The molar concentration of barium ions in a saturated aqueous solution of barium fluoride was measured to be 7.5 x 103 M at 25C. From this...

The molar concentration of barium ions in a saturated aqueous solution of barium fluoride was measured to be 7.5 x 10−3 M at 25°C. From this information, estimate the Ksp of barium fluoride at 25°C. (Disregard any hydrolysis associated with the fluoride ion).

1.7 x 10−6

8.8 x 10−6

4.3 x 10−11

5.6 x 10−5

If the theoretical molar solubility of CaCO3 in pure water at 25°C is 5.7 x 10−5 M, estimate the theoretical molar solubility of CaCO3 in "hard water" containing 3.0 x10−3 M Ca2+ ions. (Hint: what is the Ksp of CaCO3 at 25°C ?)

1.1 x10−6 M

3.2 x10−6 M

1.7 x10−7 M

2.2 x10−8 M

A saturated aqueous solution of an unknown semi-soluble hydroxide salt M(OH)2 exhibits a pH of 12.91 at 25°C. Which of the following is the most likely identity of the hydroxide salt?

Ni(OH)2 (Ksp = 5.3 x 10−16)

Zn(OH)2 (Ksp = 3.0 x 10−17)

Cd(OH)2 (Ksp = 7.4 x 10−15)

Ba(OH)2 (Ksp = 2.7 x 10−4)

For which of the following salts do you predict the experimental solubility (as determined in a lab) to exceed the theoretical solubility (as calculated by the Ksp of the salt)?

Ca3(PO4)2

Fe(OH)2

PbCl2

AgBr

Which of the following sodium salt solutions would you choose to selectively precipitate X2+ ions from a solution containing equal concentrations of X2+ and Y2+ ions?

Potentially helpful Ksp information:  

XSO4, Ksp = 1 x 10−7

YSO4, Ksp = 2 x 10−9

X(OH)2, Ksp = 1 x 10−10

Y(OH)2, Ksp = 2 x 10−5

X(IO3)2, Ksp = 2 x 10−9

Y(IO3)2, Ksp = 1 x 10−9

XCl2, Ksp = 1 x 10−2

YCl2, Ksp = 2 x 10−6

a solution of NaOH

a solution of Na2SO4

a solution of NaIO3

a solution of NaCl

If 50.0 mL of a 0.0040 M Pb(NO3)2 solution is mixed with 50.0 mL of a 0.0060 M NaI solution at 25°C, would a precipitation of PbI2 form? (Ksp of PbI2 = 7.9 x 10−9 at 25°C)

yes because Q > Ksp

no because Q < Ksp

yes because Q < Ksp

no because Q > Ksp

Which of the following statements is FALSE?

AgCl is predicted to be more soluble in 0.10 M HCN than in pure water (Kf of Ag(CN)2− = 3 x 1020)

A saturated aqueous solution of AgCl is predicted to exhibit an approximately neutral pH at 25°C

AgCl is predicted to be more soluble in pure water than in 0.10 M HCl

Ag2CO3 is predicted to be more soluble in pure water than in 0.10 M HCl

Which of the following best explains why a semisoluble hydroxide salt like Zn(OH)2 is more soluble in acidic or basic aqueous solutions than in pure water?

Because hydroxide ion OH- can accept a proton from an acid (to form H2O) and donate a proton to a base (to form O2-) and both of these processes are equally potential in aqueous solutions. 

Because the Zn2+ ion can be both a Lewis acid and a Lewis base and can therefore react with both bases and acids.

Because the actual chemical formula for zinc hydroxide is not merely Zn(OH)2 but actually Zn(H2O)2(OH)2 with acidic O—H bonds in the H2O molecules around the Zn2+ ion, as well as the basic OH− hydroxides. Hence the amphoteric nature of such hydroxide salts as "Zn(OH)2", which can therefore react with acids or with bases.

Consider the following solubility equilibrium at 25°C:

M2X(s) ↔ 2M+(aq) + X2−(aq)       Ksp < 1.0  

Which of the following statements is FALSE regarding this equilibrium?

2[M+]eq  = [X2−]eq

At 25°C, [X2−]eq  <  1.0 M

Q = Ksp

If the reaction is endothermic, an increase in temperature should increase the solubility of M2X

Which of the following describes an "ideal solution"?

An ideal solution exhibits an exothermic ΔHsolution

An ideal solution exhibits a neutral pH

An ideal solution exhibits whole-number van't Hoff factors

An ideal solution obeys PV=nRT