The potential difference between a potassium ion and saturated calomel electrode immersed in 100.0 ml of 1.073 x 10 M KCI was 362 mV. When 594 mg of NH4CI was added to the KCI solution (no volume change), the potential difference was 386 mV. Calculate the selectivity coefficient kx*,NH, for K* over NH4. Assume that there is no loss of NHa due to acid-base equilibrium effects.

The potential difference between a potassium ion and saturated calomel electrode immersed in 100.0 ml of 1.073 x 10 M KCI was 362 mV. When 594 mg of NH4CI was added to the KCI solution (no volume change), the potential difference was 386 mV. Calculate the selectivity coefficient kx,NH, for K over NH4. Assume that there is no loss of NHa due to acid-base equilibrium effects.

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

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Concept explainers

Potentiometric Titrations

Potentiometric titration is one of the important titration methods and is mainly carried out to characterize acid solutions. It follows a method similar to that of the normal titration technique, but the main difference is that there is no indicator used in the potentiometric titrations. Instead of using an indicator, electrodes are employed for monitoring the whole titration process. A scientist named Robert Behrend did the potentiometric titration for the first time, and it was done using the mercurous solution and potassium chloride.

Question

The potential difference between a potassium ion and saturated calomel electrode immersed in 100.0
ml of 1.073 x 103 M KCI was 362 mV. When 594 mg of NHẠCI was added to the KCI solution (no volume
change), the potential difference was 386 mV. Calculate the selectivity coefficient kxt,NH, for K* over
NH4. Assume that there is no loss of NH4 due to acid-base equilibrium effects.

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