General Chemistry w/Lab I
Experiment: Analysis of Hard Water
In this laboratory you will first learn to prepare a solution of EDTA and standardize it by titration. You will then use the same titration techniques to measure the hardness of two water samples, one an unknown solution prepared by the lab personnel, and the other will be one you bring from home. You will receive an additional grade dependent on your accuracy in determining the hardness of the unknown sample, as explained in the course syllabus.
People who live in other regions of the country, such as the Midwest, are very familiar with hard water. Soap doesnt lather well in hard water, and there is a constant build up of a ring of soap scum in the bathtub which has to be scrubbed to remove. What is hard water, and why does it have this effect? Hard water contains a higher than normal concentration of calcium and magnesium ions. These ions precipitate with soap, forming the soap scum build up. Additionally, since the soap molecules are being precipitated by the Ca+2 and Mg+2 ions, there is less soap available to form a lather. Taking a shower in hard water can be very frustrating!
Another effect of hard water is boiler scale. When hard water comes into contact with dissolved carbonates, a precipitate of insoluble calcium carbonate can form. This scale can build up on the inside of water pipes to such a degree that the pipes become almost completely blocked.
The following chart shows how hard water is classified. For reporting purposes, hardness is reported as parts per million (ppm) CaCO3. In other words, even though both Ca+2 and Mg+2 contribute to water hardness, it is reported as though all hardness ions are Ca+2 from CaCO3. Since Ca+2 and Mg+2 behave exactly the same, this convention is a convenient shorthand.
|Hardness (ppm CaCO3)||Classification
|< 15 ppm||Very Soft
|15 ppm - 50 ppm||Soft
|50 ppm - 100 ppm||Medium hard
|100 ppm - 200 ppm||Hard
|> 200 ppm||Very hard
The hardness of a sample of water can be measured by determining the concentration of the dissolved Ca+2 and Mg+2
ions. The procedure you will use is called a titration (See your text, pp. 153–154
& 164–165, for descriptions of titrations.) To analyze for Ca+2 and Mg+2 ions you will add a substance, Na2EDTA, which will react with the metal ions and remove them from solution. You will know when you have added enough EDTA by using an indicator in the solution; the indicator will change color when all of the Ca+2 and Mg+2 ions have reacted. The reactions of the metal ions are:
Ca+2 + EDTA-2 -----> CaEDTA
Mg+2 + EDTA-2 -----> MgEDTA
Na2EDTA is a complex molecule. Its name stands for ethylenediaminetetraacetic acid
disodium salt. The formula is Na2C10H14N2O8. Even though its name and structure are complex, you only need to know that:
- EDTA reacts with Ca+2 and Mg+2 in a one to one ratio
- The formula mass of Na2EDTA dihydrate is 372.24 g/mole
You will need to know the formula mass because you will be preparing a solution of EDTA with a specified molarity. By recording the exact volume of this solution that is needed to react with the Ca+2 and Mg+2, the concentration of these ions can be determined.
The strategy for this lab will be as follows:
- You will prepare a solution of EDTA with a concentration that is approximately 0.010 M. Because Na2EDTA comes as a dihydrate which is efflorescent
(efflorescent substances lose water of crystallization to the air), you cannot reliably weigh out the desired number of grams. The salt may be fully, or only partially, hydrated at the time you weigh it. The result is that you cannot be certain that you have prepared a 0.010 M solution. To determine the exact concentration of the solution you will need to standardize your solution in the manner described in the procedure.
- The lab technician will have prepared a standard solution which contains exactly 0.01000 M Ca+2. You will use the Ca+2 solution to standardize your EDTA solution. In other words, you will use the Ca+2 solution to determine the exact concentration of your EDTA solution by titrating the solution. You should find that the concentration of your solution is very close to 0.01000 M.
- Once the EDTA solution has been standardized, you will use it to determine the hardness of two samples of water. One will be an unknown which the instructor will provide. The other will be a sample of water that you will bring from home. The unknown analysis must come within 5% of the true value for full credit.
Finally, it is very important that all glassware you use be very clean and rinsed well. Metal ions from sources other than those solutions being tested will, of course, interfere with the tests. Furthermore, even soap seems to interfere, presumably because of ions in the soap.
In this lab you will work individually. Begin by obtaining the following from the lab cart
- A buret
- One 500 mL volumetric flask
- An unknown (record its number in your lab book)
Then complete the following steps:
- Use the 500 mL volumetric flask to prepare a 0.01000 M solution of Na2EDTA, as described by your instructor.
- Clean the buret carefully and rinse it twice with small amounts of your EDTA solution. Mount the buret on a ringstand using buret clamps. Use a funnel to fill the buret nearly to the top mark with your EDTA solution.
- Record the exact reading of the liquid level in the buret. N.B.: The burets can be read to the nearest 0.01 mL; be sure that all of your readings are accurate and consistent. Furthermore, be certain that you are reading the buret correctly. A common error is to read the buret upside down. Check with your instructor if you are unclear.
- Pipet 20.00 mL of the standard Ca+2 solution
(i.e., the 0.01000 M Ca+2) into a 125 mL Erlenmeyer flask. Add about 25 mL of deionized water.
- Add six (6) drops of indicator and one dropperfull of the buffer solution to the solution in the Erlenmeyer flask.
- Slowly add your EDTA solution from the buret until the endpoint is reached. The endpoint is the point at which the solution in the Erlenmeyer flask turns a distinct blue. As you near the endpoint, add EDTA drop by drop to avoid overshooting the endpoint. Record the exact liquid level of EDTA in the buret, and calculate the volume used in the titration by difference.
- Repeat steps 36 two more times, using a fresh sample of standard Ca+2
solution each time. If the three trials do not agree to within
2% of each other (i.e., ± 0.1 – 0.4 mL), do a fourth trial.
- Calculate the exact concentration of your EDTA solution, in moles/L, from the data recorded.
- Repeat steps 36 using your assigned unknown in place of the standard Ca+2 solution used in step 4. Three trials are recommended, however, it is up to you to decide how many trials are necessary to ensure reliable results.
- Repeat steps 36 using the sample of water from your home in place of the standard Ca+2 solution used in step 4. You may repeat this step if time permits and if you have sufficient EDTA solution. However, only one trial is required.
Analysis and Calculations
Be sure to do the following as part of your lab writeup:
- Calculate the concentration of your EDTA solution, using the known concentration of the standard Ca+2 solution, the volume of the standard Ca+2 solution, and the volume of your EDTA solution. Average the results from your three trials, assuming they agree within
± 2% of each other.
- Calculate the concentration of your unknown using the concentration of your EDTA solution, calculated above, and the amounts of solutions used.
- Do the same for your sample of water from home.
- Convert the concentration of Ca+2/Mg+2 ions to the standard reporting units of ppm CaCO3. (This calculation will be discussed as part of
a postlab discussion.) Once you have completed the calculations, report your results to your instructor. Remember that you will receive a separate grade based on the accuracy of your results.
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Last Revised: 1/2/12