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8/12/2019 Chap 7 Additional Aspects of Aqueous Equilibria
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Copyright 1999, PRENTICE HALL Chapter 17 1
Additional Aspects ofAqueous Equil ibr ia
Chapter 17
David P. White
University of North Carol ina, Wilmington
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Copyright 1999, PRENTICE HALL Chapter 17 2
The Common I on Effect
The solubility of a partially soluble salt is decreased
when a common ion is added. Consider the equilibrium established when acetic
acid, HC2H3O2, is added to water.
At equilibrium H+and C2
H3
O2
-are constantly moving
into and out of solution, but the concentrations of ions
is constant and equal.
If a common ion is added, e.g. C2H3O2-from
NaC2H3O2(which is a strong electrolyte) then[C2H3O2
-] increases and the system is no longer at
equilibrium.
So, [H+] must decrease.
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Copyright 1999, PRENTICE HALL Chapter 17 3
Buffered Solutions
Composition and Action of Buffered Solutions
A buffer consists of a mixture of a weak acid (HX)and its conjugate base (X-):
The Kaexpression is
A buffer resists a change in pH when a small amount
of OH-or H+is added.
HX(aq) H+(aq) + X-(aq)
.]X[
[HX]]H[
.[HX]
]X][H[
-
-
a
a
K
K
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Copyright 1999, PRENTICE HALL Chapter 17 4
Buffered Solutions
Composition and Action of Buffered Solutions
When OH-is added to the buffer, the OH-reacts withHX to produce X-and water. But, the [HX]/[X-] ratio
remains more or less constant, so the pH is not
significantly changed.
When H+is added to the buffer, X-is consumed to
produce HX. Once again, the [HX]/[X-] ratio is more
or less constant, so the pH does not change
significantly.
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Copyright 1999, PRENTICE HALL Chapter 17 5
Buffered Solutions
Composition and Action of Buffered Solutions
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Copyright 1999, PRENTICE HALL Chapter 17 6
Buffered Solutions
Buffer Capacity and pH
Buffer capacity is the amount of acid or baseneutralized by the buffer before there is a significant
change in pH.
Buffer capacity depends on the composition of thebuffer.
The greater the amounts of conjugate acid-base pair,
the greater the buffer capacity.
The pH of the buffer depends on Ka.
8/12/2019 Chap 7 Additional Aspects of Aqueous Equilibria
7/45Copyright 1999, PRENTICE HALL Chapter 17 7
Buffered Solutions
Buffer Capacity and pH
If Kais small (i.e., if the equilibrium concentration ofundissociated acid is close to the initial
concentration), then
.[HX]
]X[logppH
.]X[
[HX]loglog]Hlog[
]X[
[HX]
]H[
-
-
-
a
a
a
K
K
K
8/12/2019 Chap 7 Additional Aspects of Aqueous Equilibria
8/45Copyright 1999, PRENTICE HALL Chapter 17 8
Buffered Solutions
Addition of Strong Acids or Bases to Buffers
We break the calculation into two parts:stoichiometric and equilibrium.
The amount of strong acid or base added results in a
neutralization reaction:X-+ H3O
+ HX + H2O
HX + OH- X-+ H2O.
By knowing how must H3
O+or OH-was added
(stoichiometry) we know how much HX or X-is
formed.
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Copyright 1999, PRENTICE HALL Chapter 17 10
Buffered Solutions
Addition of Strong Acids or Bases to Buffers
With the concentrations of HX and X-(note thechange in volume of solution) we can calculate the pH
from the Henderson-Hasselbalch equation
acidconjugate
baseconjugate
logppH
[HX]
]X[logppH
-
a
a
K
K
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Copyright 1999, PRENTICE HALL Chapter 17 11
Acid-Base Titrations
Strong Acid-Base Titrations
The plot of pH
versus volume
during a titration
is a titration curve.
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Copyright 1999, PRENTICE HALL Chapter 17 12
Acid-Base Titrations
Strong Acid-Base Titrations
Consider adding a strong base (e.g. NaOH) to asolution of a strong acid (e.g. HCl).
Before any base is added, the pH is given by the strong acid
solution. Therefore, pH < 7.
When base is added, before the equivalence point, the pH is
given by the amount of strong acid in excess. Therefore, pH
< 7.
At equivalence point, the amount of base added is
stoichiometrically equivalent to the amount of acid
originally present. Therefore, the pH is determined by the
salt solution. Therefore, pH = 7.
