Hardness Of Water.docx

Hardness of water

Historically, water “hardness” was defined in terms of the capacity of cations in the water to replace the sodium or potassium ions in soaps and form sparingly soluble products that cause “scum” in the sink or bathtub. Most multiply charged cations share this undesirable property. In natural waters, however, the concentrations of calcium and magnesium ions generally far exceed those of any other metal ion. Consequently, hardness is now expressed in terms of the concentration of calcium carbonate that is equivalent to the total concentration of all the multivalent cations in the sample. The determination of hardness is a useful analytical test that provides a measure of the quality of water for household and industrial uses. The test is important to industry because hard water, on being heated, precipitates calcium carbonate, which clogs boilers and pipes. Water hardness is usually determined by an EDTA titration after the sample has been buffered to pH 10. Magnesium, which forms the least stable EDTA complex of all of the common multivalent cations in typical water samples, is not titrated until enough reagent has been added to complex all of the other cations in the sample. Therefore, a magnesium-ion indicator, such as Calmagite or Eriochrome Black T, can serve as indicator in water-hardness titrations. Often, a small concentration of the magnesium-EDTA chelate is incorporated in the buffer or in the titrant to ensure the presence of sufficient magnesium ions for satisfactory indicator action. Feature 177 gives an example of a kit for testing household water for hardness. -

Skoog

https://www.slideshare.net/ratneshkanungo007/determining-of-hardness-of-water Hardness of water, historically, refers to water’s capacity to replace sodium and potassium to form a soluble scum in the sink or bathtub. However in natural water there is a relatively high concentration of Magnesium and calcium than that of any other metal ions present (Skoog, 2013). Hardness of water is defined as the concentration of alkali earth metals, which mostly are Ca2+ and Mg2+, present in water which can lather from soap (Harris, 2010). There are three kinds of hardness: first is the Temporary hardness which refers to hardness induced by bicarbonate minerals such as calcium bicarbonate. This hardness can be reduced by addition of calcium hydroxides. Permanent hardness is a hardness that cannot be removed. Though ions causing the permanent hardness of water can be removed using a water softener. Total permanent hardness is calcium hardness + magnesium hardness. The hardness of these two is expressed as the equivalent of calcium carbonate.

Since EDTA can form complex ions with Ca2+ and Mg2+, it can be used to quantify the presence of calcium and magnesium in water. Thus, the total hardness of water can be approximated through the titration with standard EDTA. The solution should have a buffer of 10pH. There are excess of H+ in the solution, without this buffer, the solution will radically become acidic.

Historically, hardness of water is was the measure of the capacity of water to react with soap producing white, foamy bubble called lather. On a later date, this was defined to be the amount of metal ions present in water (Brown, 2011). Water of this much “hardness” is inefficient enough for use as opposed to “soft” water. To identify the quality of water, the quantitative analysis of its components can be done. Amongst the many contributors to the hardness of water, calcium and magnesium ions are the most abundant contributors present. Since calcium is one of the major contributors to the hardness of water, the analysis of the hardness of water can be done by determining the amount of Calcium ions present in the solution, and this is conventionally achieved with using complexometric titration of calcium ions (WHO, 2011). If the measured amount is less than 60 mg CaCO3/L, then the water is classified as soft. Between 60-12- mg CaCO3/ L, the water is classified as moderately hard. Finally, if the amount is greater than 180 g CaCO3/ L then the water is hard (McGowan, 2000). The titration of calcium ions can be done using complexometric titration of it with EDTA as titrant. At pH 10, the formation of Ca-EDTA complex is favored (log Kf = 10.69) (Martell & Smith, 1974). The indicator to be used is EBT, and the complexation of calcium ions to EBT forms a red color Reaction Prior to titration, the titrant already has a small amount of Mg ions to form a complex with EDTA. MgEDTA complex titrant can be added because as opposed to Ca-EDTA, Mg-EDTA is less favored (log Kf = 8.79), therefore when Mg-EDTA is added to the sample, it will prefer to form the Ca-EDTA complex. During the beginning of titration, the Ca-In complex won’t be reactive due to the relatively large amount of unreacted calcium ions, and these unreacted ions will replace the magnesium ions in the EDTA complex. Reacition As more amount of EDTA is added, it would reach a point where the Ca-In complex will be replaced to form Mg-In complex, still giving a red color. Reaction Since Mg-In complex is much more stable as opposed to Mg-EDTA, it will stay as it is until the equivalence point is reached where Mg-EDTA is formed, liberating the indicator, changing its color to blue, marking the end of titration. Reaction

References Brown, T. E. (2011). Chemistry: The Central Science. New Jersey: Prentice Hall. Martell, A. E., & Smith, R. M. (1974). Critical Stability Constants. In A. E. Martell, Kf of EDTA (pp. 204211). New York: Plenium Press. McGowan, N. (2000). Water processing: residential, commercial, light-industrial, 3rd ed. Illinois: Water Quality Association.

WHO. (2011). Hardness in Drinking Water. Guidelines for Drinking-water Quality, 1-3.