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When we mix together, from separate sources, the ions of a slightly soluble ionic salt, Ksp, and will continue to precipitate until the salt will precipitate if Qsp Qsp _Ksp. O is greater than; equals is less than; equals is less than; is greater than O equals; is less than equals; is greater than

When we mix together, from separate sources, the ions of a slightly soluble ionic salt, Ksp, and will continue to precipitate until the salt will precipitate if Qsp Qsp _Ksp. O is greater than; equals is less than; equals is less than; is greater than O equals; is less than equals; is greater than

Chemistry
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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### Precipitation of Slightly Soluble Ionic Salts When we mix together, from separate sources, the ions of a slightly soluble ionic salt, the salt will precipitate if \( Q_{sp} \) ____ \( K_{sp} \), and will continue to precipitate until \( Q_{sp} \) ____ \( K_{sp} \). #### Multiple Choice Options: 1. \(\circ\) is greater than; equals 2. \(\circ\) is less than; equals 3. \(\circ\) is less than; is greater than 4. \(\circ\) equals; is less than 5. \(\circ\) equals; is greater than #### Explanation: - \( Q_{sp} \) (the ion product) and \( K_{sp} \) (the solubility product constant) are key concepts in determining whether a precipitate will form in a solution. - **Step 1:** If \( Q_{sp} \) (the actual product of ion concentrations in the solution) is greater than \( K_{sp} \) (the maximum product of ion concentrations that can exist together in a solution without forming a precipitate), then the salt will precipitate because the solution is supersaturated. - **Step 2:** Precipitation will continue until \( Q_{sp} \) equals \( K_{sp} \), at which point the solution is saturated, and no more precipitate will form. In summary, when ions of a slightly soluble salt are mixed and \( Q_{sp} \) exceeds \( K_{sp} \), precipitation will occur until \( Q_{sp} \) decreases to equal \( K_{sp} \).
Transcribed Image Text:### Precipitation of Slightly Soluble Ionic Salts When we mix together, from separate sources, the ions of a slightly soluble ionic salt, the salt will precipitate if \( Q_{sp} \) ____ \( K_{sp} \), and will continue to precipitate until \( Q_{sp} \) ____ \( K_{sp} \). #### Multiple Choice Options: 1. \(\circ\) is greater than; equals 2. \(\circ\) is less than; equals 3. \(\circ\) is less than; is greater than 4. \(\circ\) equals; is less than 5. \(\circ\) equals; is greater than #### Explanation: - \( Q_{sp} \) (the ion product) and \( K_{sp} \) (the solubility product constant) are key concepts in determining whether a precipitate will form in a solution. - **Step 1:** If \( Q_{sp} \) (the actual product of ion concentrations in the solution) is greater than \( K_{sp} \) (the maximum product of ion concentrations that can exist together in a solution without forming a precipitate), then the salt will precipitate because the solution is supersaturated. - **Step 2:** Precipitation will continue until \( Q_{sp} \) equals \( K_{sp} \), at which point the solution is saturated, and no more precipitate will form. In summary, when ions of a slightly soluble salt are mixed and \( Q_{sp} \) exceeds \( K_{sp} \), precipitation will occur until \( Q_{sp} \) decreases to equal \( K_{sp} \).
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