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Copyright 1999, PRENTICE HALL Chapter 17 13
Acid-Base Titrations
Strong Acid-Base Titrations
Consider adding a strong base (e.g. NaOH) to asolution of a strong acid (e.g. HCl).
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Copyright 1999, PRENTICE HALL Chapter 17 14
Acid-Base Titrations
Strong Acid-Base Titrations
We know the pH at equivalent point is 7.00.
To detect the equivalent point, we use an indicator
that changes color somewhere near 7.00. Usually, we use phenolphthalein that changes color between
pH 8.3 to 10.0. In acid, phenolphthalein is colorless.
As NaOH is added, there is a slight pink color at theaddition point.
When the flask is swirled and the reagents mixed, the pinkcolor disappears.
At the end point, the solution is light pink.
If more base is added, the solution turns darker pink.
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Copyright 1999, PRENTICE HALL Chapter 17 15
Acid-Base Titrations
Strong Acid-Base Titrations
The equivalence point in a titration is the point atwhich the acid and base are present in stoichiometric
quantities.
The end point in a titration is the observed point. The difference between equivalence point and end
point is called the titration error.
The shape of a strong base-strong acid titration curve
is very similar to a strong acid-strong base titrationcurve.
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Copyright 1999, PRENTICE HALL Chapter 17 16
Acid-Base Titrations
Strong Acid-Base Titrations
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Copyright 1999, PRENTICE HALL Chapter 17 17
Acid-Base Titrations
Strong Acid-Base Titrations
Initially, the strong base is in excess, so the pH > 7.
As acid is added, the pH decreases but is still greater
than 7.
At equivalence point, the pH is given by the saltsolution (i.e. pH = 7).
After equivalence point, the pH is given by the strong
acid in excess, so pH < 7.
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Copyright 1999, PRENTICE HALL Chapter 17 18
Acid-Base Titrations
Weak Acid-Strong Base Titrations
Consider the titration of acetic acid, HC2H3O2andNaOH.
Before any base is added, the solution contains only
weak acid. Therefore, pH is given by the equilibrium
calculation.
As strong base is added, the strong base consumes a
stoichiometric quantity of weak acid:
HC2H3O2(aq) + NaOH(aq) C2H3O2-(aq) + H2O(l)
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Copyright 1999, PRENTICE HALL Chapter 17 19
Acid-Base Titrations
Weak Acid-Strong Base Titrations
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Copyright 1999, PRENTICE HALL Chapter 17 20
Acid-Base Titrations
Weak Acid-Strong Base Titrations
There is an excess of acetic acid before theequivalence point.
Therefore, we have a mixture of weak acid and its
conjugate base.
The pH is given by the buffer calculation.
First the amount of C2H3O2-generated is calculated, as well as the
amount of HC2H3O2consumed. (Stoichiometry.)
Then the pH is calculated using equilibrium conditions. (Henderson-
Hasselbalch.)
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Copyright 1999, PRENTICE HALL Chapter 17 21
Acid-Base Titrations
Weak Acid-Strong Base Titrations
At the equivalence point, all the acetic acid has beenconsumed and all the NaOH has been consumed.
However, C2H3O2-has been generated.
Therefore, the pH is given by the C2H3O2-solution.
This means pH > 7.
More importantly, pH 7 for a weak acid-strong base titration.
After the equivalence point, the pH is given by the
strong base in excess.
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Copyright 1999, PRENTICE HALL Chapter 17 22
Acid-Base Titrations
Weak Acid-Strong Base Titrations
For a strong acid-strong base titration, the pH beginsat less than 7 and gradually increases as base is
added.
Near the equivalence point, the pH increases
dramatically.
For a weak acid-strong base titration, the initial pH
rise is more steep than the strong acid-strong base
case. However, then there is a leveling off due to buffer
effects.
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Copyright 1999, PRENTICE HALL Chapter 17 23
Acid-Base Titrations
Weak Acid-Strong Base Titrations
The inflection point is not as steep for a weak acid-strong base titration.
The shape of the two curves after equivalence point is
the same because pH is determined by the strong base
in excess.
Two features of titration curves are affected by the
strength of the acid:
the amount of the initial rise in pH, and
the length of the inflection point at equivalence.
