Chapter Sixteen - University of Houstonchembi/chem1332ch16.pdfChapter Sixteen 29 • What is the...

Prentice Hall Hall ©© 20052005General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry

Chapter Sixteen

1

More Equilibria in Aqueous Solutions:

Slightly Soluble Salts and Complex Ions

Chapter SixteenChapter Sixteen


Chapter Sixteen

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• Solubility product constant, Ksp: the equilibrium constant expression for the dissolving of a slightly soluble solid.

The Solubility Product Constant, Ksp

Ksp = [ Ag+ ][ Cl–]

AgCl(s) Ag+(aq) + Cl–(aq)

• Many important ionic compounds are only slightly soluble in water (we used to call them “insoluble” – Chapter 4). E.g. Ag salts, sulfides

• An equation can represent the equilibrium between the compound and the ions present in a saturated aqueous solution:


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Ksp and Molar SolubilityThe solubility product constant is related to the solubility of an ionic solute

Ksp = [Ag+][Cl–]; solubility given by [Ag+]

From stoichiometry, the ion ratio is 1:1, so [Ag+] = [Cl–], both of which are unknown (x)

EOS

Ksp = x2 and [Ag+] = (Ksp)1/2

Ag+ Cl– Ag+ Cl–+


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Generally the smaller the Ksp, the less soluble the salt!

But… check the number of ions per formula unit.


Chapter Sixteen

6Instead of “x” use s to indicate solubility.Solubility is per mole of salt, not the molarity of

each dissolved ion!Two ions per formula – s2

Three ions per formula – (2s)2s = 4s3


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The Common Ion EffectLe Châtelier’s principle is followed for the shift in concentration of products and reactants upon addition of either products or reactants to a solution

The solubility of a slightly soluble ionic compound is lowered when a second solute that furnishes a common ion is added to the solution

Solubility of Ag2SO4 ↓if MgSO4 is added to solution

EOS

Ag2SO4(s) 2 Ag+(aq) + SO4–2(aq)


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Common Ion Effect Illustrated

Na2SO4(aq)

SaturatedAg2SO4(aq)

Ag2SO4precipitates

The added sulfate ion reduces the

solubility of Ag2SO4.


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Example 16.3From the Ksp value for silver sulfate, calculate its molar solubility at 25 °C.

Ag2SO4(s) 2 Ag+(aq) + SO42–(aq)

Ksp = 1.4 x 10–5 at 25 °C


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Example 16.5Calculate the molar solubility of Ag2SO4 in 1.00 M Na2SO4(aq).


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• Ions that are not common to the precipitate can also affect solubility.– CaF2 is more soluble in 0.010 M Na2SO4 than it is in

water.• Increased solubility occurs because of interionic

attractions.• Each Ca2+ and F– is surrounded by ions of

opposite charge, which impede the reaction of Ca2+ with F–.

• The effective concentrations, or activities, of Ca2+

and F– are lower than their actual concentrations.

Solubility and Activities


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• Qip can then be compared to Ksp.• Precipitation should occur if Qip > Ksp.• Precipitation cannot occur if Qip < Ksp.• A solution is just saturated if Qip = Ksp.• In applying the precipitation criteria, the

effect of dilution when solutions are mixed must be considered.

• Qip is the ion product reaction quotient and is based on initial conditions of the reaction.

Will Precipitation Occur? Is It Complete?


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Other Things To Remember About Ksp

• If the anion or cation of the salt can react with H3O+ or OH-, the solubility of the salt will increase due to the formation of weak acids or bases.

• Q > Ksp saturated solution• Q < Ksp unsaturated solution (no precipitation)• Q = Ksp precipitation will begin, additional salt

– Will cause additional precipitation. Represents maximum amount of cation and anion concentration allowed at that temperature


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• The solubility of Ag2CO3 is 1.26 x 10-4 at 25oC. Calculate the Ksp for silver carbonate.

