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The Chemistry Of Natural Water
INTRODUCTION The purpose of this experiment is to explore the hardness of the water on campus. Hard water has been a problem for hundreds of years. One of the earliest references to the hardness or softness of water is in Hippocrates' discourse on water quality in 5th Century B.C. Hard water causes many problems in both the household and the industrial world. One of the largest problems with hard water is that it tends to leave a residue when it evaporates. Aside from being aesthetically unpleasant to look at, the build up of hard water residue can result in the clogging of valves, drains and pipes. This build up is merely the accumulation of the minerals dissolved in natural water and is commonly called scale. Other than clogging plumbing, the build up of scale poses a large problem in the industrial world. Many things that are heated are often cooled by water running through piping. The build up of scale in these pipes can greatly reduce the amount of heat the cooling unit can draw away from the source causing a potentially dangerous situation. It also greatly reduces the heat efficiency of the container which holds the water, therefore much more energy is required to heat the item to the necessary temperature. In the industrial world, this could amount to large sums of money being thrown into wasted heat. In addition to clogging plumbing and reducing heating efficiency, the build up of hard water also adversely affects the efficiency of many soaps and cleansers. The reason for this is because hard water contains many divalent or sometimes even polyvalent ions. These ions react with the soap and although they do not form precipitates, they prevent the soap from doing its job. When the polyvalent ions react with the soap, they form an insoluble soap scum. This is once again quite unpleasant to look at and stains many surfaces. The reason for all these problems is that hard water tends to have higher than normal concentrations of minerals, and hence it leaves a considerable amount more residue than normal water. The concentration of these minerals is what is known as the water's Total Dissolved Solids or TDS for short. This is merely a way of expressing how many particles are dissolved in water. The TDS vary from waters of different sources, however, they are present in at least some quantity in all water, unless it has been passed through a special distillation filter. The relative TDS is easily measured by placing two drops of water, one distilled and one experimental on a hotplate and evaporating the two drops. You will notice that the experimental drop will leave a white residue. This can be compared to samples from other sources, and can be used as a crude way of measuring the relative TDS of water from a specific area. The more residue that is left behind, the more dissolved solids were present in that particular sample of water. Another, perhaps more quantitative way of determining hardness of water is by calculating the actual concentrations of divalent ions held in solution. This can be done in two ways. One is by serially titrating the water with increasing concentrations of indicator for Mg++ and Ca++ (we will be using EDTA). This will tell us the approximate concentration of all divalent ions. The method of serial titrations is accurate to within 10 parts per million (ppm). A second possible method for determining the hardness of water is by using Atomic Absorption Spectrophotometry or AA for short. AA is a method of determining the concentrations of individual metallic ions dissolved in the water. This is accomplished by sending small amounts of energy through the water sample causing the electrons to assume excited states. When the electrons drop back to their ground states, they release a photon of energy which can be measured by a machine and matched up to the corresponding element with the same E as was released. When the finding is related to the intensity of the light emitted and the amount of light absorbed, a concentration value is assigned. A quick overview of how the atomic absorption spectrophotometer works follows: First, the water sample is sucked up. Then the water sample is atomized into a fine aerosol mist. This is in turn sprayed into an extremely high intensity flame of 2300 C which is attained by burning a precise mix of air and acetylene. This mixture burns hot enough to atomize everything in the solution, solvent and solute alike. A light is emitted from a hollow cathode lamp. The light is then absorbed by the atoms and an absorption spectrum is obtained. This is matched with cataloged known values to attain a reading on concentration. Because there are so many problems with hard water, we decided that perhaps the water on Penn State's campus should be examined. My partners and I decided to test levels of divalent ions (specifically Mg++ and Ca++ ) in successive floors of dormitories. We hypothesized that the upper level dormitories would have lower concentrations of these divalent ions because being heavy metals, they would tend to settle out of solution. The Ca++ should settle out first seeing how it is heavier than the Mg++, but they will both decrease in concentration as they climb to higher floors in the dormitories. PROCEDURE We collected samples from around Hamilton Halls and West Halls. In order to be systematic, we collected samples in the morning from the water fountains near the south end of the halls. We also collected water samples from each floor for comparison. The reason we collected them in the morning was so that the Mg++ and Ca++ would be in noticeable quantities. We then went about and tested and analyzed via serial titrations and via Atomic Absorption Spectrophotometry. We also obtained a TDS sample merely for the sake of comparison, and to ensure that there were in fact dissolved solids in our water samples (without which this lab would become moot). For the serial titration, we merely mixed the water sample with EBT, and then with increasing concentrations of EDTA. The EBT served as an indicator to tell us when the concentrations of the EDTA and the divalent ions in solution were equal (actually it told us when Mg++ was taken out of solution but that served the same purpose). This allowed us to find the concentration of the divalent ions dissolved in solution. Based on this, we calculated the parts per million and the grains per gallon for each water sample. Finally, we took an AA reading for each sample. This gave us absorption values and concentration values for each of the two main metals we were observing; Ca++ and Mg++. We then plotted a graph of Atomic Absorption Standards using values that were given to us by the AA operator. These values helped us to calibrate the machine. The parts per million that we find will be based on plugging in the reported absorption value into the resulting curve from the graph of these values. The resulting concentration was used as the final value for the hardness for that particular sample. All calculations and conclusions were done based on these final values obtained for the concentration of Ca++ and Mg++. RESULTS Molarity x (100g CaCO3 / 1 mole CaCO3 ) x (1000 mg / 1g) = Xmg/1000g = ppm Grains/Gallon = ppm /17.1 Example: (1.6 x 10 -3 moles / 1 Liter) x (100g CaCO3 / 1 mole CaCO3 ) x (1000 mg / 1g) = 160 ppm 160 ppm/17.1 = 9.35 grains/gallon Conversion Factors Given by AA operator: Ca++ = 2.5 Mg++ = 4.2 Ca++ x 2.5 = CaCO3 hardness ppm value Mg++ x 100 x 4.5 = Mg CO3 hardness ppm value *NOTE: the Mg++ is x 100 because it was diluted before it was put into the AA. CONCLUSION Upon completion of this lab, it can be said that the data supports only half of the original hypothesis. Yes, the Ca++ did seem to decrease as the water got further from the source and climbed higher in the dormitories. However, the Mg++ did not. Instead it did quite the opposite and showed a general trend of increasing in concentration as it got further away from the source and higher in the dormitories. Perhaps a viable explanation could be attained if studies were done on the plumbing inside the building. Perhaps there is a high concentration of magnesium in the solder used to hold the pipes together. Perhaps it is not in the pipes but rather perhaps the people on the upper floors get up later and therefore at the time of collection, the water in the upper floors had been sitting longer than that on the lower floors. In either case, more investigation would have to be conducted in order determine what caused the unexpected results. In light of this discrepancy, the overall accuracy of the lab was very good. The numbers all seem to back each other up and correlate very well. As was mentioned in the previous section, the precision and accuracy with which this lab was carried out seems to indicate that there is very little source of error. Overall, it seems that the lab was quite well done. The hypothesis would have to be revised and as of this point, without further investigation, it would have to be reformulated to say that only the Ca++ would decrease in concentration whereas the Mg++ would increase. REFERENCES: Brown, Theodore L. et al. Chemistry The Central Science; Sixth Edition; Prentice Hall, Englewood Cliffs, NJ; 1994

 



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