Chapter 17b Solubility

7/27/2019 Chapter 17b Solubility

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CH160 General Chemistry II

Lecture Presentation

Solubility Equilibria

Chapter 17


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Why Study Solubility Equilibria?

Many natural processes involve precipitation or dissolutionof salts. A few examples:

Dissolving of underground limestone deposits (CaCO3)forms caves

Note: Limestone is water insoluble (How can this be?)

Precipitation of limestone (CaCO3) forms stalactites andstalagmites in underground caverns

Precipitation of insoluble Ca3(PO4)2 and/or CaC2O4 in thekidneys forms kidney stones

Dissolving of tooth enamel, Ca5(PO4)3OH, leads to toothdecay (ouch!)

Precipitation of sodium urate, Na2C5H2N4O2, in jointsresults in gouty arthritis.


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Many chemical and industrial processes involveprecipitation or dissolution of salts. A few examples:

Production/synthesis of many inorganic compounds involves theirprecipitation reactions from aqueous solution

Separation of metals from their ores often involves dissolution

Qualitative analysis, i.e. identification of chemical species insolution, involves characteristic precipitation and dissolutionreactions of salts

Water treatment/purification often involves precipitation of metalsas insoluble inorganic salts

Toxic Pb2+, Hg2+, Cd2+ removed as their insoluble sulfide (S2-) saltsPO4

3- removed as insoluble calcium salts

Precipitation of gelatinous insoluble Al(OH)3 removes suspendedmatter in water


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To understand precipitation/dissolution processes in

nature, and how to exploit precipitation/dissolution

processes for useful purposes, we need to look at the

quantitative aspects of solubility and solubility

equilibria.


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Solubility of Ionic Compounds

Solubility Rulesgeneral rules for predicting the solubility of ionic

compounds

strictly qualitative


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Solubility Rule Examples

All alkali metal compounds are soluble

Most hydroxide compounds are insoluble. The

exceptions are the alkali metals, Ba2+, and Ca2+Most compounds containing chloride are soluble. The

exceptions are those with Ag+, Pb2+, and Hg22+

All chromates are insoluble, except those of the alkalimetals and the NH4+ ion


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NaOH+

Fe3+

Cr3+

large excess added

Fe(OH)3 Cr(OH)3

Precipitation of both Cr3+

and Fe3+ occurs


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NaOH

+

Fe3+

Cr3+

small excess added

slowly

Fe(OH)3

Cr3+

less soluble salt

precipitates only


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Solubility Rulesgeneral rules for predicting the solubility of ionic

compounds

strictly qualitativeDo not tell how soluble

Not quantitative


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My+ yAx-

Ax-

xMy+ My+

Ax-

MxAy

Solubility Equilibrium

saturated

solution

solid


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Solubility Equilibrium

MxAy(s) xMy+(aq) + yAx-(aq)

The equilibrium constant for this reaction is thesolubility product, Ksp:

Ksp = [My+]x[Ax-]y


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Solubility Product, Ksp

Ksp is related to molar solubility


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qualitative comparisons


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Ksp used to compare relative solubilities

smaller Ksp = less soluble

larger Ksp= more soluble


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qualitative comparisons

quantitative calculations


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Calculations with Ksp

Basic steps for solving solubility equilibriumproblems

Write the balanced chemical equation for the

solubility equilibrium and the expression for Ksp

Derive the mathematical relationship between Ksp

and molar solubility (x)

Make an ICE table

Substitute equilibrium concentrations of ions into Kspexpression

Using Ksp, solve for x or visa versa, depending on

what is wanted and the information provided


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Example 1(1 on Example Problems Handout)

Calculate the Ksp for MgF2 if the molar

solubility of this salt is 2.7 x 10-3 M.

(ans.: 7.9 x 10-8)


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Calculate the Ksp for Ca3(PO4)2 (FW = 310.2)

if the solubility of this salt is 8.1 x 10-4 g/L.(ans.: 1.3 x 10-26)


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The Ksp for CaF2 (FW = 78 g/mol) is 4.0 x 10-11.

What is the molar solubility of CaF2 in water?

What is the solubility of CaF2

in water in g/L?(ans.: 2.2 x 10-4 M, 0.017 g/L)


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Precipitation

Precipitation reactionexchange reaction

one product is insoluble

ExampleOverall: CaCl2(aq) + Na2CO3(aq) --> CaCO3(s) + 2NaCl(aq)


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Precipitation




Na+ and Ca2+exchange anions


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Precipitation




Net Ionic: Ca2+(aq) + CO3

2-(aq) CaCO3

(s)


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Precipitation

Compare precipitation to solubility equilibrium

Ca2+(aq) + CO32-(aq) CaCO3(s) prec.

vsCaCO3(s) Ca

2+(aq) + CO32-(aq) sol. Equil.

saturated solution


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Precipitation

Compare precipitation to solubility equilibrium:

Ca2+(aq) + CO32-(aq) CaCO3(s)

vsCaCO3(s) Ca

2+(aq) + CO32-(aq)

saturated solution

Precipitation occurs until solubility equilibrium is

established.


