Transcript Chapter Nine

Acids, Bases, and Salts
Acids
Taste sour, are corrosive
React with bases (alkalies)
Bases
Taste bitter, feel slippery
React with acids
Salts
Are formed in reactions between acids and
bases
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9a–1
9.1 Arrhenius Acid-Base Theory
Acids produce hydrogen ions, H1+,
in water solution.
Acids are molecular compounds
(not ionized) when pure.
Common acids:
HCl(aq), H2SO4, H3PO4, HNO3, HC2H3O2
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9a–2
9.1 Arrhenius Acid-Base Theory
Bases produce hydroxide ions, OH1–, in
water solution.
Bases are ionic compounds when pure
Common bases
NaOH, KOH
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9a–3
Figure 9.1
Ionization (Arrhenius acids)
Dissociation (Arrhenius bases).
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9a–4
Figure 9.2 Litmus: a vegetable dye obtained from
lichens. Paper treated with this dye turns
from blue to red in acids (left) and from
red to blue in bases (right).
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9a–5
9.2 BrØnsted-Lowry
Acid-Base Theory
An acid can donate a proton, H1+, to another
substance.
An acid is a proton donor
Reactions of acids:
HCl(aq) + H2O  H3O1+(aq) + Cl1(aq)
H3O1+(aq) + NH3(aq)  H2O + NH41+(aq)
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9a–6
9.2 BrØnsted-Lowry
Acid-Base Theory
A base can accept a proton, H1+, from another
substance.
A base is a proton acceptor
Reactions of acids:
H2O + HCl(aq)  H3O1+(aq) + Cl1(aq)
NH3(aq) + H3O1+(aq)  NH41+(aq) + H2O
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9a–7
9.2 BrØnsted-Lowry
Acid-Base Theory
Conjugate acid-base pairs:
A reaction between and acid and a base produces
a conjugate acid and a conjugate base
HCl(aq) + H2O  H3O1+(aq) + Cl1(aq)
Acid
Base
Conjugate
Acid
Conjugate
Base
H3O1+(aq) + NH3(aq)  H2O + NH41+(aq)
Acid
Base
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Conj.
Base
Conj.
Acid
9a–8
9.2 BrØnsted-Lowry
Acid-Base Theory
Choose the acid, base, conjugate acid, and
conjugate base:
HCOOH(aq) + NH3  HCOO1(aq) + NH41+ (aq)
H2PO41(aq) + H2O  HPO42(aq) + H3O1+ (aq)
HPO42(aq) + H2O  PO43(aq) + H3O1+ (aq)
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9a–9
9.2 BrØnsted-Lowry
Acid-Base Theory
Amphoteric substances can act as both acids
and bases:
HCOOH(aq) + H2O  HCOO1(aq) + H3O1+ (aq)
NH3 (aq) + H2O  NH41+ (aq) + OH1(aq)
HPO42(aq) + OH1(aq)  PO43(aq) + H2O
HPO42(aq) + H3O1+(aq)  H2PO41(aq) + H2O
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9a–10
9.3 Mono-, Di- and Triprotic
Acids
Monoprotic acids can transfer one proton
CH3COOH + H2O  CH3COO1 + H3O1+
Diprotic acids can transfer two protons
H2CO3 + H2O  HCO31 + H3O1+
HCO31 + H2O  CO31 + H3O1+
The first proton transfer is complete before the
second one starts.
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9a–11
9.3 Mono-, Di- and Triprotic
Acids
Triprotic acids can transfer three protons
H3PO4 + H2O  H2PO41 + H3O1+
H2PO41 + H2O  HPO42 + H3O1+
HPO42 + H2O  PO43 + H3O1+
The first proton transfer is complete before the
second one starts. The second proton transfer
is complete before the third one starts.
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9a–12
9.4 Strength of Acids and Bases
Acids differ in the extent of ionization when
they are put in solution
A few acids, the strong acids, ionize completely.
Most acids, weak acids, do not ionize completely.
The equilibrium constant, Ka, is a measure of the
strength of an acid.
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9a–13
The Strong Acids
Formula
HCl(aq)
HBr(aq)
HI(aq)
HNO3
HClO4
H2SO4
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Name
Hydrochloric acid
Hydrobromic acid
Hydriodic acid
Nitric acid
Perchloric acid
Sulfuric acid*
*first proton
9a–14
Some Weak Acids
Formula
HSO41
C 9H 8O 4
HCOOH
HC3H5O3
CH3COOH
H2CO3
H2S(aq)
HCN(aq)
C6H5OH
Name
Hydrogen sulfate
Acetylsalicylic acid
Formic acid
Lactic acid
Acetic acid
Carbonic acid
Hydrosulfuric acid
Hydrocyanic acid
Phenol
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Ka
1.2 x 102
3.0 x 104
1.8 x 104
1.4 x 104
1.8 x 105
4.3 x 107
1.0 x 107
4.9 x 1010
1.3 x 1010
9a–15
Figure 9.5 A comparison of the number of H3O1+
ions present in strong acid and weak
acid solutions of equal concentration.
