Ch17 Outline (1)

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Aqueous

Equilibria

2009, Prentice-Hall, Inc.

Chapter 17

Additional Aspects ofAqueous Equilibria

Chemistry, The Central Science, 11th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

John D. Bookstaver

St. Charles Community College

Cottleville, MO

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Aqueous

Equilibria


The Common-Ion Effect

Consider a solution of acetic acid:

If acetate ion is added to the solution,

Le Chtelier says the equilibrium will

shift to the left.

CH3COOH(aq) + H2O(l) H3O+(aq)+ CH3COO

(aq)

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Aqueous

Equilibria



The extent of ionization of a weak

electrolyte is decreased by adding to

the solution a strong electrolyte that

has an ion in common with the weak

electrolyte.

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Aqueous

Equilibria



Calculate the fluoride ion concentration and pH

of a solution that is 0.20 Min HF and 0.10 MinHCl.

Kafor HF is 6.8 104.

[H3O+] [F]

[HF]Ka= = 6.8 10

-4

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Aqueous

Equilibria



Because HCl, a strong acid, is also present,

the initial [H3O+] is not 0, but rather 0.10 M.

[HF], M [H3O+], M [F], M

Initially 0.20 0.10 0Change x +x +x

At Equilibrium 0.20 x 0.20 0.10 + x 0.10 x

HF(aq) + H2O(l) H3O+(aq) + F(aq)

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Aqueous

Equilibria



= x

1.4 103

= x

(0.10) (x)(0.20)

6.8 104 =

(0.20) (6.8 104)

(0.10)

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Aqueous

Equilibria



Therefore, [F] = x= 1.4 103

[H3O+] = 0.10 + x= 0.10 + 1.4 103 = 0.10 M

So, pH = log (0.10)pH = 1.00

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Aqueous

Equilibria


Buffers

Buffers are solutions

of a weak conjugate

acid-base pair. They are particularly

resistant to pH

changes, even when

strong acid or base isadded.

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Aqueous

Equilibria


Buffers

If a small amount of hydroxide is added to an

equimolar solution of HF in NaF, for example, the HF

reacts with the OH to make F and water.

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Aqueous

Equilibria


Buffers

Similarly, if acid is added, the F reacts with it to formHF and water.

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Aqueous

Equilibria


Buffer Calculations

Consider the equilibrium constant

expression for the dissociation of a

generic acid, HA:

[H3O+] [A]

[HA]Ka=

HA + H2O H3O+ + A

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Aqueous

Equilibria


Buffer Calculations

Rearranging slightly, this becomes

[A]

[HA]Ka= [H3O+]

Taking the negative log of both side, we get

[A][HA]

log Ka=log [H3O+] + log

pKa

pH

acid

base

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Aqueous

Equilibria


Buffer Calculations

So

pKa= pH log[base]

[acid]

Rearranging, this becomes

pH = pKa+ log[base]

[acid]

This is the HendersonHasselbalchequation.

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Aqueous

Equilibria


HendersonHasselbalch Equation

What is the pH of a buffer that is 0.12 Min lactic acid, CH3CH(OH)COOH, and

0.10 Min sodium lactate? Ka for lactic

acid is 1.4 104.

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Aqueous

Equilibria


HendersonHasselbalch Equation

pH = pKa+ log[base]

[acid]

pH = log (1.4 104) + log (0.10)(0.12)

pH

pH = 3.77

pH = 3.85 + (0.08)

pH

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Aqueous

Equilibria


pH Range

The pH range is the range of pH values

over which a buffer system works

effectively.

It is best to choose an acid with a pKaclose to the desired pH.

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Aqueous

Equilibria


When Strong Acids or Bases Are

Added to a Bufferit is safe to assume that all of the strong acid

or base is consumed in the reaction.

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Aqueous

Equilibria


Addition of Strong Acid or Base to a

Buffer

1. Determine how the neutralization

reaction affects the amounts of

the weak acid and its conjugatebase in solution.

2. Use the HendersonHasselbalch

equation to determine the new

pH of the solution.

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Aqueous

Equilibria


Calculating pH Changes in Buffers

A buffer is made by adding 0.300 mol

HC2H3O2 and 0.300 mol NaC2H3O2 to

enough water to make 1.00 L of

solution. The pH of the buffer is 4.74.Calculate the pH of this solution after

0.020 mol of NaOH is added.