• Ag2CO3(s) 2 Ag+ + CO32-

• Initial 0 0• Change +2s +s• Equil 2s s• Ksp = [Ag+]2[CO3

2-] = (2s)2(s) = 4s3

• Ksp = (4)(1.26 x 10-4)3 = 8.00 x 10-12

• If MgF2 has a Ksp = 5.2 x 10-11 what is the solubility of this salt?


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• If 100mL of 0.0200M BaCl2 is mixed with 50.0mL of 0.0300M Na2SO4 will BaSO4precipitate? Ksp = 1.1 x 10-10

• 1) First calculate the concentrations of each ion.• M2 = M1V1/V2

• Ba2+ =(0.0200)(100mL)/(150mL) = 0.0133M • For SO4

2-= (0.0300M)(50.0mL)/(150mL) = 0.0100M

• Q = [Ba2+][SO42-] = (0.0133M)(0.0100M) = 1.33

x 10-4

• Will precipitation happen?


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• We generally consider precipitation to be “complete” if about 99.9% of the target ion is precipitated (0.1% or less left in solution).

• Three conditions generally favor completeness of precipitation:

1. A very small value of Ksp. (few ions in solution)2. A high initial concentration of the target ion.3. A concentration of common ion that greatly exceeds that

of the target ion. (use soluble salts, e.g. alkali metal salts)

To Determine Whether Precipitation Is Complete


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Example 16.9To a solution with [Ca2+] = 0.0050 M, we add sufficient solid ammonium oxalate, (NH4)2C2O4(s), to make the initial [C2O4

2–] = 0.0051 M. Will precipitation of Ca2+ as CaC2O4(s) be complete?

CaC2O4(s) Ca2+(aq) + C2O42–(aq) Ksp = 2.7 x 10–9


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Example 16.4 A Conceptual ExampleWithout doing detailed calculations, but using data from Table 16.1, establish the order of increasing solubility of these silver halides in water: AgCl, AgBr, AgI.


Chapter Sixteen

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Chapter Sixteen

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Selective PrecipitationAgNO3 added to

a mixture containing Cl–

and I–


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Effect of pH on SolubilityThe solubility of an ionic solute may be greatly affected by pH if an acid–base reaction also occurs as the solute dissolves

Precipitate anions of weak acids will dissolve somewhat when the pH is lowered; precipitate anions of strong acids are not affected by pH

EOS

Cl– is the conjugate base of HCl, a strong acid – pH doesn’t affect the solubility of AgCl


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• If the anion of a precipitate is that of a weak acid, the precipitate will dissolve somewhat when the pH is lowered:

• If, however, the anion of the precipitate is that of a strong acid, lowering the pH will have no effect on the precipitate.

Added H+ reacts with, and removes, F–; LeChâtelier’s

principle says more F– forms.

H+ does not consume Cl– ; acid does not affect the

equilibrium.

Effect of pH on Solubility

CaF2(s) Ca2+(aq) + 2 F–(aq)

AgCl(s) Ag+(aq) + Cl–(aq)


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Equilibria Involving Complex Ions

Silver chloride becomes more soluble, not less soluble, in high concentrations of chloride ion.


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• A complex ion consists of a central metal atom or ion, with other groups called ligands bonded to it.

• The metal ion acts as a Lewis acid (accepts electron pairs).

• Ligands act as Lewis bases (donate electron pairs).

• The equilibrium involving a complex ion, the metal ion, and the ligands may be described through a formation constant, Kf:

Complex Ion Formation

Ag+(aq) + 2 Cl–(aq) [AgCl2]–(aq)

[AgCl2]–

Kf = –––––––––– = 1.2 x 108

[Ag+][Cl–]2

Common ligands


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Complex Ion FormationConcentrated NH3

added to a solution of pale-blue Cu2+ …

… forms deep-blue Cu(NH3)4

2+.


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Chapter Sixteen

27Complex Ion Formation

and Solubilities

AgCl is insoluble in water.

But if the concentration of NH3 is made high

enough …

… the AgCl forms the soluble

[Ag(NH3)2]+ ion.

