Chemistry of Natural Waters

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The Chemistry of Natural Waters|
Chem 111 Sec 104|
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Hyunjung Hwang|
11/6/2012|

TA: Sarah Boehm, Group members: Rachel Hoffman, Dan Hirt|

Introduction
Water hardness is a major part of overall water quality that affects many industrial and domestic water users. Water is considered hard when there are high concentrations of the divalent cations Magnesium and Calcium; water hardness is considered as the sum of both the calcium and magnesium concentrations and expressed as calcium carbonate in milligrams per liter (mg/L).1 Hard water requires more soap and synthetic detergents for home laundry and washing, and contributes to scaling in boilers and industrial equipment. Calcium and magnesium carbonates tend to be deposited as off-white solids on the surfaces of pipes and the surfaces of heat exchangers. This precipitation (formation of an insoluble solid) is principally caused by thermal decomposition of bi-carbonate ions but also happens to some extent even in the absence of such ions. The resulting build-up of scale restricts the flow of water in pipes. In boilers, the deposits impair the flow of heat into water, reducing the heating efficiency and allowing the metal boiler components to overheat. In a pressurized system, this overheating can lead to failure of the boiler7. Water softening with a commercial water conditioning agent is used to reduce these affects. The goal of this experiment is to determine hardness of water samples and test the sample with a common softening technique. Water softening methods mainly rely on the removal of Ca2+ and Mg2+ from a solution. Commercial water-softening appliances intended for household use depend on an ion-exchange resin in which hardness ions are exchanged for sodium ions. Water softening may be desirable where the source of water is hard.8

Hardness in water is made up primarily of two elements: calcium and magnesium. Both naturally exist in groundwater and surface water supplies. These minerals, calcium and magnesium, dissolve in rainwater as it passes through soil and rock formations.2 Periods of low precipitation can cause hardness levels to increase for short periods of time. These levels usually decrease after rainfall or snowmelt due to dilution in the raw water sources. Calcium and magnesium in hard water pose no health problems and can promote stronger bones. However, hard water conveys some benefits to health by reducing the solubility of potentially toxic metal ions such as lead and copper. Conversely, removal of calcium and magnesium components through advanced processes has the potential to increase sodium level in the drinking water, which could be harmful for those who have high blood pressure. Softer water is also more corrosive and might shorten the life of home plumbing.3

The water samples for this experiment were obtained from Starbucks in downtown of State College, HUB building, and Pollock Dining Commons. All three locations are in State College. State College is located in the center of the state and is part of the Nittany Valley geographic area which is made up of the same rocks as piedmont lowland but is complexly folded and faulted because of the mountainous terrain.4 State College gets its water from seven wells and the Shingletown Gap Reservoir.5 To give a variety among samples, HUB and Pollock were selected to compare the hardness of water on campus. Starbucks was selected to compare the commercial café serving water and the fountain waters on campus.

Atomic Absorption spectroscopy (AA) and Ethylenediaminetetraacetic acid titration (EDTA) were used to determine the hardness of the water samples in the experiment. AA measures the amount of absorbance in the sample. This is done by turning the water into an aerosol and then shinning a monochromatic light through the sample to measure how much light was absorbed. The wavelength of the light corresponds to an excited electron state for the metal being examined and the...
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