Acid-Base Equilibria
Chapter 16
The __________________is the shift in equilibrium caused by
the addition of a compound having an __________in common
with the______________ substance.
The presence of a common ion
_____________the ionization of a
________acid or a __________base.
Consider mixture of CH3COONa (_______electrolyte) and
CH3COOH (_______ acid).
CH3COONa (s)
Na+ (aq) +)
CH3COOH (aq)
H+ (aq) +)
common
ion
16.2
A ________________is a solution of:
1. A weak ______or a weak _______and
2. The _________of the weak acid or weak base
Both must be present!
A ____________has the ability to _________changes in
___upon the addition of small amounts of either _______or
__________.
Consider an equal molar mixture of CH3COOH and CH3COONa
CH3COOH (aq)
H+ (aq) + CH3COO- (aq)
Adding more _______creates a shift left IF
enough _________ions are present
16.3
Which of the following are buffer systems? (a) KF/HF
(b) KCl/HCl, (c) Na2CO3/NaHCO3
(a) KF is a weak acid and F- is its conjugate base
_______________
(b) HCl is a strong acid
_______________
(c) CO32- is a weak base and HCO3- is it conjugate acid
_______________
16.3
What is the pH of a solution containing 0.30 M HCOOH
and 0.52 M HCOOK?
Mixture of weak acid and conjugate base!
HCOOH (aq)
H+ (aq) + HCOO- (aq)
Initial (M)
Change (M)
Equilibrium (M)
Ka for HCOOH = 1.8 x 10 -4
16.2
What is the pH of a solution containing 0.30 M HCOOH
and 0.52 M HCOOK?
Mixture of weak acid and conjugate base!
HCOOH (aq)
H+ (aq) + HCOO- (aq)
Initial (M)
Change (M)
Equilibrium (M)
Common ion effect
16.2
HCl
HCl + CH3COO-
H+ + ClCH3COOH + Cl-
16.3
Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer
system. What is the pH after the addition of 20.0 mL of
0.050 M NaOH to 80.0 mL of the buffer solution?
Initial
Change
End

16.3
Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer
system. What is the pH after the addition of 20.0 mL of
0.050 M NaOH to 80.0 mL of the buffer solution?
final volume = 80.0 mL + 20.0 mL = 100 mL
start (M)
end (M)
pH = 9.20
16.3
Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer
system. What is the pH after the addition of 20.0 mL of
0.050 M NaOH to 80.0 mL of the buffer solution?
[
[
[
start (M)
end (M)
=
16.3
Chemistry In Action: Maintaining the pH of Blood
16.3
Titrations
In a ____________a solution of accurately known
_____________is added gradually added to another solution of
__________concentration until the chemical reaction between
the two solutions is complete.
___________________– the point at which the reaction is complete
___________– substance that changes color at the
________ (hopefully close to the equivalence point)
Slowly add base
to unknown acid
UNTIL
The indicator
changes color
(________)
4.7
Strong Acid-Strong Base Titrations
NaOH (aq) + HCl (aq)
OH- (aq) + H+ (aq)
H2O (l) + NaCl (aq)
H2O (l)
100%
ionization!
___________
_
16.4
_________Acid-________ Base Titrations
CH3COOH (aq) + NaOH (aq)
CH3COOH (aq) + OH- (aq)
CH3COONa (aq) + H2O (l)
CH3COO- (aq) + H2O (l)
At equivalence point (pH ______ 7):
CH3COO- (aq) + H2O (l)
OH- (aq) + CH3COOH (aq)
16.4
_________Acid-__________ Base Titrations
HCl (aq) + NH3 (aq)
H+ (aq) + NH3 (aq)
NH4Cl (aq)
NH4Cl (aq)
At equivalence point (pH _____7):
NH4+ (aq) + H2O (l)
NH3 (aq) + H+ (aq)
16.4
Acid-Base Indicators
16.5
The titration curve of a strong acid with a strong base.
