Hard water analysis
Abstract
This experiment aimed at carrying out a laboratory analysis of the hardness of tap water.. The water sample test done involved a known concentration of CaCO3 solution to find out the EDTA indicator concentration. The conclusion made after the test indicated that the water was ‘hard water’.
Introduction
Water is described as being hard if the amount of magnesium and calcium dissolved in it is very high. This means that the water has high levels of these ions which in turn cause the hardness. The calcium and magnesium ions combine with soap molecules chemically causing the decreased cleansing action. The hardness in fresh water normally ranges between 15 – 375 mg/L as CaCO3 as a solution. The quantity of minerals found in tap water normally varies and is determined by the geological conditions of the area where water originates. Water that flows through areas with rich deposits of limestone absorbs the minerals from the limestone. Where there is large amount of white residue remaining after boiling, it means the water has large quantity of minerals. Hard water affects almost all cleansing tasks since the quantity of hardness minerals in the water has a lot of effects on how much soap is to be used for the purpose of cleansing. However, most of these minerals found in water are not toxic but on the other hand, they are beneficial and calcium in particular is a component of aquatic plants cell walls or bones of marine organisms. Magnesium is also vital nutrient for most plant life and a part of chlorophyll.
Procedure
- Add 25.0 ml of calcium solution in four different Erlenemyer flasks.
- Add 25.0 ml of tap water in four different Erlemeyer flasks
- Into each of these flasks, add buffer (pH – 10) solution amounting to 1ml
- Into each of the flasks, add 2 drops of EBT indicator
- Then Swirl the flask until the liquid turns purple
- Swirl flask till the liquid inside turns purple.
- Clean and fill 50 mL burette EDTA solution. Record initial reading.
- Use the first of the four flasks of Ca2+solution as a practice, run and titrate till the solution turns sky blue.
- Titrate the remaining three flasks of Ca2+.solution. For each record both the starting and ending volume of the burette.
- Repeat step 7 with a flask of tap water.
- Repeat step 8 with the remaining three flasks of tap water.
Results and analysis
The test of the water hardness was done in three trials which involved two readings. The readings include the initial Burette reading and the final Burette reading and the measurements were in ml.
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Trial 1 |
Trial 2 |
Trial 3 |
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Sample volume (mL) |
25.0 ml |
25.0 ml |
25.0 ml |
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Burette reading, initial (mL) |
0 |
1.80 ml |
1.50 ml |
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Burette reading, final (mL) |
1.80 ml |
1.50 ml |
1.20 ml |
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Volume of Na2H2Y titrant (mL) |
0.20 ml |
0.30 ml |
0.30 ml |
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Mol Na2HlY = mol hardening ions, Ca2+ and Mg2+ (mol/) |
5.0 x 10-5 mol
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7.5 x 10-5 mol |
7.5 x 10-5 mol |
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Mass of the same CaCO3 |
0.00500g |
0.0075g |
0.00757g |
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ppm CaCO3 (mgCaCO3/L sample |
200ppm |
302.7 ppm |
302.7 ppm |
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Average ppm CaCO3 |
248 |
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Average gpg CaCO3, |
15.7 gpg |
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Standard deviation of ppm CaCO3, |
58.9 |
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Relative standard deviation of ppm CaCO3 ( %RSD) |
22.0 % |
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Discussion
After calcium is added into water it reacts quite vigorously with water at room temperature in an exothermic reaction. The reaction gives of bubbles of hydrogen and forms a white precipitate which is calcium hydroxide together with an alkaline solution of the same precipitate. The calcium hydroxide is a bit soluble in water. Calcium occurs in water naturally having dissolved from various rocks such as marble limestone, fluorite, calcite and limestone. It is this element that determines the hardness of water, because it can be found in water as calcium ions (Beran, 251). Thus the calcium in water added to the flasks exists in form of calcium ions. The EDTA used was in anionic form. After the addition of the buffer solution to the flasks at pH 10, the Ca2+ (aq) ion forms a complex solution with the indicator as the CalIn+ that was seen to be wine red. As more of the stronger indicator is added, the complex CaY2- which is now blue replaces the Caln+ (aq). When the sharp color change is noticed, it marks the end point of the titration process. In the experiment the standardization of the NA2H2Y solution is as results of its reaction with a determined quantity of calcium ion in the primary standardized solution of calcium. The resulting standardized solution for Na2H2Y is then used in titration of the water sample hardening ions to the indicator.
The calcium complex has a higher stability constant than magnesium so calcium reacts first, and the magnesium later(Beran, 251). The reaction of magnesium and calcium and the indicator has the ration at around pH 10. In the experiment the average ppm of CaCO3 was found to be 148. The range for total water hardness in water ranges is usually between 15- 375 mg/l of the CaCO3. This means that the experiment proofed the water to be hard. Hardening ions that are present in the natural waters results from the rainwater that is slightly acidic and which normally flows over the deposits of compositions that are varying. This acidic water then reacts with the carbonate salts of magnesium and calcium and various rocks that normally contain iron (Beran, 251). The hardening ions, that is, Ca2+ and Mg2+ make up compounds that are insoluble with soups which in turn makes many of the detergent used not to be very effective. This hard water is also responsible for the formation of boiler scale on water heating appliances such as pots and kettles. Since the limestone deposits and other calcium minerals are naturally available in large quantities, it is not a wonder that the ions of calcium are and magnesium ions form the major components of the hard water’s dissolved solids (Kelter, Michael and Andrew, 767)
Conclusion
In order to determine the hardness of the tap water within the chemistry lab, two primary steps where required. First, a known concentration of approximately 0.002 M CaC0g solution was used in order to find the concentration of an EDTA indicator. That value was approximately 0.015 M. Using this information three samples of tap water where then titrated. The tap water hardness was determined to be a average of 175.3 ppm. This number indicates that the lab tap water is "Hard Water." However, it should be noted that relative standard deviation is -8% which indicates that the results are somewhat imprecise. This could be due to inconsistencies between the four experimenter’s methods of performing individual tasks of the experiment
References
Beran, J A. Laboratory Manual for Principles of General Chemistry. Hoboken, NJ: Wiley, 2011. Print. 250 The Geochemical Interpretation of Water Analyses. 250-251
Kelter, Paul B, Michael D. Mosher, and Andrew Scott. Chemistry: The Practical Science. Boston: Houghton Mifflin, 2009. Print. 766-767
