What concentration of SO32– is in equilibrium with Ag2SO3(s) and 3.60 × 10-3 M Ag ? The Ksp is 8.46×10–12.
To find the concentration of \( \text{SO}_3^{2-} \) in equilibrium with Ag2SO3(s) and 3.60 × 10-3 M Ag, we need to use the solubility product constant (Ksp).
The balanced equation for the dissolution of Ag2SO3 is:
\[ Ag2SO3(s) \rightleftharpoons 2Ag^+(aq) + SO3^{2-}(aq) \]
The Ksp expression for this reaction is:
\[ Ksp = [Ag^+]^2 \cdot [SO3^{2-}] \]
Where \( [Ag^+] \) represents the concentration of Ag+ ions in the solution, and \( [SO3^{2-}] \) represents the concentration of sulfate ions in the solution.
Given that the concentration of Ag+ ions is 3.60 × 10-3 M, we can substitute these values into the Ksp expression:
\[ 8.46 × 10^{-12} = (3.60 × 10^{-3})^2 \cdot [SO3^{2-}] \]
Now, we can solve for the concentration of \( [SO3^{2-}] \):
\[ [SO3^{2-}] = \frac{8.46 × 10^{-12}}{(3.60 × 10^{-3})^2} \]
Calculating this expression will give us the concentration of \( \text{SO}_3^{2-} \) in equilibrium with Ag2SO3(s) and 3.60 × 10-3 M Ag.