The chemistry of natural waters

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The Chemistry of Natural Waters
Jennifer Hendriks
Chemistry 14
Spring 2006

Water is one of the most important, life-sustaining compounds that can be found nearly anywhere on the planet. This compound of hydrogen and oxygen has extraordinary properties that make it essential to all species’ well-being. A problem that occurs that reduces water’s usefulness and qualityis the reality of hard water. Water is classified as hard water when the concentrations of calcium ion (Ca2+), magnesium ion (Mg2+), or other divalent cations exceeds a set standard1. A common problem due to hard water is the formation of soap scum. However, when hard water is heated, such as the case with industrial-strength boilers and other water heating equipment, deposits form on keysurfaces. The deposits that form, also called scale, contain large quantities of calcite crystals (CaCO3)1. The formation of scale is due to the equation shown below2.
Ca2+ (aq) + 2HCO3-(aq) ( CaCO3(s) + H2O + CO2
The scale deposits narrow pipelines and cause blockages. Scale is expensive to remove if removal is even possible1.
In Experiment 10: The Chemistry of Natural Waters found in PSUChemtrek1, a group analysis was conducted for four different water samples taken from various sources. The four sources that were analyzed include tap water from Pittsburgh, Pennsylvania; tap water from State College, Pennsylvania; tap water from Delaware; and a sample taken from a nearby man-made lake, Stone Valley. The primary focus here will be the Stone Valley sample. It was hypothesized thatlake water will be the softest sample due to the naturally occurring low-concentration of ions1. The three tap water samples were expected to be rather close. This is because municipal water is treated similarly in varying locations. However, the State College tap water was expected to be harder than the Pittsburgh and Delaware taps, because State College, Pennsylvania is located in a valley andlower elevations usually produce harder water. This is because when the rainwater comes down, it usually must travel ‘downhill’ before settling at the base of the valley. While running, it picks up more and more cations on its way. At the base of the valley it is then pumped from the ground and treated in the case of municipal water. Because it is harder to begin with, it was also expected to beharder after treatment.
The hardness of each sample was tested for cation concentration using two primary methods. The first test consists of titrations using ethylenediaminetetraacetic acid (EDTA)1. EDTA titrations use a dye and acid-base indicator called eriochrome black T (EBT) that forms a blue chelate with the reaction of Mg2+.1 An EDTA titration is described below in the equation from PSUChemtrek:1
HD2- + Mg2+ + Ca2+ ( MgD- + H+ + Ca2+ ( CaEDTA + MgEDTA + HD2-
Blue Red (immediately) (eventually) Blue

The EDTA titration process begins with a known volume of a water sample to be analyzed. The sample’s pH must be adjusted to 10 with an NH3/NH4 buffer. The pH is adjusted to ensure an endpoint once the titration is complete3.Mg2+ in the water sample will react with the EBT to form a wine red chelate1. EDTA is now added and reacts first with Ca2+ and then with Mg2+. Once all of the Mg2+ has been reacted with, the indicator forms a blue color. This change marks the endpoint of the titration1.
The second test was that of Atomic Absorption Spectrophotometry (AA)1. The AA analysis was used to determine theconcentrations of cations, specifically Ca2+ and Mg2+. The process that AA uses to analyze a water sample is described below1. AA requires that the sample is heated to 2300ºC. From here, the sample is atomized and made into a fine aerosol. Monochromatic light of a specific wavelength is passed through the sample. The wavelength corresponds to the electronic energy level of a specific ion, such as...