Ch. 19 – Acids & Bases
I. Introduction to
Acids & Bases
A. Properties

electrolytes
 electrolytes

sour taste

bitter taste

turn litmus red

turn litmus blue

react with metals
to form H2 gas

slippery feel

vinegar, milk, soda,
apples, citrus fruits

ammonia, lye,
antacid, baking soda
B. Definitions
 Arrhenius
• Acids contain hydrogen
• Acids form hydronium ions (H3O+) in
aqueous solution
HCl + H2O 
+
H3O
H
H
Cl
acid
O
H
H
–
+
O
H
+
–
Cl
Cl
H
B. Definitions
 Arrhenius
• Bases contain a hydroxide group
• Bases form hydroxide ions (OH-) in
aqueous solution
H 2O
NaOH 
base
+
Na
+
OH
B. Definitions
 Brønsted-Lowry
• Acids are proton (H+) donors
• Bases are proton (H+) acceptors
HCl + H2O 
acid
–
Cl
+
+
H3O
base
conjugate base
conjugate acid
B. Definitions
 Brønsted-Lowry
• Conjugate Acids are the result after
a base accepts a hydrogen ion
• Conjugate Bases are the result after
an acid donates a hydrogen ion
HBr + NaOH  NaBr + H2O
acid
base
conjugate base
conjugate acid
B. Definitions
H2O + HNO3  H3O+ + NO3–
B
A
CA
CB
H2O + NH3  NH4+ + OHA
B
 Amphoteric
CA
CB
– can be an acid or a base
B. Definitions

Give the conjugate base for each of the following:
-
HF
F
H3PO4
H2PO4
+
H3O
H2O
 Polyprotic
– an acid with more than one H+
B. Definitions

Give the conjugate acid for each of the following:
Br
-
HBr
HSO4
H2SO4
2CO3
HCO3
C. Strength
-
+
 Strong
Acid/Base
• 100% ionized in water
• strong electrolyte
HCl
HNO3
H2SO4
HBr
HI
HClO4
NaOH
KOH
KOH
RbOH
CsOH
Ca(OH)2
Ba(OH)2
C. Strength
 Weak
Acid/Base
• does not ionize completely
• weak electrolyte
HF
CH3COOH
H3PO4
H2CO3
HCN
-
+
NH3
Ch. 19 – Acids & Bases
II. pH
(p. 644 – 658)
pH water equilibrium
 Pure
water ionizes to a small extent to
produce hydrogen ions and hydroxide ions
 According to LeChatlier’s principle if an acid
is dissolved in water the equilibrium will shift
to the left decreasing the hydroxide ion
concentration.
 If a base is dissolved in water this
decreases the hydrogen ion concentration.
A. Ionization of Water
H2O (l)+ H2O
(l)
H3
+
O (aq)
Self-Ionization of Water
+
OH (aq)
A. Ionization of Water
Kw = [H3O+][OH-] = 1.0  10-14
Ion Product Constant for Water
• The ion production of water, Kw = [H3O+][OH–]
• Pure water contains equal concentrations of H+
and OH– ions, so [H3O+] = [OH–]
• For all aqueous solutions, the product of the
hydrogen-ion concentration and the hydroxide-ion
concentration equals 1.0 x 10-14
A. Ionization of Water
 Find
the hydroxide ion concentration of
3.0  10-2 M HCl.
HCl
→
H+
+
Cl3.0  10-2M
3.0  10-2M
[H3O+][OH-] = 1.0  10-14
[3.0  10-2][OH-] = 1.0  10-14
[OH-] = 3.3  10-13 M
A. Ionization of Water
 Find
the hydronium ion concentration of
1.4  10-3 M Ca(OH)2.
Ca(OH)2 →
Ca2+
+
2 OH1.4  10-3M
2.8  10-3M
[H3O+][OH-] = 1.0  10-14
[H3O+][2.8  10-3] = 1.0  10-14
[H3O+] = 3.6  10-12 M
B. pH Scale
14
0
7
INCREASING
ACIDITY
NEUTRAL
pH =
INCREASING
BASICITY
+
-log[H3O ]
pouvoir hydrogène (Fr.)
“hydrogen power”
B. pH Scale
pH of Common Substances
B. pH Scale
pH =
+
-log[H3O ]
pOH =
-log[OH ]
pH + pOH = 14
B. pH Scale
 What
is the pH of 0.050 M HNO3?
pH = -log[H3O+]
pH = -log[0.050]
pH = 1.30
Acidic or basic? Acidic
B. pH Scale
 What
is the pH of 0.050 M Ba(OH)2?
[OH-] = 0.100 M
pOH = -log[OH-]
pOH = -log[0.100]
pOH = 1.00
pH = 13.00
Acidic or basic?
Basic
B. pH Scale
 What
is the molarity of HBr in a solution
that has a pOH of 9.60?
pH + pOH = 14
pH = -log[H3O+]
pH + 9.60 = 14
4.40 = -log[H3O+]
pH = 4.40
-4.40 = log[H3O+]
Acidic
[H3O+] = 4.0  10-5 M HBr