What Happens When You Add Calcium Carbonate To Magnesium Oxide In Water

Hard Water

Last updated
Relieve as PDF

Page ID: 601

Hard water contains high amounts of minerals in the form of ions, especially the metals calcium and magnesium, which can precipitate out and cause bug in h2o cconducting or storing vessels like pipes. Hard water tin can be distinguished from other types of water past its metal, dry sense of taste and the dry out feeling it leaves on skin. It is responsible for the scum rings seen in bathtubs, also as the inability of soap to lather.

Types of Hard Water

Hard water is water containing high amounts of mineral ions. The most common ions institute in hard water are the metal cations calcium (Ca² ⁺) and magnesium (Mgⁱⁱ ⁺), though atomic number 26, aluminum, and manganese may too be found in sure areas. These metals are water soluble, meaning they will dissolve in water. The relatively high concentrations of these ions can saturate the solution and consequently cause the equilibrium of these solutes to shift to the left, towards reactants. In other words, the ions can precipitate out of the solution. This displacement of minerals from the solution is responsible for the calcination often seen on water faucets, which is a precipitation of calcium or magnesium carbonate. Hard water may also react with other substances in the solution, such as lather, and grade a precipitate called "scum." There are ii defined types of hard h2o, temporary and permanent, which are described below.

Temporary Hard Water

Temporary hard water is hard water that consists primarily of calcium (Ca² ⁺) and bicarbonate (HCO₃ ^-) ions. Heating causes the bicarbonate ion in temporary hard water to decompose into carbonate ion (CO₃ ^ii-), carbon dioxide (CO_two), and water (H₂O). The resultant carbonate ion (CO₃ ^two ^-) can and so react with other ions in the solution to form insoluble compounds, such as CaCO₃ and MgCO₃. The interactions of carbonate ion in the solution also crusade the well-known mineral build-up seen on the sides of pots used to eddy h2o, a rust known as "boiler calibration." Increasing the temperature of temporary difficult water, with its resultant decomposition of the bicarbonate ion, signifies a shift in the equilibrium equation (shown beneath). The high temperature causes the equilibrium to shift to the left, causing atmospheric precipitation of the initial reactants.

\[CaCO_{3 \; (due south)} + CO_{2 \; (aq)} + H_2O_{(l)} \rightleftharpoons Ca^{2+}_{(aq)} + 2HCO^-_{iii \; (aq)} \tag{i}\]

This shift is responsible for the white calibration observed in the boiling containers described above, every bit well as the mineral deposits that build up inside water pipes, resulting in inefficiency and fifty-fifty explosion due to overheating. The CaCO_iii or other scale does non completely dissolve dorsum into the h2o when it is cooled because it is relatively insoluble, as shown by its pocket-sized solubility constant. For this reason, this type of hard water is "temporary" because boiling can remove the hardness by displacing the offending ions from solution.

\[CaCO_{iii \; (south)} \rightleftharpoons Ca^{ii+}_{(aq)} + CO^{2-}_{3 \; (aq)} \tag{2a}\]

\[K_{(sp)} = 2.eight \times 10^{-ix} \tag{2b}\]

Permanent Hard Water

Permanent difficult water consists of high concentrations of anions, like the sulfate anion (SO_four ^two-). This type of hard water is referred to as "permanent" because, different temporary difficult water, the hardness cannot exist removed just past boiling the water and thereby precipitating out the mineral ions. Still, the name is deceiving equally "permanent" hard water can exist softened by other means. The scale caused by permanent hard water has detrimental effects similar to those seen with temporary hard water, such as obstruction of water flow in pipes. Permanent difficult water is besides responsible for the bathtub "ring," or soap scum, seen after showering or bathing. As previously mentioned, permanent hard water contains calcium and magnesium cations.These cations react with soap to grade insoluble compounds that are and so deposited on the sides of the tub. Additionally, the reaction of these cations with soap is the reason it is difficult for lather to foam or soap well in hard water. The equation below gives an example of the reaction of magnesium ion with components of soap, in this case stearate (C₁₈H₃₅O₂ ² ^-), to grade the insoluble compound magnesium stearate, which is responsible for the infamous soap scum.

\[2(C_{18}H_{35}O_2)^{ii-}_{(aq)} + Mg^{2+}_{(aq)} \longrightarrow Mg(C_{18}H_{35}O_2)_{ii \; (s)} \tag{iii}\]

Effects on the body

Though the taste of hard water may exist unpleasant to some, it has many health benefits when compared to soft h2o. Two of the near prevalent minerals in hard water are calcium and magnesium. Both calcium and magnesium are considered essential nutrients, meaning that they must be provided in the nutrition in lodge to maintain healthy body function. Calcium is a critical component of bones, and has many positive effects on the trunk, such as prevention of serious life-threatening and painful ailments like osteoporosis, kidney stones, hypertension, stroke, obesity, and coronary artery disease. Magnesium as well has positive wellness effects. Inadequate amounts of magnesium in the body increment the risks for some forms of health problems, such as hypertension, cardiac arrhythmia, coronary center disease, and diabetes mellitus. Studies washed on the wellness effects of hard and soft water have shown that people who drink greater amounts of soft h2o take much higher incidences of heart affliction, every bit well as higher blood pressure level and cholesterol levels, and faster heart rates than those who beverage mostly hard h2o. Furthermore, soft water is corrosive to pipes, which may allow for toxic substances like lead to contaminate drinking h2o.