Even though AgCl is insoluble in water, it does dissolve in aqueous ammonia to form the complex ion [Ag(NH3)2]+


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Kf = formation constantKd = dissociation constant

Kd = 1/Kf

If a complexing agent is present, the solubility of an

ion will be greatly enhanced!


Chapter Sixteen

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• What is the value of the equilibrium constant if AgBr is dissolved in a solution of thiosulfateinstead of pure water. The equilibria are

• AgBr(s) Ag+ + Br- Ksp = 5.4 x 10-14

• Ag+ + 2 S2O32- Ag(S2O3)2

3- Kf = 2.0 x 1013

• To get the new K value simply add the equations and do what to the K values?


Chapter Sixteen

30Example 16.13Calculate the concentration of free silver ion, [Ag+], in an aqueous solution prepared as 0.10 M AgNO3 and 3.0 M NH3.

Ag+(aq) + 2 NH3(aq) [Ag(NH3)2]+(aq) Kf = 1.6 x 107


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Example 16.15What is the molar solubility of AgBr(s) in 3.0 M NH3?

AgBr(s) + 2 NH3(aq) [Ag(NH3)2]+(aq) + Br–(aq)

Kc = 8.0 x 10–6


Chapter Sixteen

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• Water molecules are commonly found as ligands in complex ions (H2O is a Lewis base).

• The electron-withdrawing power of a small, highly charged metal ion can weaken an O—H bond in one of the ligand water molecules.

• The weakened O—H bond can then give up its proton to another water molecule in the solution (hydrolysis).

• The complex ion acts as an acid.

Complex Ions in Acid–Base Reactions

[Na(H2O)4]+ [Al(H2O)6]3+ [Fe(H2O)6]3+


Chapter Sixteen

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Ionization of a Complex Ion

[Fe(H2O)6]3+ + H2O [Fe(H2O)5OH]2+ + H3O+

The highly-charged iron(III) ion withdraws electron density

from the O—H bonds.

Ka = 1 x 10–7


Chapter Sixteen

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• Certain metal hydroxides, insoluble in water, are amphotericamphoteric; they will react with both strong acids and strong bases.

• Al(OH)3, Zn(OH)2, and Cr(OH)3 are amphoteric.

Amphoteric Species


Chapter Sixteen

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Amphoteric Substances

The hydroxides of Al, Zn, and Cr are amphoteric

Like hydroxides, oxides that can react with and dissolve in either an acid or a base are often said to be amphoteric

EOS

The oxides of Al, Zn, and Cr behave toward acids and bases much as their hydroxides do


Chapter Sixteen

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• Acid–base chemistry, precipitation reactions, oxidation–reduction, and complex ion formation all apply to an area of analytical chemistry called classical qualitative inorganic analysis.

• “Qualitative” signifies that the interest is in determining what is present.– Quantitative analyses are those that determine how much of a

particular substance or species is present.• Although classical qualitative analysis is not used as

widely today as are instrumental methods, it is still a good vehicle for applying all the basic concepts of equilibria in aqueous solutions. -

• Applies knowledge of color changes, precipitates, and equilibrium as a determination scheme.

Qualitative Inorganic Analysis


Chapter Sixteen

37Qualitative Analysis Outline

In acid, H2S produces very little S2–, so only

the most-insoluble sulfides precipitate.

In base, there is more S2–, and the less-insoluble

sulfides also precipitate.

Some hydroxides also precipitate here.


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Group 1 CationsIf aqueous HCl is added to an unknown solution of cations, and a precipitate forms, then the unknown contains one or more of these cations: Pb2+, Hg2

2+, or Ag+ - These are the only ions to form insoluble chlorides.

The supernatant liquid and any precipitate are saved for further analysis. If there is no precipitate, then Group 1 ions must be absent from the mixture.

Test for Hg2

+2

EOS

Test for Pb+2

Any precipitate is separated from the mixture and further tests are performed to determine which of the three Group 1 cations are present.


