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Introduction to Chemistry


Testing Water Hardness

Adapted from the American Chemical Society, Chemistry in Context Laboratory Manual, third ed., McGraw Hill, pub., 2000.


Water from lakes, rivers, or wells that contains significant concentrations of calcium ions (Ca2+) and magnesium ions (Mg2+) is know as hard water.  These ions are responsible for an assortment of problems ranging from aesthetic ones, such as bathtub rings and soap scum, to more critical ones like plugged steam lines and damaged hot water heaters.  Calcium ions get into water if the water comes in contact with limestone (calcium carbonate); magnesium ions likewise enter water when ground water passes through minerals that contain magnesium ions.  

In order to measure the combined concentrations of calcium and magnesium, you will be performing titrations.  In a titration, a known volume of a calcium or magnesium ion solution is measured out; then a solution which reacts with those ions (EDTA) is added slowly until just enough has been added to react with all of the calcium and/or magnesium ions as indicated by a color change in an indicator solution.  If one knows the volume of EDTA solution added and its concentration it is possible to calculate how much calcium and magnesium ions must have been present to react with the EDTA.  From that information it is possible to calculate the concentration of the calcium and magnesium in the original solution. 




Set up: 
  1. Obtain six small test tubes.
  2. Fill one test tube about 1/3 full with reference solution: 0.500 mg CaCO3 per mL.  Label it.
  3. Fill two other test tubes similarly with water samples.  Label them.
  4. Fill a third test tube about one third full with EDTA solution.  Label it.
  5. Fill one test tube about 1/4 full with pH 10 buffer solution.  Label it.
  6. Fill the last test tube about 1/4 full with calmagite indicator.  Label it.
  7. Place a labeled graduated pipet in each of the first three test tubes.
  8. Place a labeled thin stem pipet in each of the last three test tubes.

Before beginning: Technique practice

Using a graduated pipet and water, practice dispensing water one drop at a time.  Avoid getting half drops because of bubbles: if there are bubbles near the end of the pipet, dispense those drops into a well set aside for waste.

Practice filling a graduated pipet to the 1 mL line.

Part 1: Reference solution titration
  1. Test the color change of the indicator:
    1. Place 8-10 drops of water into one well of the wellplate.
    2. Add one drop of pH 10 buffer solution.
    3. Add one drop of calmagite indicator and stir.
    4. If the solution is not blue, add one drop EDTA solution and stir.  This is the color which will mark the end of each titration.  It should be clear blue with no trace of red.
  2. Place exactly 1 mL of calcium reference solution into each of 4 wells.
  3. To the first well, add one drop of buffer and one drop of indicator.
  4. To the first well, add EDTA drop by drop, counting the drops and stirring after each addition, until the color starts to change from red to purple.  Wait a few moments to see if the color continues to change (the reaction is slow); then slowly add additional drops until the color becomes pure blue. Keep track of the total number of drops used.  When you reach the end point (clear blue color), record the number of drops it took.
  5. Repeat steps 3-4 for the other 3 wells filled in step 2.
  6. If you are uncertain about the results for any of the 4 wells, repeat for a fifth well.

Part 2: Titration of water samples of unknown hardness
NOTE: Some water samples may not give a good color change due to iron or copper in the water.  These samples may need to be titrated to a purple color rather than to blue.  Check with the instructor if this appears to be the case.

  1. Place exactly 1 mL of a water sample into each of 4 wells.
  2. Titrate each well as in part 1 above, steps 3-4. 
  3. Repeat for one other sample.


  1. Look carefully at the data from each of the sets of titrations to see whether the results are consistent or whether any one result should be eliminated because it appears to be an "outlier", significantly different from the others.  If so, make a note beside it in your lab notebook.
  2. Calculate the average number of drops of EDTA needed to titrate each solution: the reference and each water sample, omitting any results you think are not valid (see step one above.)
  3. Calculate the mg of hardness per drop of EDTA: divide the hardness of the reference solution (mg CaCO3) by the average number of drops EDTA added in part 1.  This is the amount of Ca and Mg one drop of EDTA reacts with, your calibration factor.
  4. Calculate the hardness of each water sample: multiply the average number of drops of EDTA used for each sample by the calibration factor.  This will be the mg of Ca and Mg per mL of water for each sample.
  5. Convert the numbers from step 4 to mg per liter by multiplying by 1000.
  6. Compare your data with that of the rest of the class.


  1. Were there significant differences in hardness in the samples tested?
  2. In these titrations, there may be an uncertainty of one drop in identifying the exact end point.  If 50 drops of EDTA solution were used in a titration, what percent uncertainty in the hardness is contributed by the 1 drop uncertainty?
  3. What other techniques could be used to increase the accuracy of the titration?

Lab Report:

  1. An introduction explaining the purpose of the lab and the process of titration.   
    3 points
  2. A brief summary of what you did.    
    1 point
  3. Data    
    1 point
  4. Calculations   
    3 points  
  5. Conclusions: What was the hardness of each water sample you tested in mg/liter?  How do these compare with the samples tested by the rest of the class?    
    3 points
  6. Answers to questions.    
    3 point
  7. Metacognitive analysis: What contributed most to your learning in this lab?  What was confusing or unhelpful?  What would you change?
    1 point

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