How to soften hard water

Some wish to soften hard h2o to control its irritating, and in many cases damaging, effects. The diminished ability of soap to soap is not only annoying, but tin can besides be potentially harmful economically. Businesses that depend on the foaming of soap, such as car washes and pet groomers, may wish to soften hard water to avoid excessive use of soap due to a decreased ability to lather. Likewise, it is frequently necessary to soften water that comes into contact with pipes to avoid the destructive and compromising build-upwardly of deposits. Also, many people may find the calcifying effects that hard h2o has on faucets and other items unfavorable and choose to soften the water to forestall such mineral deposits from forming. Still others may dislike the viscous, dry feeling left past the precipitation of soap scum onto the skin. Whatever the reasons, there are many processes available to soften hard water.

Ion Exchange

One way to soften water is through a process chosen ion exchange. During ion exchange, the unwanted ions are "exchanged" for more than acceptable ions. In many cases, it is desirable to replace the hard h2o ions, such as Ca² ⁺ and Mg² ⁺, with more agreeable ions, like that of Na⁺. To practice this, the hard water is conducted through a zeolite or resin-containing column, which binds the unwanted ions to its surface and releases the more tolerable ions. In this process, the hard water ions get "fixed" ions because of their attachment to the resin material. These fixed ions readapt the desirable ions (Na⁺), now referred to as counterions, from the column, thus exchanging the ions in the water. This process is illustrated in Effigy ane.

Unfortunately, this procedure has the disadvantage of increasing the sodium content of drinking water, which could be potentially chancy to the health of people with sodium-restricted diets.

Lime Softening

Another process is called lime softening. In this procedure, the compound calcium hydroxide, Ca(OH)₂, is added to the difficult water. The calcium hydroxide, or "slaked lime," raises the pH of the water and causes the calcium and magnesium to precipitate into CaCO₃ and Mg(OH)₂. These precipitates tin and then be hands filtered out due to their insolubility in water, shown below past the small solubility constant of magnesium hydroxide (the solubility product constant for calcium carbonate is shown above). After precipitation and removal of the offending ions, acrid is added to bring the pH of the water back to normal.

\[Mg(OH)_{2 \; (s)} \rightleftharpoons Mg^{2+}_{(aq)} + 2OH^-_{(aq)} \tag{4a}\]

\[K_{(sp)} = 1.eight \times 10^{-11} \tag{4b}\]

Chelation

Chelating agents tin too be used to soften difficult water. Polydentate ligands, such as the popular hexadentate ligand EDTA, bind the undesirable ions in hard water. These ligands are especially helpful in binding the magnesium and calcium cations, which every bit already mentioned are highly prevalent in hard h2o solutions. The chelating agent forms a very stable ring complex with the metal cations, which prevents them from interacting with any other substances that may be introduced to the solution, such every bit lather. In this way, chelators are able to diminish the negative effects associated with hard water. A simplified equation representing the chelation of the metal calcium cation (Ca² ⁺) with the hexadentate ligand EDTA is shown below. The big value of the formation constant (One thousand_f) reflects the tendency of the reaction to keep to completion in the forrard direction.

\[Ca^{2+} + EDTA^{4-} \longrightarrow [Ca(EDTA)]^{two-} \tag{5a}\]

\[K_f = iv.ix \times 10^{10} \tag{5b}\]

Reverse Osmosis

The final process, reverse osmosis, uses loftier pressures to forcefulness the h2o through a semipermeable membrane. This membrane is more often than not intended to be impermeable to annihilation other than h2o. The membrane serves to filter out the larger ions and molecules responsible for the water'southward hardness, resulting in softened water. During this process, the water is forced from an area with a loftier concentration of solute in the form of dissolved metal ions and like compounds, to an area that is very depression in the concentration of these substances. In other words, the water moves from a state of hardness to a softer composition as the ions causing the h2o'southward hardness are prevented passage through the membrane. Contrary Osmosis does have a disadvantage of wasting wastewater compared to other water handling methods. This process is shown in Figure ii beneath. Notation that this figure describes the desalination of salt water. Withal, the process for softening hard h2o is the same.

Practice Issues

1. Proper noun the ii chief types of difficult water. (Highlight blue area for the answers)

Temporary and Permanent.

2. What makes "hard" water hard?

The presence of loftier concentrations of minerals, typically in the course of metal cations.

3. What are the ii most prevalent ions in hard water? How are these important to the proper part of the trunk?

Calcium (Ca² ⁺) and magnesium (Mg² ⁺) ion. They are important because both are essential nutrients, which means they are necessary for the proper part of the trunk and are too important for the prevention of many diseases and other ailments.

four. Proper name the iv described processes for softening difficult h2o. Information technology is also important to understand the essential steps of each process.

Ion exchange, chelation, lime softening, and reverse osmosis

v. What is a major disadvantage of ion exchange when Na⁺ is used as a counterion?

It increases the concentration of sodium in the h2o, a potential adventure for people with sodium-restricted diets.

References

Calcium and Magnesium in Drinking H2o: public health significance. Geneva, Switzerland: WHO press. 2009.
Health Effects of Drinking Water Treatment Technologies. Chelsea, MI: Lewis Publishers. 1989.
Lewis Alan, Scott. Condom Drinking H2o. San Francisco, CA: Scott Alan Lewis. 1996.
Crawford, T. and Grand. Crawford (1967). "Prevalence and pathological changes of ischaemic middle-disease in a hard-water and in a soft-water area." The Lancet 289(7484): 229-232.
Stitt, F., D. Clayton, et al. (1973). "Clinical and biochemical indicators of cardiovascular affliction among men living in difficult and soft water areas." The Lancet 301(7795): 122-126.
Gardiner, J. (1976). "Complexation of trace metals by ethylenediaminetetraacetic acid (EDTA) in natural waters." Water Inquiry 10(half dozen): 507-514.

Contributors

Andrea Kubisch, Courtney Korff (UCD)