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Copyright 1999, PRENTICE HALL Chapter 17 24
Acid-Base Titrations
Weak Acid-Strong Base Titrations
The weaker the acid,
the smaller theequivalence point
inflection.
For very weak acids, it
is impossible to detectthe equivalence point.
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Copyright 1999, PRENTICE HALL Chapter 17 25
Acid-Base Titrations
Weak Acid-Strong Base Titrations
Titration of weak bases with strong acids have similarfeatures to weak acid-strong base titrations.
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Copyright 1999, PRENTICE HALL Chapter 17 26
Acid-Base Titrations
Titrations of Polyprotic Acids
In polyprotic acids, each ionizable proton dissociatesin steps.
Therefore, in a titration there are nequivalence points
corresponding to each ionizable proton.
In the titration of Na2CO3with HCl there are two
equivalence points:
one for the formation of HCO3-
one for the formation of H2CO3.
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Copyright 1999, PRENTICE HALL Chapter 17 27
Acid-Base Titrations
Titrations of Polyprotic Acids
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Copyright 1999, PRENTICE HALL Chapter 17 28
Solubil i ty Equi l ibr ia
Solubility-Product Constant, Ksp
Consider
for which
Ksp is the solubility product. (BaSO4 is ignored
because it is a pure solid so its concentration is
constant.)
BaSO4(s) Ba2+(aq) + SO4
2-(aq)
]SO][Ba[ -242spK
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Copyright 1999, PRENTICE HALL Chapter 17 29
Solubil i ty Equi l ibr ia
Solubility-Product Constant, Ksp
In general: the solubility product is the molarconcentration of ions raised to their stoichiometric
powers.
Solubility is the amount (grams) of substance that
dissolves to form a saturated solution.
Molar solubility is the number of moles of solute
dissolving to form a liter of saturated solution.
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Copyright 1999, PRENTICE HALL Chapter 17 30
Solubil i ty Equi l ibr ia
Solubility and Ksp
To convert solubility to Ksp
solubility needs to be converted into molar solubility
(via molar mass);
molar solubility is converted into the molarconcentration of ions at equilibrium (equilibrium
calculation),
Kspis the product of equilibrium concentration of
ions.
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Copyright 1999, PRENTICE HALL Chapter 17 31
Solubil i ty Equi l ibr ia
Solubility and Ksp
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Copyright 1999, PRENTICE HALL Chapter 17 32
Factors That Affect Solubil i ty
Common-Ion Effect
Solubility is decreased when a common ion is added.
This is an application of Le Chteliers principle:
as F- (from NaF, say) is added, the equilibrium shiftsaway from the increase.
Therefore, CaF2(s) is formed and precipitation occurs.
As NaF is added to the system, the solubility of CaF2
decreases.
CaF2(s) Ca2+(aq) + 2F-(aq)
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Copyright 1999, PRENTICE HALL Chapter 17 33
Factors That Affect Solubil i ty
Common-Ion Effect
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Copyright 1999, PRENTICE HALL Chapter 17 34
Factors That Affect Solubil i ty
Solubility and pH
Again we apply Le Chteliers principle:
If the F-is removed, then the equilibrium shifts towards the
decrease and CaF2dissolves.
F-can be removed by adding a strong acid:
As pH decreases, [H+] increases and solubility increases.
The effect of pH on solubility is dramatic.
CaF2(s) Ca2+(aq) + 2F-(aq)
F-(aq) + H+(aq) HF(aq)
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Copyright 1999, PRENTICE HALL Chapter 17 35
Factors That Affect Solubil i ty
Solubility and pH
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Copyright 1999, PRENTICE HALL Chapter 17 36
Factors That Affect Solubil i ty
Formation of Complex Ions
Consider the formation of Ag(NH3)2+:
The Ag(NH3)2+is called a complex ion.
NH3(the attached Lewis base) is called a ligand.
The equilibrium constant for the reaction is called the
formation constant, Kf:
Focus on Lewis acid-base chemistry and solubility.
Ag+(aq) + 2NH3(aq) Ag(NH3)2(aq)
.]][NH[Ag
])Ag(NH[
2323fK
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Copyright 1999, PRENTICE HALL Chapter 17 37
Factors That Affect Solubil i ty
Formation of Complex Ions
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Copyright 1999, PRENTICE HALL Chapter 17 38
Factors That Affect Solubil i ty
Formation of Complex Ions
Consider the addition of ammonia to AgCl (whiteprecipitate):
The overall reaction is
Effectively, the Ag+(aq) has been removed from
solution.
