Transcript AP Notes Chapter17

ACIDS
&
BASES
Arrhenius Theory
1. in aqueous solution
+
2. Acid: produces H
3. Base: produces OH
Acid
HA

 H3O+
H2O
A
+
+
HA +
O
H
O
H
H
H
H
+ A
-
HCl(g) + H2O H3O+(aq) + Cl-(aq)
CH3COOH(l) + H2O = H3O+(aq) + CH3COO-(aq)
careless, but often seen
HCl 
+
H +
Cl
CH3COOH H+ + CH3COO-
Base
NaOH(s)


H2O
Na+(aq) + OH-(aq)
Arrhenius acid/base reaction
acid + base  H2O + a salt
HA + MOH  HOH + MA
Monoprotic acid: HCl
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
H+ + Cl- + Na+ + OH- 
H+ + OH-  H2O
HCl 
H2O + Na+ + Cl-
Diprotic acid: H2SO4
H2SO4 (aq) + 2NaOH (aq)  2H2O(l) + Na2SO4 (aq)
H+ + OH-  H2O
H2SO4 
Triprotic acid: H3PO4
Polyprotic
H3PO4(aq) + 3NaOH(aq)  3H2O(l) + Na3PO4(aq)
H3PO4 + 3 OH-  3 H2O + PO43-
H3PO4 
Bronsted-Lowry Theory
1. aqueous & nonaqueous solutions
2. Acid: species donating a proton
HCl  H+ + ClH2SO4  H+ + HSO4CH3COOH  H+ + CH3COO-
Bronsted-Lowry Theory
3. Base: species accepting a proton
OH- + H+  HOH
H2O + H+  H3O+
NH3 + H+  NH4+
Conjugate acid-base pairs
acid1 + base1
acid2 + base2
conjugate pairs
HF + HOH
Conjugate acid-base pairs
acid1 + base1
acid2 + base2
conjugate pairs
HF + HOH
+
H3O
+
F
ALL Arrhenius reactions
are Bronsted-Lowry
reactions
HCl + NaOH  H2O + NaCl
NOT all Bronsted reactions
are Arrhenius reactions
CH3COOH + NH3 
NH4+ + CH3COO-
Amphiprotic = Amphoteric
Can act as either an acid or a base
HCl + HOH  H3 +
NH3 + HOH  NH4+ + OHNH3 + OH-  NH2- + HOH
HOH + HOH  H3O+ + OH+
O
Cl
ACID STRENGTH
Relative ability of a
compound to
donate a proton
Base strength is
considered a result,
not a cause
REVIEW
Strong acid
100% dissociation
Weak acid
<100% dissociation
Notice this is NOT related
to concentration
Electronegativity is
the most significant
factor influencing
the strength of
acids & bases
HF > HCl > HBr > HI
as acids in nonaqueous solvents,
or as pure gases
Look at difference in
electronegativities
2.1 H - F
2.1 H - Cl
2.1 H - Br
2.1 H - I
4.0
3.0
2.8
2.5
Most “ionic” is
the most acidic
Nonpolar
Polar
Ionic
ED 0
ED 1.7
ED 4.0
However,
as acids
in aqueous solution
HF < HCl = HBr = HI
2.1 H - O 3.5
competition!
2.1 H - F 4.0
2.1 H - Cl 3.0
2.1 H - Br 2.8
2.1 H - I 2.5
Is methane acidic
as a gas or in
aqueous solution?
2.1 H - C 2.5
The strength of
oxy-acids are also
dependent on
electronegativity.
Oxy-acids and
bases have the
same fundamental
structure
NaOH: Na - O - H
0.9 3.5 2.1
HClO: Cl - O - H
3.0 3.5 2.1
In water, the more
“ionic” bond
dissociates, forming
the acid or base
NaOH: Na - O - H
0.9 3.5 2.1
HClO: Cl - O - H
3.0 3.5 2.1
Are alcohols acids
or bases?
C-O-H
2.5 3.5 2.1
Acids in
homologous series
are of different
strength
Acid Strength
H2SO4 > H2SO3
HNO3 > HNO2
HClO4 > HClO3 > HClO2 > HClO
Structurally
H2SO4 = O2S(OH)2
H2SO3 = OS(OH)2
Need to examine
formal charge of
central atom.
Acid Strength
CH3COOH> CH3CH2OH
CF3COOH > CH3COOH
Need to examine
inductive effect
of neighboring
atoms.
pH
pK
Ka , Kb , Kw
2H2O
+
H3O

