(aq) + +

Definitions
Arrhenius Acids and Bases
• Acids – produce hydronium ion in solution
• Bases – produce hydroxide ion in solution
Note: These definitions are specifically written for aqueous
solutions.
HA
B
-
+
H2O
H3O+
+
H2O
HB
+
+
A-
OH
Acid and Base Reactions
HCl
+
H2O
H2O
NaOH(s)
H3O+
Na+(aq)
+
+
Cl-
OH-(aq)
Table 18.1
Fig. 18.2
Hydronium Ion and Intermolecular Interactions
Acid – Base Neutralization
+
H3O
Acid
2H2O
OH
+
+
+
NaOH
NaCl
Base
Salt
H
HCl
-
+
-
H2O
OH
+
H2O
Water
Fig. 18.3
Strong
vs.
Weak
Acids
Fig. 18.4
Acid Strength
Effects [H3O+] and
Reaction Rate which
Depends upon [H3O+]
1M HCl
1M CH3COOH
HA(aq)
Zn(s)
+
+
H2O(l)
H3O+(aq)
+
A-(aq)
2H3O+(aq)
Zn+2(aq)
+
H2(g)
Strong Acids and Strong Bases
• Strong Acids
• Hydrohalic Acids: HCl, HBr, and HI
• Oxoacids:
• # of O’s exceed # acid protons (H’s) by two or more
• Strong Bases
• Soluble compounds with O2- or OH• M2O or MOH
Group 1A metals
• MO or M(OH)2
Group 2A metals
K = [Products]/[Reactants]
Law of Mass Action
(Law of Chemical Equilibrium)
aA + bB
cC + dD
c
d
[C] [D]
Qc =
=
K
c
a
b
[A] [B]
For Acid Dissociation Reaction
HA(aq)
Qc =
+
[H3O
H3O+(aq)
H2O(l)
+
(aq)
][A
(aq)
A-(aq)
+
]
[HA (aq) ] [H2O(l) ]
K c [H2O(l) ] =
Ka =
[H3O
+
(aq)
][A
[HA (aq) ]
(aq)
]
+
[H3O(aq)
][A (aq)
]
[HA (aq) ]
= Ka
Acid Dissociation
Constant
Table 18.2
Stronger Acid 
Higher [H3O+] 
Larger Ka
Smaller Ka 
Lower % HA
Dissociated 
Weaker Acid
[Acid Strength:
% Dissociation
Ka Magnitude]
Prob. 18.1
Page 764
Autoionization of Water
Ka =
K W = [H3O
+
(aq)
[H3O
+
(aq)
][OH
(aq)
]
[H2O(l) ]2
][OH
(aq)
]
Ion Product of Water
Fig. 18.5
Aqueous Solutions
+
1 x 10-14 = [H3O(aq)
][OH (aq)
]
Neutral Solution:
+
[H3O(aq)
] = [OH (aq)
] = 1 x 10-7
Prob. 18.2
Fig. 18.6
pH Scale
+
K W = [H3O(aq)
][OH (aq)
]
1 x 10
-14
= [H3O
+
(aq)
][OH
pH = - log[H3O
(aq)
+
(aq)
]
]
Fig. 18.7
Aqueous
Solution
Characteristics
pOH = - log[OH (aq)
]
Prob. 18.3
Table 18.3
pK a = - log K a
Fig. 18.9
Brønsted-Lowry Definitions
Fig. 18.10
Conjugate
Acid-Base
Pairs
Table 18.4
Table 18.5
Fig. 18.11
Base Dissociation Reactions
Conjugate Base
Conjugate Acid
Base Dissociation Constant
Conjugate
Base(aq)
+
H2O(l)
Conjugate
Acid(aq)
+
OH-(aq)
Product of [Products]
Kb =
Product of [Reactants]
[Conj. Acid][OH - ] [HA][OH - ]
Kb =
=
[Conj. Base]
[A - ]
A-(aq)
+
H2O(l)
HA(aq)
+
OH-(aq)
Table 18.6
Kb
pKb
[OH-]
pOH
A-(aq)
+
H2O(l)
HA(aq)
HA(aq)
+
H2O(l)
A-(aq)
2H2O(l)
H3O+(aq)
+
+
+
OH-(aq)
OH-(aq)
Kb
H3O+(aq)
Ka
Ka x Kb = Kw
Fig. 18.12
Acid – Base
Strength
and
the
Periodic Table
Fig. 18.13
Salts of Weak Acids and Weak Bases
Consider Cations and Anions
Table 18.8
Lewis Acids and Bases
Lewis Acid – electron pair acceptor
Lewis Base – electron pair donor
Page 793
Metals act as Lewis acids in water
Equilibria of Acid-Base Buffer
Systems
Buffer: resists a change in pH upon addition of Acid or Base
conjugate acid-base pair of a weak acid
HA(aq)
+
Ka =
A-(aq)
H2O(l)
[H3O
+
(aq)
][A
[HA (aq) ]
(aq)
]
+
H3O+(aq)
Acid Dissociation
Constant
Table 19.1
Common Ion Effect
Fig. 19.2
HA(aq)
H2O(l)
+
HC2H3O2(aq)
+
H2O(l)
A-(aq)
+
C2H3O2-(aq)
H3O+(aq)
+
H3O+(aq)
Henderson-Hasselbalch Equation
Ka =
[H3O
+
(aq)
][A
(aq)
]
[HA (aq) ]
[A
]
+
-log K a = -log [H 3O ] - log
[HA]
[A - ]
pK a = pH - log
[HA]
-
[A ]
pH = pK a + log
[HA]
Definition of pKa for a buffer
solution.
pKa = pH
when
[conjugate acid] = [conjugate base]
Prob. 19.2
How to prepare a buffer at a specific pH:
-
[A ]
pH = pK a + log
[HA]
Calculate the proportions of NaH2PO4 and Na2HPO4 needed to
Prepare a buffer solution of pH=7.00. The pKa of H2PO4- is 7.21
 HPO4−2 
= 7.21 + log
7.00
 H 2 PO4−1 
 HPO4−2 
=
log
7.00 − 7.21
−1
 H 2 PO4 
take antilog of both sides:
 HPO4−2 
0.62
=
1
 H 2 PO4−1 
Use 0.62 mol of Na2HPO4
for every mol of NaH2PO4
in any total amount.
Fig. 19.4
Acid-Base Indicators
Color changes at pKa values.
Fig. 19.6
Titration of a Strong Acid with a Strong Base
Fig. 19.7
Titration of a
Weak Acid
with a
Strong Base
Fig. 19.9
Titration of a Weak Polyprotic Acid with a Strong Base