By Le Chteliers principle, the forward reaction (the
dissolving of AgCl) is favored.
AgCl(s) Ag+(aq) + Cl-(aq)
Ag+
(aq) + 2NH3(aq) Ag(NH3)2(aq)
AgCl(s) + 2NH3(aq) Ag(NH3)2(aq) + Cl-(aq)
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Copyright 1999, PRENTICE HALL Chapter 17 39
Factors That Affect Solubil i ty
Amphoterism
Amphoteric oxides will dissolve in either a strong acidor a strong base.
Examples: hydroxides and oxides of Al3+, Cr3+, Zn2+,
and Sn2+.
The hydroxides generally form complex ions with four
hydroxide ligands attached to the metal:
Hydrated metal ions act as weak acids. Thus, the
amphoterism is interrupted:
Al(OH3)(s) + OH-(aq) Al(OH)4
-(aq)
F Th Aff S l bil i
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Copyright 1999, PRENTICE HALL Chapter 17 40
Factors That Affect Solubil i ty
Amphoterism
Hydrated metal ions act as weak acids. Thus, theamphoterism is interrupted:
Al(H2O)63+(aq) + OH-(aq) Al(H2O)5(OH)
2+(aq) + H2O(l)
Al(H2O)5(OH)
2+
(aq) + OH
-
(aq) Al(H2O)4(OH)2+
(aq) + H2O(l)Al(H2O)4(OH)
+(aq) + OH-(aq) Al(H2O)3(OH)3(s) + H2O(l)
Al(H2O)3(OH)3(s) + OH-(aq) Al(H2O)2(OH)4
-(aq) + H2O(l)
P i it ti d S ti f I
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Copyright 1999, PRENTICE HALL Chapter 17 41
Precipitation and Separation of I ons
At any instant in time, Q = [Ba2+][SO42-]. If Q< Ksp, precipitation occurs until Q= Ksp.
If Q= Ksp, equilibrium exists.
If Q> Ksp, solid dissolves until Q= Ksp.
Based on solubilities, ions can be selectively removed
from solutions.
Consider a mixture of Zn2+(aq) and Cu2+(aq). CuS
(Ksp= 6 10-37)is less soluble than ZnS (Ksp= 210-25), CuS will be removed from solution before ZnS.
BaSO4(s) Ba2+(aq) + SO4
2-(aq)
P i it ti d S ti f I
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Copyright 1999, PRENTICE HALL Chapter 17 42
Precipitation and Separation of I ons
As H2S is added to the green solution, black CuS
forms in a colorless solution of Zn2+
(aq). When more H2S is added, a second precipitate of
white ZnS forms.
Selective Precipitation of Ions
Ions can be separated from each other based on their
salt solubilities.
Example: if HCl is added to a solution containing Ag+
and Cu2+, the silver precipitates (Kspfor AgCl is 1.810-10) while the Cu2+remains in solution.
Removal of one metal ion from a solution is called
selective precipitation.
Q lit ti A l i f M t l l i El t
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Copyright 1999, PRENTICE HALL Chapter 17 43
Qualitative Analysis for Metall ic Elements
Qualitative analysis is
designed to detect the
presence of metal ions. Quantitative analysis is
designed to determine how
much metal ion is present.
Q lit ti A l i f M t l l i El t
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Copyright 1999, PRENTICE HALL Chapter 17 44
Qualitative Analysis for Metall ic Elements
We can separate a complicated mixture of ions into
five groups: Add 6 M HCl to precipitate insoluble chlorides (AgCl,
Hg2Cl2, and PbCl2).
To the remaining mix of cations, add H2S in 0.2 MHCl to
remove acid insoluble sulfides (e.g. CuS, Bi2S3, CdS, PbS,HgS, etc.).
To the remaining mix, add (NH4)2S at pH 8 to remove base
insoluble sulfides and hydroxides (e.g. Al(OH)3, Fe(OH)3,
ZnS, NiS, CoS, etc.).
To the remaining mixture add (NH4)2HPO4 to remove
insoluble phosphates (Ba3(PO4)2, Ca3(PO4)2, MgNH4PO4).
The final mixture contains alkali metal ions and NH4+.
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End of Chapter 17
Additional Aspects of
Aqueous Equil ibr ia