K eq
+
OH

[ H3 O ][ OH ]

2
[ H2 O]
Keq [H2
2
O]
Kw = [H3
= [H3
+
O
+
O
][OH ]
][OH ]
where
o
-14
Kw (25 C ) = 1 x 10
in a neutral solution
[H3O+ ] = [OH-]
1 x 10-14 = [H3O+ ]2 = [OH-]2
+
[H3O
]=
[OH ]
=1x
-7
10
pX = -log X
pK = -log K
pH =
pOH =
+
-log [H3O ]
-log [OH ]
leveling effect of
+
H2O limits [H3O ]
& [OH ] to that
controlled by H2O
upper limit [H3
lower limit [H3
+
O
+
O
]=1
]=
-14
1 x 10
pH scale
acid
0
neutral
7
base
14
+
highest [H3O ] on left
lowest [H3O+ ] on right
+
[H3O
] and
[OH ]
must be considered
together
Kw = [H3
+
O
][OH ]
-log Kw = -log
+
{[H3O
][OH ]}
-log Kw =
{-log [H3O+ ]} + {-log[OH-]}
pKw = pH + pOH
but Kw = 1 x
-14
10
14 = pH + pOH
Relationship between
conjugate
acids & bases
HA + H2O
A
+ H2O
H3
+
O
HA +
+
A
OH


[ H3 O ][ A ]
Ka 
[ HA ]

[ HA ][ OH ]
Kb 

[A ]



[H3O ][ A ] [HA][ OH ]
K a  Kb 


[HA]
[A ]



