Water “hardness” was analyzed in this experiment, through the determination of CaCO3 concentration. This was achieved by the titration of an unknown solution using a standardized 0.1M EDTA, and addition of Eriochrome Black T to the unknown, to indicate the endpoint of the titration. The average concentration of CaCO3 obtained was 1034 ppm, with a standard deviation of 2.4495. The results indicate that the unknown solution can be considered as hard water.
The hardness of water is defined in terms of its cation content, which includes calcium, magnesium, iron, zinc, and other polyvalent metal ions. These metal ions interfere with the use of the water for many applications. For example, these ions diminish the effectiveness of soap and detergent for cleansing operations; they diminish the drinking quality of water, and they contribute to the accumulation of insoluble salt deposits in storage vessels or plumbing1.
Water hardness can be easily determined by titration with the chelating agent EDTA (ethylenediaminetetraacetic acid), where its completely deprotonated anion forms a 1:1 complex with metal ions such as Ca2+ and Mg2+. In a titration, to establish the concentration of a metal ion, the EDTA that is added combines quantitatively with the cation to form the complex. The end point occurs when essentially all of the cation has reacted. The equation is shown below.
M + EDTA -----> M(EDTA)-complex
Both calcium and EDTA solutions are colorless, so it is necessary to add a metallochromic indicator to determine the endpoint of the titration, such as Eriochrome Black T which forms a stable wine-red complex ion (1). As EDTA is added, it binds calcium ion more strongly than Eriochrome Black T allowing the indicator to return (through shades of violet) to a pure blue color establishing the endpoint of the titration (2). Mg2+ is added to the titration solution to enhance the sharpness of the titration endpoint, since calcium ion does not form a stable complex with the indicator.
M2+ + HIn2+ + H2O MIn- + H3O+ (1) blue red
MIn- + HY3+ MY2- + HIn2- (2) red blue
Before the beginning of the experimental titrations, a standard calcium solution must be prepared in order to standardize the EDTA used in the titration of the unknown sample. In order to prepare the calcium standard, we must first accurately weigh 0.5 g of dried, pure CaCO3 into a 250mL beaker (this must be dried to a constant weight). Add approximately 25mL of distilled water, then 1mL of concentrated HCl carefully. Cover solution with watchglass spaced with glass hooks until all CaCO3 has dissolved. In order to remove the CO2 in the solution, we must next evaporate the solution on a hot plate to a volume of about 2mL maintaining the watchglass on the beaker (The acidification and boiling step is to remove carbonates, which if present, will precipitate CaCO3 when the solution is made basic. The precipitate obscures the end point). Rinse watchglass and transfer remaining solution quantitatively into a 500mL volumetric flask and make up to volume with water.
Next, the 0.1M EDTA solution to be used in titration must be prepared by weighing out approximately 4g of reagent grade disodium EDTA into a 250mL beaker. Add 0.1 g of magnesium chloride hexahydrate (to create more stable complex thus resulting in a better color change of the indicator, enhancing the sharpness of the endpoint), five pellets of NaOH and about 200mL of distilled water to dissolve. The EDTA will take approximately half an hour to dissolve. Filter the EDTA solution into a 1L bottle (if completely dissolved, there is...
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