Determination of the Solubility Product Constant for Calcium Sulfate: the Effect of Ionic Strengths of Electrolyte Solutions

Abstract In this experiment, the Ksp for calcium sulfate dihydrate, CaSO4·2H2O, by titrating 4 times a calcium sulfate dihydrate solution with diprotic EDTA, H2(EDTA)2-. For each trial we found the Ksp by means of molarities and activities. The results for the Ksp using only molarities was very different than the Ksp using activities. The average Ksp using molarity only was 2.26 x 10-4 and the average Ksp using activity turned out to be 2.31 x 10-5. The actual Ksp however, is 3.14 x 10-5. A percent error of 26.6 % was calculated.

Introduction

Experimental In this experiment, a saturated calcium sulfate was already made and ready to use. 25.00 mL of this solution was then mixed with 10 mL of an ammonia buffer and 1 drop of eriochrome black 1 indicator. We then filled our burette 0.0168 M EDTA solution and began to titrate with the calcium sulfate solution prepared. This procedure was done 4 times.
Results and Discussion
| |Ksp |
|Trial 1 |2.25 x 10-4 |
|Trial 2 |2.28 x 10-4 |
|Trial 3 |2.27 x 10-4 |
|Trial 4 |2.26 x 10-4 |

By finding the moles of EDTA used, the concentrations of Ca2+ and SO42- can be calculated. By finding the molarities of the two, the Ksp can also be calculated and the result are shown on Table 1. below
Table 1. Ksp Using Molarities

When comparing these values to the actual Ksp of 3.14 x 10-5 it is safe to say that they are not close at all. The problem here can be said to be that the