Problem: You add 69.4 mL of a 0.447 AgNO3 solution to solution to 119.5 mL of a 0.128 M NaCl solution. Determine the concentration of Ag+ and Cl- ions in solution as if all remain dissolved. After, use your Ksp expression to determine if a precipitate will form. Which of the following represents what will occur and why? A) A precipitate will form, because Q > KB) A precipitation will form, because Q < KC) A precipitation will not form, because Q > KD) A precipitation will not form, because Q < K 

We are being asked to determine whether a precipitate will form when you add 69.4 mL of a 0.447 AgNO₃ solution to solution to 119.5 mL of a 0.128 M NaCl solution

We will do the following steps:

Step1: Write the net ionic equation

Step 2: Calculate initial concentrations

Step 3: Calculate Q

Step 4: Compare Q and K

$\boxed{\mathbf{Q}\mathbf{ }\mathbf{=}\mathbf{ }\frac{\mathbf{product}}{\mathbf{reactant}}}$

Depending on if Q is greater than or less than K, our reaction will shift to attain equilibrium by reaching the equilibrium constant K:

If Q = K → the reaction is at equilibrium

If Q < K → the reaction shifts in the forward direction to reach equilibrium

If Q > K → the reaction shifts in the reverse direction to reach equilibrium