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19 Nov 2019

Given that the average concentration of calcium ion in natural waters is 3.8 x 10-4 M, and that the solubility product Ksp for CaF2 is 4 x 10-11, calculate the maximum molar concentration of fluoride ion that can be dissolved before precipitation begins. Show that this molarity is equivalent to about 6 ppm. (Hint: Write the solubility product expression for CaF2 and substitute the molar value for [Ca2+] into it.)

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Beverley Smith
Beverley SmithLv2
10 Aug 2019

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Related questions

Consider a saturated solution of calcium fluoride in 0.092 M potassium nitrate. Complete the following tasks, and then answer the question.a) Write the chemical equation corresponding to Ksp.b) Write the defining expression for the thermodynamic solubility product constant, Kspo.c) Write the defining expression for the conditional solubility product constant, Ksp'.d) For the Ca2+ and F- ions, write the mathematical relations defining activity of each ion in terms of its concentration and activity coefficient.e) Based on your answers above, derive a relationship for Ksp' in terms of Kspo and activity coefficients, and calculate the value of Ksp', given that Kspo = 3.9 x 10-11. (Note: Compute the activity coefficient for each ion using the Debye-Huckel equation. You can disregard the contribution of the dissolved CaF2 to the overall ionic strength. Also disregard the formation of any HF in the equilibrium solution.) Question:Calculate the molar solubility of the calcium fluoride.

B2. A chemist mixes equal volumes of an aqueous solution of CaCl2 (1.34  10–4 M ) and an aqueous solution of Na2CO3 (1.34  10–4 M) to prepare a saturated solution of calcium carbonate (no precipitate is observed).

(a) Calculate the concentration of each of the ions in this saturated solution (assuming 100% dissociation).

[Ca2+] = [Na+] = [Cl–] = [CO32–] =

(b) Assuming that the chemist has calculated the molar solubility for calcium carbonate accurately, write the expression for the solubility product and calculate the value for Ksp for calcium carbonate.

1.In gravimetric analysis, it is possible to use the common-ion effect to favor the production of solid precipitate. In this problem you will test this approach to see how effective it is for the gravimetric determination of calcium. Suppose the reaction flask in the gravimetric analysis experiment contains 175 mL of solution before filtering through the Gooch crucible. A precipitate of calcium oxalate monohydrate is collected in the crucible and dried, where:

CaC2O4 * H2O <--> Ca2^+ + C2O4^2- + H2O Ksp=1.3x10^-8

The measured dry mass of CaC2O4·H2O precipitate is 0.4324 g. In this problem you will determine the mass of Ca2+ that remains in the filtrate and is therefore unaccounted for in the precipitate. Specifically, please do the following:

(a) Determine the moles of oxalate ion in the 175-mL reaction flask before precipitation. Use the experimental information in the lab manual. For the purposes of this problem, assume the oxalate ion is totally deprotonated.

(b) Determine the equilibrium molarity of dissolved Ca2+ after the precipitation has occurred. [Hint: assume the precipitation first goes to completion; i.e., all the calcium ion from the original sample is in the precipitate. Then set up an ICE table to determine the amount (in moles/L) of the calcium ion that goes back into solution via solubility equilibrium subsequent to the precipitation. Note carefully that that the oxalate (“common ion”) is in excess before the precipitation.]

(c) Determine the mass (in g) of dissolved Ca2+ that remains in solution following filtration.

(d) Determine the mass (in g) of calcium ion in the solid precipitate.

(e) From your answers to (c) and (d), determine what percentage of the original calcium mass is lost to the filtrate.

(f) Repeat parts (b) through (e) under conditions in which the oxalate is not in excess; in other words, set the initial oxalate concentration in your ICE table to zero.

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