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Aqueous

Equilibria



Before the reaction, since

mol HC2H3O2 = mol C2H3O2

pH = pKa= log (1.8 105) = 4.74

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Aqueous

Equilibria



The 0.020 mol NaOH will react with 0.020 mol of the

acetic acid:

HC2H3O2(aq) + OH(aq) C2H3O2(aq) + H2O(l)

HC2H3O2 C2H3O2 OH

Before reaction 0.300 mol 0.300 mol 0.020 mol

After reaction 0.280 mol 0.320 mol 0.000 mol

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Aqueous

Equilibria



Now use the HendersonHasselbalch equation to

calculate the new pH:

pH = 4.74 + log (0.320)(0.200)

pH = 4.74 + 0.06pH

pH = 4.80

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Aqueous

Equilibria


Titration

A pH meter or

indicators are used to

determine when the

solution has reachedthe equivalence point,

at which the

stoichiometric amount

of acid equals that ofbase.

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Aqueous

Equilibria


Titration of a Strong Acid with a

Strong BaseFrom the start of the

titration to near the

equivalence point,

the pH goes upslowly.

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Aqueous

Equilibria



Strong BaseJust before (and

after) the equivalence

point, the pH

increases rapidly.

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Aqueous

Equilibria



Strong BaseAt the equivalence

point, moles acid =

moles base, and the

solution containsonly water and the

salt from the cation

of the base and the

anion of the acid.

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Equilibria



Strong BaseAs more base is

added, the increase

in pH again levels

off.

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Equilibria


Titration of a Weak Acid with a

Strong Base Unlike in the previous

case, the conjugatebase of the acidaffects the pH when it

is formed.

At the equivalencepoint the pH is >7.

Phenolphthalein iscommonly used as anindicator in thesetitrations.

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Aqueous

Equilibria



Strong Base

At each point below the equivalence point, the

pH of the solution during titration is determined

from the amounts of the acid and its conjugatebase present at that particular time.

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Aqueous

Equilibria



Strong BaseWith weaker acids,

the initial pH is

higher and pH

changes near theequivalence point

are more subtle.

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Aqueous

Equilibria


Titration of a Weak Base with a

Strong Acid

The pH at the

equivalence point in

these titrations is < 7. Methyl red is the

indicator of choice.

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Aqueous

Equilibria


Titrations of Polyprotic Acids

When one

titrates a

polyprotic acid

with a basethere is an

equivalence

point for each

dissociation.

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Aqueous

Equilibria


Solubility Products

Consider the equilibrium that exists in a

saturated solution of BaSO4 in water:

BaSO4(s) Ba2+(aq) + SO4

2(aq)

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Aqueous

Equilibria


Solubility Products

The equilibrium constant expression for

this equilibrium is

Ksp= [Ba2+] [SO4

2]

where the equilibrium constant, Ksp, iscalled the solubility product.

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Aqueous

Equilibria


Solubility Products

Ksp is notthe same as solubility.

Solubility is generally expressed as the mass

of solute dissolved in 1 L (g/L) or 100 mL

(g/mL) of solution, or in mol/L (M).

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Aqueous

Equilibria


Factors Affecting Solubility


If one of the ions in a solution equilibrium

is already dissolved in the solution, the

equilibrium will shift to the left and thesolubility of the salt will decrease.

BaSO4(s) Ba2+(aq) + SO42(aq)

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Aqueous

Equilibria



pH

If a substance has a

basic anion, it will be

more soluble in an

acidic solution.

Substances with

acidic cations are

more soluble in

basic solutions.

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Aqueous

Equilibria



Complex Ions Metal ions can act as Lewis acids and form

complex ions with Lewis bases in the solvent.

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Equilibria



Complex Ions

The formation

of these

complex ions

increases the

solubility of

these salts.

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Aqueous

Equilibria



Amphoterism

Amphoteric metal

oxides and hydroxides

are soluble in strong

acid or base, because

they can act either as

acids or bases.

Examples of suchcations are Al3+, Zn2+,

and Sn2+.

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Aqueous

Equilibria


Will a Precipitate Form?

In a solution,

IfQ= Ksp, the system is at equilibriumand the solution is saturated.

IfQ< Ksp, more solid can dissolveuntil Q= Ksp.

IfQ> Ksp, the salt will precipitate until

Q= Ksp.

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Aqueous

Equilibria

Selective Precipitation of Ions

One can use

differences in

solubilities of

salts to separateions in a

mixture.

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Ch17 Outline (1)