<p>A solution contains 0.018 moles each of I-, Br-, and Cl-.
When the solution is mixed with 200 mL of 0.24 M AgNO3, how much AgCl(s)
precipitates out, and what is [Ag+]? Assume no volume change. </p>
<p>Ksp AgI = 1.5 * 10-16
Ksp AgBr = 5.0 * 10-13
Ksp AgCl = 1.6 * 10-10 </p>
<p>Here is my work so far.....</p>
<pre><code> AgI --> Ag+ + I-
</code></pre>
<p>Before 0.24M 0.09M<br>
After 0.15M 0M </p>
<p>Equilibrium 0.15 + x x</p>
<p>(0.15 + x) x = 1.5 * 10-16 change is neglible, so [Ag+] = 0.15 M </p>
<p>AgI --> Ag+ + Br-
Before 0.15M 0.09M<br>
After 0.06M 0M </p>
<p>At equilibrium, change is neglible. </p>
<p>AgI --> Ag+ + Cl-
Before 0.06M 0.09M<br>
After 0M 0.03M
Equilibrium x 0.03 + x</p>
<p>x(0.03 + x) = 1.6 * 10-10
0.03x = 1.6 x 10-10
x = 5.333 x 10-9 </p>
<p>[Ag+] = 5.333 x 10-9 M</p>
<p>So it the above answer [Ag+] ions remaining in solution?
How to find out how much AgCl precipitate? </p>
<p>Thank you in advance</p>