What is the approximate concentration of free Ag* ion at equilibrium when 1.85x102 mol silver nitrate is added to 1.00 L of solution that is 1.390 M in S2032. For [Ag(S203)2], Kf= 2.9x1013. [Ag'] = M
What is the approximate concentration of free Ag* ion at equilibrium when 1.85x102 mol silver nitrate is added to 1.00 L of solution that is 1.390 M in S2032. For [Ag(S203)2], Kf= 2.9x1013. [Ag'] = M
Chemistry
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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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![**Problem Statement:**
What is the approximate concentration of free Ag⁺ ion at equilibrium when 1.85×10⁻² mol silver nitrate is added to 1.00 L of solution that is 1.390 M in S₂O₃²⁻? For [Ag(S₂O₃)₂]³⁻, K_f = 2.9×10¹³.
\[ \text{[Ag⁺]} = \boxed{\rule{2cm}{0.15mm}} \, \text{M} \]
**Explanation:**
This problem is about determining the concentration of free silver ions (Ag⁺) in a solution at equilibrium. The setup involves adding a certain amount of silver nitrate to a solution containing thiosulfate ions (S₂O₃²⁻). The equilibrium constant (K_f) for the formation of the complex [Ag(S₂O₃)₂]³⁻ is provided.
To solve this problem, you would typically use equilibrium calculations that involve:
1. Writing the balanced chemical equation for the formation of the complex.
2. Setting up the expression for the equilibrium constant, K_f.
3. Using initial concentrations and stoichiometry to find the equilibrium concentrations.
Here, the provided data includes:
- Initial moles of silver nitrate: 1.85×10⁻² mol
- Volume of the solution: 1.00 L
- Initial concentration of S₂O₃²⁻: 1.390 M
- Formation constant (K_f): 2.9×10¹³
Using these values, you can calculate the approximate concentration of free Ag⁺ at equilibrium.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F445c8372-a397-41f1-88ec-d7a9c1d0e786%2Fd201e1a6-63f4-4fe7-b385-5d76fdae029c%2F7bgxbh_processed.jpeg&w=3840&q=75)
Transcribed Image Text:**Problem Statement:**
What is the approximate concentration of free Ag⁺ ion at equilibrium when 1.85×10⁻² mol silver nitrate is added to 1.00 L of solution that is 1.390 M in S₂O₃²⁻? For [Ag(S₂O₃)₂]³⁻, K_f = 2.9×10¹³.
\[ \text{[Ag⁺]} = \boxed{\rule{2cm}{0.15mm}} \, \text{M} \]
**Explanation:**
This problem is about determining the concentration of free silver ions (Ag⁺) in a solution at equilibrium. The setup involves adding a certain amount of silver nitrate to a solution containing thiosulfate ions (S₂O₃²⁻). The equilibrium constant (K_f) for the formation of the complex [Ag(S₂O₃)₂]³⁻ is provided.
To solve this problem, you would typically use equilibrium calculations that involve:
1. Writing the balanced chemical equation for the formation of the complex.
2. Setting up the expression for the equilibrium constant, K_f.
3. Using initial concentrations and stoichiometry to find the equilibrium concentrations.
Here, the provided data includes:
- Initial moles of silver nitrate: 1.85×10⁻² mol
- Volume of the solution: 1.00 L
- Initial concentration of S₂O₃²⁻: 1.390 M
- Formation constant (K_f): 2.9×10¹³
Using these values, you can calculate the approximate concentration of free Ag⁺ at equilibrium.
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