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Precipitation

Ca2+(aq) + CO32-(aq) CaCO3(s)

vs

CaCO3(s) Ca2+(aq) + CO3

2-(aq)

saturated solution

Key to forming ionic precipitates: Mix ions so

concentrations exceed those in saturated solution

(supersaturated solution)


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Predicting Precipitation

To determine if solution is supersaturated:

Compare ion product (Q or IP) to Ksp

ForMxAy(s) xMy+(aq) + yAx-(aq)

Q = [My+]x[Ax-]y

Q calculated for initial conditions

Q >Ksp supersaturated solution, precipitation

occurs, solubility equilibrium established (Q =

Ksp

)

Q = Ksp saturated solution, no precipitation

Q < Ksp unsaturated solution, no precipitation


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Predicting Precipitation

Basic Steps for Predicting Precipitation

Consult solubility rules (if necessary) to determine

what ionic compound might precipitate

Write the solubility equilibrium for this substance

Pay close attention to the stoichiometry

Calculate the moles of each ion involved before

mixing

moles = M x L or moles = mass/FW

Calculate the concentration of each ion involved

after mixing assuming no reaction

Calculate Q and compare to Ksp


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Example 4(7 and 8 on Example Problems Handout)

Will a precipitate form if (a) 500.0 mL of 0.0030

M lead nitrate, Pb(NO3)2, and 800.0 mL of

0.0040 M sodium fluoride, NaF, are mixed, and

(b) 500.0 mL of 0.0030 M Pb(NO3)2 and 800.0

mL of 0.040 M NaF are mixed?

(ans.: (a) No, Q = 7.5 x 10-9; (b) Yes, Q = 7.5 x 10-7)


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Solubility RulesAll alkali metal compounds are soluble

The nitrates of all metals are soluble in water.

Most compounds containing chloride are soluble. Theexceptions are those with Ag+, Pb2+, and Hg2

2+

Most compounds containing fluoride are soluble. The

exceptions are those with Mg2+, Ca2+, Sr2+, Ba2+, and

Pb2+

Ex. 4: Possible precipitate = PbF2 (Ksp = 4.1 x 10-8)


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Example 5(10 on Example Problem Handout)

A student carefully adds solid silver nitrate,

AgNO3, to a 0.0030 M solution of sodium

sulfate, Na2SO4. What [Ag+] in the solution is

needed to just initiate precipitation of silver

sulfate, Ag2SO4 (Ksp = 1.4 x 10-5)?

(ans.: 0.068 M)


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Factors that Affect Solubility

Common Ion Effect

pH

Complex-Ion Formation


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Factors that Affect Solubility

Common Ion Effect

pH

Complex-Ion Formation

These sure sound

familiar. Where

have I seen them

before?


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Common Ion Effect and Solubility

Consider the solubility equilibrium of AgCl.

AgCl(s) Ag+(aq) + Cl-(aq)

How does adding excess NaCl affect the

solubility equilibrium?

NaCl(s) Na+(aq) + Cl-(aq)


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Consider the solubility equilibrium of AgCl.


How does adding excess NaCl affect the

solubility equilibrium?

NaCl(s) Na+

(aq) + Cl-

(aq)2 sources of Cl-

Cl- is common ion


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Example 6(11 on Example Problem Handout)

What is the molar solubility of AgCl (Ksp = 1.8 x

10-10) in a 0.020 M NaCl solution? What is the

molar solubility of AgCl in pure water?(ans.: 8.5 x 10-9, 1.3 x 10-5)


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How does adding excess NaCl affect thesolubility equilibrium of AgCl?

AgCl in H2O

1.3 x 10-5 M

+ 0.020 M NaCl

Molars

olubility

Molarsolubility

AgCl in 0.020

M NaCl

8.5 x 10-9 M


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Why does the molar solubility of AgCl decrease

after adding NaCl?

Understood in terms of LeChateliers principle:

NaCl(s) --> Na+ + Cl-


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after adding NaCl?



AgCl(s) Ag+ + Cl-


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after adding NaCl?



AgCl(s) Ag+ + Cl-

Common-Ion Effect


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pH and Solubility

How can pH influence solubility?

Solubility of insoluble salts will be affected by pH

changes if the anion of the salt is at least moderately

basic

Solubility increases as pH decreases

Solubility decreases as pH increases


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pH and Solubility

Salts contain either basic or neutral anions:basic anions

Strong bases: OH-, O2-

Weak bases (conjugate bases of weak molecular acids):

F-, S2-, CH3COO-, CO3

2-, PO43-, C2O4

2-, CrO42-, etc.

Solubility affected by pH changes

neutral anions (conjugate bases of strong

monoprotic acids)Cl-, Br-, I-, NO3

-, ClO4-

Solubility not affected by pH changes


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pH and Solubility

Example:Fe(OH)2

Fe(OH)2(s) Fe2+(aq) + 2OH-(aq)


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pH and Solubility

Example:Fe(OH)2-Add acid



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pH and Solubility



2H3O+(aq) + 2OH-(aq) 4H

2O


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pH and Solubility



2H3O+(aq) + 2OH-(aq) 4H

2O

Which way does this reaction shift the solubility equilibrium? Why?

