(Start on Page 29-8, paragraph)

Estimating the sample size for Part 2a. You will use the conductivity reading to estimate a value of the sample size for the titration in part 2a, the next part of the experiment. The reasoning behind the estimate goes like this: is used as a proxy for the sum of concentrations of and because most natural waters contain more than and is usually the predominant anion. A 0.0010 M solution of has a conductivity of 208 microSiemens/cm at . We will aim to use at least 1.5 mL of EDTA titrant-in order to be able to do several titrations from one reagent filled buret- to titrate the in the sample. Assuming our EDTA solution is 0.0100 M , we estimate that a sample the size of 15 mL would be about right:

So we can establish an approximate working relationship between the conductivity- proportional to the concentration-and the sample size:

For convenience, you will round the calculated sample volume to the nearest mL (but you must be sure to measure the sample volume for all titrations to an accuracy of ).

The estimated volume is approximate because there may be cations beside and in the tap water or bottled water samples that also contribute to the conductivity. For example, some bottled waters, particularly mineral waters with a high content of dissolved solids, often contain significant amounts of or that contribute to the conductivity. For such waters, the concentration of would be smaller than estimated from the conductivity. Then our sample size would have to be made larger-but you wouldn't know this until you make a first trial titration.

Fill a microburet with EDTA. (Read the solution labels carefully so that you do not mistakenly fill the microburet with solution.) Then put into a Erlenmeyer flask a volume of tap water calculated using Equation (11), measured with an accuracy of (after rounding the calculated value to the nearest mL ). Add 15.0 mL of buffer. Using a graduated plastic transfer pipet, add 1.5 mL of . Finally, add 4 drops of calmagite indicator.

Calcium is titrated with EDTA at pH 12.5 , where magnesium ion is precipitated as . The endpoint is easily detected as a color change from pink to pure blue. To avoid
using diethylamine, as called for in the procedure originally described by Hildebrand and Reilley (see the Bibliography), we will obtain a solution pH of 12.5 by addition of NaOH .
Procedure:
Use the same microburet filled with EDTA that was used in part 2a. Put into a Erlenmeyer flask a volume of tap water equal to that used in part 2a, measured to an accuracy of . Add 15 mL of 0.10 M NaOH and stir the solution, then add 5 drops of calcon indicator. Immediately titrate with EDTA contained in the microburet. (The indicator is somewhat unstable at high pH .) The color change at the end point is from pink to pure blue. Repeat until you have completed at least two titrations that agree to within .

Use the sample volumes that were found satisfactory for the EDTA titrations of tap water and bottled water in part 2 a. Using a graduated cylinder or graduated pipet, place a sample of tap water in a Erlenmeyer flask. (Record the sample volume, accurate to .) Add drops of bromocresol green indicator. If a magnetic stirrer is available, place the flask on the stirrer over a sheet of white paper to better see the color change. Otherwise swirl the flask by hand while adding the titrant. Using a microburet filled with 0.020 M HCl , titrate the sample from blue to a greenish yellow. (The equivalence point comes at approximately pH 4.6 ). The ratio water sample should agree within for the two titrations. If it does, move on to the titration of bottled water; if not, titrate a third sample.

Repeat the entire procedure for at least two samples of bottled water, measuring the sample volume to an accuracy of .