system and and is usually released into the soil or drainage system- something that can have damaging consequences to the environment, especially in arid regions. For this reason, many jurisdications prohibit such release, and require users to dispose of the spent brine at an approved site or to use a commercial service company.
It is often considered desirable to soften hard water. This is because the calcium and magnesium causing hardness partly block the oil emulsifying action simple soap formulations use in the cleaning action. The calcium and magnesium form an insoluble precipitate observed as a soap scum and extra large amounts of soap have to be used to counteract this. Most modern soaps and detergents contain ingredients that at least partly prevent this effect and detergents are available that are chemically completely unaffected by the hardness. This makes hardness removal/softening an optional rather than a necessary water treatment except possibly in the case of extremely hard water. 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. Another reason for this is to avoid adding sodium or potassium from the softener to cold water taken for human consumption while still providing softening for hot water used in washing and bathing.
Process of water softening
A water softener works on the principle of cation or ion exchange in which ions of the hardness minerals (mainly calcium and magnesium ions) 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 â†’ 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 â†’ 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 potasium sulphate 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 â†’ 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 "clean" feeling. Soft water lathers better than hard water but leaves a "slippery feeling" on the skin after use with soap. Some providers of water softening equipment claim that the "slippery feeling" after showering in soft water is due to "clean skin" and the absence of 'friction-causing' soap scum.
However, the chemical explanation is that softened water, because of its sodium content, has a much reduced ability to combine with the soap film on the body; therefore, the soap is much more difficult to rinse off. Solutions are to use less soap or a synthetic liquid body wash.
Potential Health Effectss
Hard water is not a health hazard. In fact, the National Research Council in US (National Academy of Sciences) states that hard drinking water generally contributes a small amount toward total calcium and magnesium human dietary needs. They further state that in some instances, where dissolved calcium and magnesium are very high, water could be a major contributor of calcium and magnesium to the diet.
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 FrantiÅ¡ek KoÅ¾íÅ¡ek, M.D., Ph.D. National Institute of Public Health, Czech Republic there is an overview of the topic which, 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.
While some studies suggest a correlation between hard water and lower cardiovascular disease mortality, other studies do not suggest a correlation. The National Research Council states that results at this time are inconclusive and recommends that further studies should be conducted.
A UK nationwide study, funded by the Department of Health, is investigating anecdotal evidence that childhood eczema may be correlated with hard water.
Very hard 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.