16.5
Which indicator(s) would you use for a titration of HNO2
with KOH ?
_______acid titrated with _________base.
At equivalence point, will have __________________acid.
At equivalence point, pH _____ 7
Use __________or__________________
16.5
Finding the Equivalence Point
(calculation method)
• ________Acid vs. ________Base
– ________% ionized! pH ____No__________!
• ________Acid vs. _______ Base
– ______is____________; Need ____for conjugate
________equilibrium
• _________Acid vs. ________Base
– Base is________________; Need ____for
conjugate ________equilibrium
• _________ Acid vs. ______ Base
– Depends on the ______of both; could be
conjugate______, conjugate ______, or pH ___
Exactly 100 mL of 0.10 M HNO2 are titrated with 100 mL of
a 0.10 M NaOH solution. What is the pH at the
equivalence point ?
start (moles)
HNO2 (aq) + OH- (aq)
NO2- (aq) + H2O (l)
end (moles)
Final volume = 200 mL
NO2- (aq) + H2O (l)
Initial (M)
Change (M)
Equilibrium (M)
Kb =
=
=
OH- (aq) + HNO2 (aq)
_____________Ion Equilibria and Solubility
A __________is an ion containing a central metal cation
bonded to one or more molecules or ions.
Co2+ (aq) + 4Cl- (aq)
Co(H2O)2+
6
CoCl42- (aq)
CoCl24
16.10
16.10
Complex Ion Formation
• These are usually formed from a
transition metal surrounded by ligands
(polar molecules or negative ions).
• As a "rule of thumb" you place twice the
number of ligands around an ion as the
charge on the ion... example: the dark
blue Cu(NH3)42+ (ammonia is used as a
test for Cu2+ ions), and Ag(NH3)2+.
• Memorize the common ligands.
Common Ligands
Ligands
Names used in the ion
H2O
NH3
aqua
ammine
OHClBrCNSCN-
hydroxy
chloro
bromo
cyano
thiocyanato (bonded through
sulphur)
isothiocyanato (bonded through
nitrogen)
Names
• Names: ligand first, then cation
Examples:
– tetraamminecopper(II) ion: Cu(NH3)42+
– diamminesilver(I) ion: Ag(NH3)2+.
– tetrahydroxyzinc(II) ion: Zn(OH)4 2-
• The charge is the sum of the parts
(2+) + 4(-1)= -2.
When Complexes Form
• Aluminum also forms complex ions as do some post
transitions metals. Ex: Al(H2O)63+
• Transitional metals, such as Iron, Zinc and Chromium,
can form complex ions.
• The odd complex ion, FeSCN2+, shows up once in a
while
• Acid-base reactions may change NH3 into NH4+ (or vice
versa) which will alter its ability to act as a ligand.
• Visually, a precipitate may go back into solution as a
complex ion is formed. For example, Cu2+ + a little
NH4OH will form the light blue precipitate, Cu(OH)2. With
excess ammonia, the complex, Cu(NH3)42+, forms.
• Keywords such as "excess" and "concentrated" of any
solution may indicate complex ions. AgNO3 + HCl forms
the white precipitate, AgCl. With excess, concentrated
HCl, the complex ion, AgCl2-, forms and the solution
clears.
Coordination Number
• Total number of bonds from the ligands
to the metal atom.
• Coordination numbers generally range
between 2 and 12, with 4
(tetracoordinate) and 6
(hexacoordinate) being the most
common.
Some Coordination Complexes
molecular
formula
Lewis
base/ligand
Lewis acid donor
atom
coordination
number
Ag(NH3)2+
NH3
Ag+
N
2
[Zn(CN)4]2- CN-
Zn2+
C
4
[Ni(CN)4]2-
CN-
Ni2+
C
4
[PtCl6] 2-
Cl-
Pt4+
Cl
6
Ni2+
N
6
[Ni(NH3)6]2+ NH3