A solution contains 0.020 mole each of I¯, Br¯, and Cl¯, When the solution is mixed with 200. mL of 0.29 M AgNO3, what mass of AgCI(s) precipitates out, and what is [Ag*]? Assume no volume change. AgI: Ksp = 1.5 x 10¬16 AgBr: Kp = 5.0 × 10¬13 AgCl: К,р — 1.6 х 10-10 Mass = g (Ag*]=| M
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![**Chemical Precipitation Problem**
A solution contains 0.020 mole each of \( \text{I}^- \), \( \text{Br}^- \), and \( \text{Cl}^- \). When the solution is mixed with 200. mL of 0.29 M \( \text{AgNO}_3 \), what mass of \( \text{AgCl(s)} \) precipitates out, and what is \([\text{Ag}^+]\)?
**Given Solubility Product Constants (Ksp):**
- \( \text{AgI: } \quad K_{\text{sp}} = 1.5 \times 10^{-16} \)
- \( \text{AgBr: } \quad K_{\text{sp}} = 5.0 \times 10^{-13} \)
- \( \text{AgCl: } \quad K_{\text{sp}} = 1.6 \times 10^{-10} \)
**Questions:**
- Mass = \(\boxed{} \, \text{g} \)
- \([\text{Ag}^+ ] = \boxed{} \, \text{M} \)](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F3216a391-9091-4165-86c6-f10847a236cf%2F273ba84a-30ca-4f3b-a766-18112db28117%2F4e358ha_processed.png&w=3840&q=75)
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