Given the following solubility constants, which list arranges the solutes in order of increasing solubility?
CaCO3: Ksp = 2.8 × 10-9
Ca(OH)2: Ksp = 5.5 × 10-6
CaSO4: Ksp = 9.1 × 10-6
CaF2: Ksp = 5.3 × 10-9
A. CaSO4 < Ca(OH)2 < CaCO3 < CaF2
B. CaSO4 < Ca(OH)2 < CaF2 < CaCO3
C. CaCO3 < CaF2 < Ca(OH)2 < CaSO4
D. CaF2 < CaCO3 < Ca(OH)2 < CaSO4
It's C
To determine the order of increasing solubility, we need to compare the solubility product constants (Ksp) of the given solutes.
The solute with the lower Ksp value is less soluble than the one with the higher Ksp value.
Comparing the Ksp values given:
CaCO3: Ksp = 2.8 × 10^-9
Ca(OH)2: Ksp = 5.5 × 10^-6
CaSO4: Ksp = 9.1 × 10^-6
CaF2: Ksp = 5.3 × 10^-9
Comparing the Ksp values, we can see that CaF2 has the lowest solubility, followed by CaCO3, Ca(OH)2, and CaSO4.
Therefore, the correct order of increasing solubility is:
CaF2 < CaCO3 < Ca(OH)2 < CaSO4
So, the correct answer is option D.
To determine the order of solubility for the given solutes, we need to compare their solubility constants (Ksp) values. The solute with the highest Ksp will be the most soluble, while the solute with the lowest Ksp will be the least soluble.
Comparing the given solubility constants:
CaCO3: Ksp = 2.8 × 10^-9
Ca(OH)2: Ksp = 5.5 × 10^-6
CaSO4: Ksp = 9.1 × 10^-6
CaF2: Ksp = 5.3 × 10^-9
From the given solubility constants, we can see that CaSO4 has the highest Ksp, followed by Ca(OH)2, CaF2, and CaCO3 which has the lowest Ksp. Therefore, the correct order of increasing solubility is CaCO3 < CaF2 < Ca(OH)2 < CaSO4.
Hence, the correct option is C. CaCO3 < CaF2 < Ca(OH)2 < CaSO4.