[ H3 O ][ A ] [ HA ][ OH ]
Ka  Kb 

[ HA ]
[A ]
Ka x Kb = [H3O+ ][OH-] = Kw
.
Ka Kb
.
Ka Kb
=
+
[H3O
= Kw
][OH ]
= Kw
SUMMARY
+
pH = -log [H3O ]
pOH = -log [OH ]
+
-14
[H3O ][OH ] = 1 x 10
pH + pOH = 14
.
Ka Kb = Kw
Applications of
Acid-Base
Concepts
for weak acids & bases,
refer to
Appendix H for Ka &
Appendix I for Kb
values in
Kotz & Treichel
1. What is the pH of
a solution that is
0.025 M KOH?
2. What is the pH of a
0.20 M acetic acid
solution?
3. 100 mL of 0.10 M
CH3COOH are mixed
with 20.0 mL of 0.10 M
NaOH. What is the pH
of the solution?
4. Calculate the
percent ionization of
0.10 M methylamine
(CH3NH2).
pH of
Salts & Oxides
What effect does
the addition of a
salt to water have
upon the pH of the
water?
H2O equilibrium is
the prime factor in
the behavior of
solutions.
pH of a salt solution
is dependent upon
the strength of the
salt as an electrolyte.
Example 1
NaCl(s) + HOH
NaOH(aq) + HCl(aq)
Example 1
NaCl(s) + HOH  NaOH (aq) + HCl(aq)
strong base
strong acid
Na+ + OH- + H+ + Cl-  Na+ + HOH + Cl-
thus, NaCl in water
has NO effect on pH
Example 2
NaCN(s) + HOH
NaOH(aq) + HCN(aq)
Example 2
NaCN(s) + HOH 
NaOH (aq) + HCN(aq)
strong base
weak acid
 Na+ + OH- + HCN
CN
is the anion of
the weak acid HCN
CN
+ HOH
HCN +
OH
NaCN(s) + HOH Na+ + OH- + HCN
strong base
weak acid
thus, NaCN in water
produces a/n ??
solution
thus, NaCN in water
produces a BASIC
solution
5. What is the pH of a
0.010 M sodium
cyanide solution?
Example 3
NH4Cl(s) + HOH NH4OH (aq) + HCl(aq)
Example 3
NH4Cl(s) + HOH NH4OH (aq) + HCl(aq)
weak base
strong acid
NH4OH + H+ + Cl-
+
NH4
is the cation of the
weak base NH4OH
NH4+ + HOH
NH3 + H3O+
NH4Cl(s) + HOH -> NH4OH + H+ + Clweak base
strong acid
thus, NH4Cl in water
produces a/n ??
solution
thus, NH4Cl in water
produces an ACID
solution
6. What is pH of a
0.10 M ammonium
chloride solution?
Example 4
NH4CN(s) + HOH
NH4OH (aq) + HCN(aq)
weak base
weak acid
+
NH4 is the cation of
the weak base NH4OH
+
NH4 +
HOH
NH3 +
+
H3O
CN
is the anion of
the weak acid HCN
CN
+ HOH
HCN +
OH
thus, NH4CN in
water produces a/n
?? solution
The pH of a solution
formed from the cation of
a weak base and the anion
of a weak acid is
dependent on the relative
strength of the weak acid
and weak base.
-10
10
Ka(HCN) = 6.2 x
[Text: Table 5.1]
Appendix H A-23
-5
Kb(NH4OH) = 1.8 x 10
[Text: Table 5.3]
Appendix I A-25
thus, NH4CN in
water produces a/n
?? solution
thus, NH4CN in water
produces a BASIC
solution, because the
weak base is stronger
(ionizes more) than the
weak acid
Acidity of Oxides
SO2 + HOH
??
SO2 + HOH
H2SO3
[O2]
H2SO4
SO2 + HOH
H2SO3
[O2]
H2SO4
Covalent oxides are acidic &
are referred to as
acid anhydrides
Na2O + HOH
??
Na2O + HOH
2NaOH(aq)
Na2O + HOH
2NaOH(aq)
Ionic oxides are basic
& are referred to as
basic anhydrides
Lewis
Acid-Base
Theory
Acid
substance capable of
accepting an e- pair
Lewis acid
must have an empty
valence level orbital
+
H
i.e.
has an empty 1s
orbital which can accept
an e pair
+
H
Thus,
is an acid
under all three theories
Arrhenius
Bronsted-Lowry
Lewis
Lewis Acid-Base Theory
Acid: substance capable
of accepting an e pair
Base
substance capable of
donating an e pair
Examples of
Lewis bases
OH
F
, NH3 ,
all have unbonded pairs
of e available for
donation
Elements of Group 13
(3A) form compounds
that make excellent
Lewis acids
another “typical”
Lewis acid-base
reaction:
Reaction of a Lewis Acid and
Lewis Base
 New
bond formed using
electron pair from the
Lewis base.
 Coordinate covalent
bond
 Notice geometry change
on reaction.
Lewis Acids & Bases
Formation of hydronium ion is
also an excellent example.
H
+
ACID
•• ••
O—H
H
BASE
••
H O—H
H
•Electron pair of the new O-H bond
originates on the Lewis base.
Lewis Acid/Base Reaction
H3BO3 + H2O
?
H2BO3
+
+
H3O
NO!
H3BO3 + 2H2O
-
B(OH)4 + H3
+
O
is Al(OH)3 an
acid or base?
Amphoterism of Al(OH)3
Lewis Acids & Bases
This explains AMPHOTERIC
nature of some metal hydroxides.
Al(OH)3(s) + 3 H+  Al3+ + 3 H2O
Here Al(OH)3 is a Brønsted base.
Al(OH)3(s) + OH-  Al(OH)4-
Here Al(OH)3 is a Lewis acid.
Al
••
O—H
••
••
3+
Transition metal ions
also very good
Lewis Acids
Lewis Acids & Bases
Other good examples involve metal ions.
Co2+
ACID
•• ••
O—H
H
BASE
Co
2+
•• ••
O—H
H
2+
Zn
2+
Zn
..
+ HOH
..
=> [Ar]
0
4s
?
10
3d
0
4p
Zn(H2O)4
2+
Reaction of NH3 with Cu2+(aq)
Formation of complex
ions is a Lewis
acid-base reaction
Lewis Acid-Base Interactions
in Biology
 The
heme group in
hemoglobin can
interact with O2 and
CO.
 The Fe ion in
hemoglobin is a Lewis
acid
 O2 and CO can act as
Heme group
Lewis bases
Inclusiveness of the Acid/Base Definitions
Lewis
Bronstead
Arrhenius