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Self-Ionization of Water
• Even pure water conducts some electricity. This is due to
the fact that water self-ionizes:
Self ionization
of water movie
• The equilibrium constant for this process is called the ion product of water (Kw)
• At 25 °C, Kw = [H3O+][OH–] = 1.0 x 10–14 M2
⇒ Equilibrium concentrations of pure water: [H3O+] = [OH–] = 1.0 x 10–7 M
• The equilibrium constant Kw is very important because it applies to all aqueous
solutions—acids, bases, salts, and nonelectrolytes—not just to pure water.
pH and pOH
pH is a measure of the strength of an pH = –log[H3O+] and
acid; low pH = stronger acid
[H3O+] = 10–pH
pOH is a measure of the strength of
a base; low pOH = stronger base
Negative pH is also possible!
pOH = –log[OH–] and
[OH–] = 10–pOH
The pH scale extends over 14
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Ionization Constant Relationships
• pKw = –log Kw
• Kw = [H+][OH–] = 1 × 10–14 M2
⇒ at 25 ºC, pKw = 14.00
⇒ pKw = pH + pOH = 14.00
• For conjugate acid-base pairs:
4
A student determines the pH of milk of magnesia, a suspension of solid
magnesium hydroxide in its saturated aqueous solution, and obtains a
value of 10.52. What is the molarity of Mg(OH)2 in its saturated aqueous
solution? The suspended, undissolved Mg(OH)2(s) does not affect the
measurement.
Is the solution 1.0 x 10–8 M HCl acidic, basic, or neutral?
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5
6
Ordinary vinegar is approximately 1 M CH3COOH and it has a pH of about 2.4.
Calculate the expected pH of 1.00 M CH3COOH(aq), and show that the
calculated and measured pH values are in good agreement.
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7
Polyprotic Acids
• A polyprotic acid has more than one
ionizable H atom per molecule.
– Sulfuric acid, H2SO4
– Carbonic acid, H2CO3
– Phosphoric acid, H3PO4
Diprotic
Diprotic
Triprotic
• The protons of a polyprotic acid dissociate
in steps, each step having a different value
of Ka.
H2SO4
• Values of Ka decrease successively for a
given polyprotic acid:
K = 7.1⋅10-3 M
a1
Ka1 > Ka2 > Ka3 , etc.
Why?
Ka2 = 6.3⋅10-8 M
Ka3 = 4.3⋅10-13 M
H3PO4
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Calculate the following concentrations in an aqueous solution that is
5.0 M H3PO4:
(a) [H3O+] (b) [H2PO4–] (c) [HPO42–] (d) [PO43–]
What is the approximate pH of 0.71 M H2SO4?
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9
Ions as Acids and Bases
A carbonate ion accepts a proton
from water, leaving behind an OH–
and making the solution basic.
• When Na2CO3 dissolves, it completely
dissociates into ions:
Na2CO3 → 2Na+ + CO32 –
• Carbonate ion reacts to produce OH– :
CO32 – + H2O
HCO3– + OH–
• This reaction raises [OH–] above 10–7 M,
and [H3O+] decreases accordingly.
⇒ pH > 7
• We say that CO32– underwent hydrolysis.
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Ions as Acids and Bases
• Salts of strong
NaCl, KNO3
– Why?
acids and strong bases form neutral solutions:
• Salts of strong
acids and weak bases form acidic solutions: NH4NO3
– Why?
• Salts of weak acids and strong
bases form basic solutions: KNO2, NaClO
– Why?
• Salts of weak acids and weak bases form solutions that may be acidic, neutral,
or basic; it depends on the relative strengths of the cations and the anions:
NH4NO2, CH3COONH4.
– Why?
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Ions as Acids and Bases
What conclusion can you draw about the
equilibrium constants for the hydrolysis
reactions in CH3COONH4(aq)?
l
tra
u
Ne
sic
a
B
NaCl(aq)
CH3COONa(aq)
l
tra
u
Ne
c
idi
c
A
NH4Cl(aq)
CH3COONH4(aq)
* with bromthymol blue as indicator
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The Common Ion Effect
• Consider a solution of acetic acid.
• If we add acetate ion as a second solute (i.e., sodium
acetate), the pH of the solution increases:
LeChâtelier’s principle: What
happens to [H3O+] when the
equilibrium shifts to the left?
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13
The Common Ion Effect
The common ion effect is the suppression of the ionization of
a weak acid or a weak base by the presence of a common ion
from a strong electrolyte.
1. Acetic acid solution
at equilibrium: a few
H3O+ ions and a few
CH3COO– ions
2. When acetate ion
is added, and
equilibrium
reestablished: more
acetate ions, but
fewer H3O+ ions
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a) Calculate the pH of an aqueous solution that is both 1.00 M CH3COOH
and 1.00 M CH3COONa.
b) Compare the pH of a solution of only 1.00 M CH3COOH.
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Summary of Concepts
• Water undergoes limited self-ionization producing H3O+
and OH–.
• pH = –log[H3O+]
pOH = –log[OH–] pKw = –logKw
– The pH in both pure water and in neutral solutions is 7.
– Acidic solutions: pH < 7; basic solutions: pH > 7.
• In aqueous solutions at 25 oC, pH + pOH = 14.00
• Hydrolysis reactions cause certain salt solutions to be
either acidic or basic
• A strong electrolyte that produces an ion common to the
ionization equilibrium of a weak acid (or a weak base)
suppresses the ionization of the weak acid (base)
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