Question

Which of the following compounds precipitates from a solution that has the concentrations indicated? (See Appendix J for Ksp values.)

a) CaCO3 at [Ca2+] = 1.5×10−3M and [CO3^2-] = 2.0×10−3M
b) AgCl at [Ag+] = 8.0×10−9M and [Cl-] = 1.0×10−3M
c) Fe(OH)2 at [Fe2+] = 4.0×10−5M and [OH-] = 2.0×10−5M
d) BaSO3 at [Ba2+] = 2.5×10−4M and [SO3^2-] = 1.0×10−3M

Answer 1

To determine if a compound precipitates, calculate the ion product (Q) and compare it to the solubility product constant (Ksp). If Q > Ksp, precipitation occurs. For CaCO3, Q = (1.5×10–3)(2.0×10–3) = 3.0×10–6, and this should be compared to Ksp to assess precipitation.

Step-by-step explanation:

To determine whether a compound precipitates from a solution, we compare the ion product (Q) of the concentrations of the ions in the solution to the solubility product constant (Ksp) for the compound. If Q is greater than Ksp, the solution is supersaturated, and precipitation will occur to restore equilibrium.

For example, let's calculate the ion product for CaCO3 using the given concentrations:

Q = [Ca2+][CO32-–] = (1.5×10–3M)(2.0×10–3M) = 3.0×10–6M2

We would then compare this value to the Ksp of CaCO3 provided in Appendix J to determine whether precipitation will occur.