Materials:

• 0.01 M Na₂H₂EDTA*2H₂O
• pH 10.0 buffer
• Eriochrome Black T or Calmagite indicator
• 50 % (w/w) NaOH
• Calcium Indicator
• Titration Equipment
• 3 water samples: distilled, softened (Biology building), unsoftened (off-campus rental house)

TOTAL HARDNESS DETERMINATION

1. Fill the burette with 0.01 M EDTA.

2. Pipette exactly 25 mL of tap water into a clean flask or beaker. Add 25 mL of deionized water. Add a few drops of the pH 10 buffer specifically prepared for hardness testing. The buffer should bring the pH to 10.0 to 10.1. (Think carefully about which buffers you will use to calibrate the pH meter.)

3. Add 6 drops of Calmagite indicator solution or an appropriate amount of dry-powder indicator formulation (three scoops).

4. Titrate immediately with EDTA (within 5 minutes of buffer addition). The solution will change from a wine red color to purple to dark blue. When the last reddish tinge disappears, add the last few drops at 3- to 5-second intervals. At the end point the solution is blue. Daylight, or a daylight fluorescent lamp, is recommended, because ordinary incandescent lights tend to produce a reddish tinge in the blue at the end point.

Alternatively, BEFORE you begin titrating, pour 10 mL of solution into a clean beaker, titrate the remaining solution until the blue endpoint, then add the 10 mL back to the original solution and titrate to the blue endpoint again.

5. Either repeat the tap water titration to increase your precision or compare your data with another group to test accuracy. If your data is very different from your neighbors repeat steps 1 through 4.

6. Perform a blank titration by using distilled water instead of tap water. Subtract the volume of EDTA needed to titrate the blank from the EDTA volumes for the tap water samples.

7. Record the mL of EDTA consumed and calculate total hardness as mg CaCO₃/L.

mg CaCO₃/L = (A x B x 1000 mL/L)/mL sample
A=mL EDTA
B=mg CaCO₃ equivalent to 1.00 mL EDTA

B can be calculated as follows:

B=100 mg CaCO₃/mmole x molarity of EDTA x 1000 mmole/mole x 1L/1000 mL for this experiment where the molarity of EDTA is 0.01 M
B=100 mg CaCO₃/mmole x 0.01 mole/L x 1000 mmole/mole x 1L/1000 mL
B=1.0 mg CaCO₃/mL EDTA

8. Fill the burette with EDTA.

9. Add to a 50 ml sample two mL of 50 % (w/w) NaOH enough to create a pH of 12 to 13. Add two scoops (0.1 to 0.2 g of solid or 1 to 2 drops of solution) of calcium indicator (Eriochrome Blue Black R or Solochrome Dark Blue).

When EDTA is added to water containing both calcium and magnesium, it combines first with the calcium. Calcium can be determined directly, with EDTA, when the pH is made sufficiently high that the magnesium is largely precipitated as the hydroxide and an indicator is used that combines with calcium only. Several indicators give a color change when all of the calcium has been complexed by the EDTA at a pH of 12 to 13.

10. Indicator should change from salmon pink to orchid purple at the end point. Record the ml of EDTA consumed.

11. Calculate the Ca²⁺ hardness as follows:

as mg CaCO₃/L (detailed in step 7) and
as mg Ca²⁺ = [mL EDTA X mg CaCO₃ equiv. to 1.0 mL EDTA X 400.8] / (mL sample).

1. (10 points) Given what you know about the proportion of alkalinity caused by CO₃^-2 at different pH's, and the solubility of metal carbonates, why is it important to get the titration over with in less than 5 minutes?

2. (5 points) What is the advantage to the alternative method mentioned in step 4 of the procedure?

3. (10 points) What is the magnesium, calcium, and total hardness of Meadville (unsoftened) and Biology building (softened) tap water? Show all your calculations, including those in steps 7 and 12 of the procedure. What is the relationship between the hardness of Meadville’s water and its ability to buffer acid precipitation? How does the acid precipitation get buffered?

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