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• Precipitated PbCl2 is slightly soluble in hot water.• The precipitate is washed with hot water, then aqueous

K2CrO4 is added to the washings.• If Pb2+ is present, a precipitate of yellow lead chromate

forms, which is less soluble than PbCl2.• (If all of the precipitate dissolves in the hot water, what

does that mean?)

Cation Group 1 (cont’d)

Analyzing for Pb2+


Chapter Sixteen

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• Next, any undissolved precipitate is treated with aqueous ammonia.

• If AgCl is present, it will dissolve, forming Ag(NH3)2+ (the

dissolution may not be visually apparent).• If Hg2

2+ is present, the precipitate will turn dark gray/ black, due to a disproportionation reaction that forms Hg metal and HgNH2Cl.

• The supernatant liquid (which contains the Ag+, if present) is then treated with aqueous nitric acid.

• If a precipitate reforms, then Ag+ was present in the solution.

Cation Group 1 (cont’d)

Analyzing for Ag+ and Hg22+


Chapter Sixteen

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PbCl2 precipitates when HCl is added.

Hg2Cl2 reacts with NH3 to form black Hg metal and HgNH2Cl.

The presence of lead is confirmed by adding chromate ion; yellow PbCrO4 precipitates.

Group 1 Cation Precipitates


Chapter Sixteen

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• Once the Group 1 cations have been precipitated, hydrogen sulfide is used as the next reagent in the qualitative analysis scheme.

• H2S is a weak diprotic acid; there is very little ionization of the HS– ion and it is the precipitating agent.

• Hydrogen sulfide has the familiar rotten egg odor that is very noticeable around volcanic areas.

• Because of its toxicity, H2S is generally produced only in small quantities and directly in the solution where it is to be used.

Hydrogen Sulfide in theQualitative Analysis Scheme


Chapter Sixteen

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• The concentration of HS– is so low in a strongly acidic solution, that only the most insoluble sulfides precipitate.

• These include the eight metal sulfides of Group 2.• Five of the Group 3 cations form sulfides that are soluble

in acidic solution but insoluble in alkaline NH3/NH4+.

• The other three Group 3 cations form insoluble hydroxides in the alkaline solution.

• The cations of Groups 4 and 5 are soluble.• Group 4 ions are precipitated as carbonates.• Group 5 does not precipitate; these must be determined by

flame test.

Cation Groups 2, 3, 4, and 5


Chapter Sixteen

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Group 2 Cations

H2S is added to the supernatant liquid – insoluble sulfide formation is the positive test

EOS

The concentration of HS– is so low in a strongly acidic solution (HCl) that only the most insoluble sulfides precipitate


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Group 3 Cations

EOS

The other three group 3 cationsform hydroxide precipitates in the alkaline solution

Of the eight cations in Group 3, five form sulfides that are soluble in acidic solution but insoluble in an alkaline ammonia/ammonium chloride buffer solution


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Group 4 and 5 CationsThe group 4 cations are precipitated as carbonates from a buffered alkaline solution

The cations of group 5 remain soluble in the presence of all common reagents

EOS

Within each of the qualitative analysis groups, additional reactions to dissolve group precipitates and to separate and selectively precipitate individual cations for identification and confirmation are used


Chapter Sixteen

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Summary of Concepts

• The solubility product constant, Ksp, represents equilibrium between a slightly soluble ionic compound and its ions in a saturated aqueous solution

• The common ion effect is responsible for the reduction in solubility of a slightly soluble ionic compound

EOS

• Precipitation is assumed to be complete if no more than 0.1% of the target ion remains in solution


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Summary of Concepts

• The solubilities of some slightly soluble compounds depends strongly on pH

• Certain solutes become more soluble in the presence of species that can serve as ligands in complex ions

• The ability of water as ligand molecules to donate protons accounts for the acidic character of some complex ions

EOS

• Precipitation, acid–base, redox, complex-ion formation, and amphoteric behavior are all used for the qualitative analysis of common cations

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