Chapter 18 Solubility. Equilibria of Slightly Soluble Ionic Compounds Explore the aqueous equilibria of slightly soluble ionic compounds. Chapter 5. Precipitation.

Chapter 18 Solubility

Equilibria of Slightly Soluble Ionic Compounds

• Explore the aqueous equilibria of slightly soluble ionic compounds.

• Chapter 5. Precipitation Reactions:

AgNO3(aq) + NaCl(aq) ??

• Want to consider quantitative predictions

The Formation & Dissolution of Precipitates

• Solubility: maximum amount of solute that will dissolve in a given amount of solvent (depends on solvent, temperature and pressure)

No compound is infinitely soluble andno compound is perfectly insoluble.

Solute Solubility

(g solute/100 g solvent)

Qualitative Solubility

Description

Less than 0.1 Insoluble

0.1 – 1 Slightly soluble

1 – 10 Soluble

Greater than 10 Very soluble

The Formation & Dissolution of Precipitates

• Saturated solution: contains maximum concentration of solute– Equilibrium between undissolved and dissolved solute.

• Solutes (even those called “soluble”) have a limited solubility in a particular solvent.

• Slightly soluble (often called “insoluble”) ionic compounds have a relatively low solubility – Reach equilibrium with little solute dissolved– Heterogeneous equilibrium

Why is this important?

Dissolving and Precipitation occurs around us:– Tooth enamel dissolves in acidic soln (tooth decay)

– Ppt of certain salts in kidneys causes kidney stones

– Waters of Earth contains dissolved salts as water passes over and through the ground

– Ppt of CaCO3 from groundwater is responsible for cave formation.

Let’s look at the factors that affect solubility!

Solubility-Product Constant (Ksp)

• Solubility-product constant (Ksp): equilibrium constant for equilibrium between slightly soluble ionic solid and a solution of its ions– Indicates how soluble the solid is in water

• Solubility: quantity that dissolves to form a saturated solution (g/L)

• Molar solubility: number of moles of solute that dissolves in forming a liter of saturated solution of solute (mol/L)

• Solubility depends on concentrations of other ions and pH but Ksp is a constant.

Solubility-Product Constant (Ksp)

Practice: Write an ionic equation for the dissolution, and the equation for the solubility product for:

(a) Calcium carbonate

(b) Magnesium hydroxide

(c) Ag3PO4

• Magnitude of Ksp is measure of how far to the right dissolution proceeds at equilibrium (saturation).

– Used to compare solubilities if same total number of ions

Ksp of Selected Ionic Compounds (25 °C)

Name, Formula Ksp

Aluminum hydroxide, Al(OH)3 3 x 10-34

Cobalt(II) carbonate, CoCO3 1.0 x 10-10

Iron(II) hydroxide, Fe(OH)2 4.1 x 10-15

Lead(II) fluoride, PbF2 3.6 x 10-8

Lead(II) sulfate, PbSO4 1.6 x 10-8

Mercury(I) iodide, Hg2I2 4.7 x 10-29

Silver sulfide, Ag2S 8 x 10-48

Zinc iodate, Zn(IO3)2 3.9 x 10-6

See Appendix D in your book for a much more extensive list.

Example

Predict which of the following compounds will

have the greatest molar solubility in water

A) AgCl Ksp = 1.8 x 10-10

B) AgBr Ksp = 5.0 x 10-13

C) AgI Ksp = 8.3 x 10-17

D) all have the same molar solubility

Solubilities and Solubility Products

• Ksp for a slightly soluble solid can be determined from its solubility – as long as there is no other reaction

• Example 1: Lead(II) sulfate (PbSO4) is a key component in lead-acid car batteries. Its solubility in water at 25 °C is 4.25 x 10-3 g/100 mL solution. What is the Ksp of PbSO4?

• Example 2: Determine the molar solubility of MgF2 from its solubility product (Ksp = 6.4 x 10-9).

Example

Calculate the molar solubility of calcium

fluoride, CaF2

Ksp = 3.7 x 10-11

Factors that Affect Solubility: Common Ion Effect

•The presence of a common ion decreases the solubility of a slightly soluble ionic compound.

•The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction.

AgCl(s) ⇌ Ag+(aq) + Cl(aq)

adding NaCl( ) shifts equilibrium positionaq

The effect of a common ion on solubility

PbCrO4(s) Pb2+(aq) + CrO42-(aq) PbCrO4(s) Pb2+(aq) + CrO4

2-(aq)

CrO42- added

Example

The solubility of Ca(OH)2 in water is 0.012 M.

What is its solubility in 0.10 M Ca(NO3)2?

Ksp of Ca(OH)2 is 6.5 x 10-6

Le Châtelier’s Principle

Determine the effects of solubility when each of

the following is added to a mixture of the

slightly soluble solid NiCO3 and water at

equilibrium:

(a)Ni(NO3)2 (c) K2CO3

(b)KClO4 (d) HNO3

Effect of pH on solubility

• [H3O+] can have a profound effect on the solubility of an ionic compound.

• Solubility of slightly soluble salts containing basic anions increases as [H+] increases– More basic anion…more solubility is influenced by pH

• Predict the effect on solubility of adding a strong acid

CaCO3(s) ⇌ Ca2+(aq) + CO3

2-(aq)

AgCl(s) ⇌ Ag+(aq) + Cl-

(aq)

Ca(OH)2(s) ⇌ Ca2+(aq) + 2OH-

(aq)

If limestone (CaCO3) deposit is near surface…sinkhole

If limestone (CaCO3) deposit is well below the surface….caves

A view inside Carlsbad Caverns, New Mexico

Copyright © Houghton Mifflin Company.All rights reserved. 1–19

Cango Caves

Copyright © Houghton Mifflin Company.All rights reserved. 1–20

Sudwala Caves

Example

Calculate the molar solubility of MgF2 in 0.10 M

MgCl2 at 25 C.

Ksp of MgF2 = 7.4 x 10-11

Predicting the Formation of a Precipitate

Compare Qsp to Ksp to predict if a precipitate will form

and, if not, what concentrations of ions will cause it to

do so.

Qsp = Ksp soln is saturated & no changes occur

Qsp > Ksp ppt forms until soln is saturated

Qsp < Ksp soln is unsaturated & no ppt forms

Practice

• Example 1: Determine whether CaHPO4 will precipitate from a solution with [Ca2+] = 0.0001 M and [HPO4

2-] = 0.001 M.

• Example 2: Does silver chloride precipitate when equal volumes of a 2 x 10-4 M solution of AgNO3 and a 2 x 10-4 M solution of NaCl are mixed.

AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)

Practice

Will a precipitate form when 0.10 L of

8.0 x 10-3 M Pb(NO3)2 is added to 0.40 L of

5.0 x 10-3 M Na2SO4?

Ksp for PbSO4 = 6.3 x 10-7

Concentration Necessary to Form a Ppt

• We can also determine the concentration of an ion necessary for precipitation to begin.

• Assume that precipitation begins when Qsp = Ksp

• Example: If a solution contains 0.0020 mol CrO42-

per liter, what concentration of Ag+ ion must be added as AgNO3 before Ag2CrO4 begins to precipitate. (Neglect any increase in volume upon adding the solid silver nitrate.)