Will a precipitate be observed if 0.10 mol Ag* and 0.001 mol SO4²- are added to make 1.0 L of solution? (Ksp = 1.4 x 10-5) yes, because Q < Ksp no, because Q > Ksp no, because Q < Ksp
Will a precipitate be observed if 0.10 mol Ag* and 0.001 mol SO4²- are added to make 1.0 L of solution? (Ksp = 1.4 x 10-5) yes, because Q < Ksp no, because Q > Ksp no, because Q < Ksp
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 Chemistry Problem**
**Question:**
Will a precipitate be observed if 0.10 mol Ag⁺ and 0.001 mol SO₄²⁻ are added to make 1.0 L of solution? (Ksp = 1.4 × 10⁻⁵)
**Options:**
- ○ yes, because Q < Ksp
- ○ no, because Q > Ksp
- ○ no, because Q < Ksp
When attempting to determine if a precipitate will form in a solution, it is essential to compare the ion product (Q) with the solubility product constant (Ksp). If Q > Ksp, a precipitate will form because the solution is supersaturated. If Q < Ksp, no precipitate will form because the solution is unsaturated. If Q = Ksp, the solution is at equilibrium and is saturated.
**Calculation Process:**
To solve this, first, determine the molar concentrations of Ag⁺ and SO₄²⁻ in the solution. Since the volume of the solution is 1.0 L:
\[ [Ag^+] = \frac{0.10 \text{ mol}}{1.0 \text{ L}} = 0.10 \text{ M} \]
\[ [SO₄²⁻] = \frac{0.001 \text{ mol}}{1.0 \text{ L}} = 0.001 \text{ M} \]
Next, calculate the reaction quotient (Q) for the precipitation reaction:
\[ Q = [Ag^+] \times [SO₄²⁻] \]
\[ Q = (0.10 \text{ M}) \times (0.001 \text{ M}) \]
\[ Q = 1.0 \times 10^{-4} \]
Compare Q with the given Ksp:
\[ Q = 1.0 \times 10^{-4} \]
\[ Ksp = 1.4 \times 10^{-5} \]
Since \( Q > Ksp \), the solution is supersaturated, and a precipitate will form. Therefore, the correct option is:
- ○ no, because Q > Ksp](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Ff3c22a28-8e9a-42a1-ac26-114f2e495bee%2Fcfaf85df-ddf4-4899-98bf-cd115f3909c0%2Fxb9zmbs_processed.png&w=3840&q=75)
Transcribed Image Text:**Precipitation Chemistry Problem**
**Question:**
Will a precipitate be observed if 0.10 mol Ag⁺ and 0.001 mol SO₄²⁻ are added to make 1.0 L of solution? (Ksp = 1.4 × 10⁻⁵)
**Options:**
- ○ yes, because Q < Ksp
- ○ no, because Q > Ksp
- ○ no, because Q < Ksp
When attempting to determine if a precipitate will form in a solution, it is essential to compare the ion product (Q) with the solubility product constant (Ksp). If Q > Ksp, a precipitate will form because the solution is supersaturated. If Q < Ksp, no precipitate will form because the solution is unsaturated. If Q = Ksp, the solution is at equilibrium and is saturated.
**Calculation Process:**
To solve this, first, determine the molar concentrations of Ag⁺ and SO₄²⁻ in the solution. Since the volume of the solution is 1.0 L:
\[ [Ag^+] = \frac{0.10 \text{ mol}}{1.0 \text{ L}} = 0.10 \text{ M} \]
\[ [SO₄²⁻] = \frac{0.001 \text{ mol}}{1.0 \text{ L}} = 0.001 \text{ M} \]
Next, calculate the reaction quotient (Q) for the precipitation reaction:
\[ Q = [Ag^+] \times [SO₄²⁻] \]
\[ Q = (0.10 \text{ M}) \times (0.001 \text{ M}) \]
\[ Q = 1.0 \times 10^{-4} \]
Compare Q with the given Ksp:
\[ Q = 1.0 \times 10^{-4} \]
\[ Ksp = 1.4 \times 10^{-5} \]
Since \( Q > Ksp \), the solution is supersaturated, and a precipitate will form. Therefore, the correct option is:
- ○ no, because Q > Ksp
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