Understood in terms of LeChatliers principle


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pH and Solubility



2H3O+(aq) + 2OH-(aq) 4H2O

Decrease=stress

More Fe(OH)2 dissolves in response

Solubility increases

Stress relief = increase [OH-]


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pH and Solubility

Example:Fe(OH)2


2H3O+(aq) + 2OH-(aq) 4H

2O(l)

Fe(OH)2(s) + 2H3O+(aq) Fe2+(aq) + 4H2O(l)overall


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pH and Solubility

Example:Fe(OH)2


2H3O+(aq) + 2OH-(aq) 4H

2O(l)

Fe(OH)2(s) + 2H3O+(aq) Fe2+(aq) + 4H2O(l)overall

solubility increasesdecrease pH

solubility decreasesincrease pH


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pH, Solubility, and Tooth Decay

Enamel (hydroxyapatite) = Ca10(PO4)6(OH)2

(insoluble ionic compound)

Ca10(PO4)6(OH)2 10Ca2+(aq) + 6PO4

3-(aq) + 2OH-(aq)


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Enamel (hydroxyapatite) = Ca10(PO4)6(OH)2(insoluble ionic compound)

Ca10(PO4)6(OH)2 10Ca2+(aq) + 6PO4

3-(aq) + 2OH-(aq)

weak base strong base


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bacteria in mouth

+ food organic acids

(Yummy)

metabolism

(H3O+)


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Ca10(PO4)6(OH)2(s) 10Ca2+(aq) + 6PO4

3-(aq) + 2OH-(aq)

OH-(aq) + H3O+(aq) 2H2O(l)

PO43-(aq) + H3O

+(aq) HPO43-(aq) + H2O(l)


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Ca10(PO4)6(OH)2(s) 10Ca2+(aq) + 6PO4

3-(aq) + 2OH-(aq)

OH-(aq) + H3O+(aq) 2H2O(l)

PO43-(aq) + H3O+(aq) HPO43-(aq) + H2O(l)

Decrease=stress

More Ca10(PO4)6(OH)2dissolves in response


Leads to tooth decay

Decrease=stress


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Tooth Decay


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Why fluoridation?

F-replaces OH- in enamel

Ca10

(PO4)6(F)

2(s)

10Ca2+(aq) + 6PO

4

3-(aq) + 2F-(aq)

fluorapatite


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Why fluoridation?


Ca10

(PO4)6(F)

2(s)

10Ca2+(aq) + 6PO

4

3-(aq) + 2F-(aq)

Less soluble (has

lower Ksp) thanCa10(PO4)6(OH)2

weaker base than OH-

more resistant to acidattack

Factors together fight tooth decay!


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Why fluoridation?


Ca10

(PO4)6(F)

2(s)

10Ca2+(aq) + 6PO

4

3-(aq) + 2F-(aq)

F- added to drinking water as NaF or Na2SiF6

1 ppm = 1 mg/L

F- added to toothpastes as SnF2, NaF, or Na2PO3F

0.1 - 0.15 % w/w


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Complex Ion Formation and Solubility

Metals act as Lewis acids (see Chapter 15)Example

Fe3+(aq) + 6H2O(l) Fe(H2O)63+(aq)

Lewis acid Lewis base


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Fe3+(aq) + 6H2O(l) Fe(H2O)63+(aq)

Complex ion

Complex ion/complex contains central metal ion bonded

to one or more molecules or anions called ligands

Lewis acid = metal

Lewis base = ligand


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Fe3+(aq) + 6H2O(l) Fe(H2O)63+(aq)

Complex ion

Complex ions are often water soluble

Ligands often bond strongly with metals

Kf>> 1: Equilibrium lies very far to right.


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Metals act as Lewis acids (see Chapter 15)Other Lewis bases react with metals also

Examples

Fe3+

(aq) + 6CN-

(aq)

Fe(CN)63-

(aq)

Ni2+(aq) + 6NH3(aq) Ni(NH3)62+(aq)

Ag+(aq) + 2S2O32-(aq) Ag(S2O3)2

3-(aq)

Lewis acid Lewis base Complex ion




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Complex-Ion Formation and Solubility

How does complex ion formation influence

solubility?

Solubility of insoluble salts increases with addition

of Lewis bases if the metal ion forms a complex withthe base.


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Example

AgCl



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Example

AgCl-Add NH3


Ag+(aq) + 2NH3(aq) Ag(NH3)2+(aq)


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Example

AgCl-Add NH3



Which way does this reaction shift the solubility equilibrium? Why?


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Example

AgCl-Add NH3



Decrease=stressMore AgCl dissolves in response



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Example

AgCl



AgCl(s) + 2NH3(aq) Ag(NH3)2+(aq) + Cl-(aq)overall

Addition of ligand

solubility increases

Summary: Factors that Influence Solubility


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y y

Common Ion Effect

Decreases solubility

pH

pH decreases

Increases solubility

pH increases

Decreases solubility

Salt must have basic anionComplex-Ion Formation

Increases solubility


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Chapter 17b Solubility

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