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9a–16
The Strong Bases
Group 1A Hydroxides
LiOH
NaOH
KOH
RbOH
CsOH
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Group 2A Hydroxides
Ca(OH)2
Sr(OH)2
Cs(OH)2
9a–17
The Weak Base, Ammonia
NH3 + H2O  NH41+ + OH1
Kb = 1.8 x 105
NH41+ + H2O  NH3 + H3O1+
Ka = 5.6 x 1010
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9a–18
9.5 Salts
A salt is a compound containing a metal or
polyatomic cation, and a nonmetal or
polyatomic anion (except OH1).
NaCl, NH4Cl, BaSO4, CaCO3, Al2(SO4)3
Salts are formed in neutralization reactions
between acids and bases.
HCl(aq) + NaOH(aq)  H2O + NaCl(aq)
2 Al(OH)3(s) + 3 H2SO4(aq) 
6 H2O + Al2(SO4)3 (aq)
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9a–19
9.6 Acid-Base Neutralization
Reactions
Neutralization reactions between an acid and
a base produce a salt and water
HCl(aq) + NaOH(aq)  H2O + NaCl(aq)
2 Al(OH)3(s) + 3 H2SO4(aq) 
6 H2O + Al2(SO4)3 (aq)
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9a–20
H2SO4(aq) + Ba(OH)2(aq)  BaSO4 (s) + 2 H2O
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9a–21
Figure 9.7 Formation of water by the transfer of
protons from H3O1+ ions to OH1 ions.
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9a–22
9.7 Self-Ionization of Water
2 H2O  H3O1+ + OH1
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Kw = 1.0 x 1014
9a–23
9.7 Self-Ionization of Water
2 H2O  H3O1+ + OH1
Kw = 1.0 x 1014
Kw = ion product constant for water
Kw = 1.0 x 1014 = [H3O1+] [OH1]
In pure water, [H3O1+] = [OH1] = 1.0 x 107M
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9a–24
Figure 9.9
The relationship between [H3O1+] and
[OH1] in aqueous solution is an inverse
proportion; when [H3O1+] is increased,
[OH1] decreases, and vice versa.
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9a–25
9.8 The pH Concept
[H3O1+] can vary over a wide range, and is
often low. Often, you need scientific
notation to express it. This isn't always
convenient.
A simpler way to write [H3O1+] is pH
pH = log [H3O1+]
[H3O1+] = 10pH
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9a–26
9.8 The pH Concept
Give the pH for
[H3O1+] =
=
=
=
=
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0.010 M
4.2 x 103 M
1.0 x 107 M
6.8 x 1010 M
1.0 x 1012 M
9a–27
9.8 The pH Concept
Give [H3O1+] for
pH =
=
=
=
=
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3.00
4.50
6.85
7.00
10.75
9a–28
Figure 9.10
Most fruits and vegetables
are acidic. Tart or sour
taste is an indication that
such is the case.
Nonintegral pH values for
selected foods are as
shown here.
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9a–29
Figure 9.11
Relationships among pH
values, [H3O1+], and
[OH1] at 24º C.
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9a–30
Figure 9.12
The pH values of selected
common liquids. The
lower the numerical value
of the pH, the more acidic
the substance.
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9a–31
Table 9.5
The Normal pH Range of Selected Body
Fluids.
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9a–32
Figure 9.13
A pH meter gives an accurate measurement of pH values. The pH of vinegar is
2.32 (left). The pH of milk of magnesia in
water is 9.39 (right).
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9a–33
Recognizing Acids, Bases, and Neutral
Solutions (III-3)
A solution is acidic if
[H3O1+] > 1.0 x 107
Ph < 7.00
A solution is basic if
[H3O1+] < 1.0 x 107
Ph > 7.00
A solution is neutral if
[H3O1+] = 1.0 x 107
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Ph = 7.00
9a–34
9.9 Buffers
A buffer is a solution that resists major changes
in pH when acids or bases are added to it.
A buffer contains
A weak acid to react with added base
A weak base to react with added acid
Most often, the acid and base are conjugate pairs
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9a–35
Table 9.6
A Comparison of pH Changes in Buffered
and Unbuffered Solutions.
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9a–36
(a)
(b)
The buffered solution on the left and the
unbuffered solution on the right are both at pH 8.
After the addition of 1 mL of 0.01 M HCl solution,
the pH of the buffered solution has not changed
much. The unbuffered solution has become
acidic, as shown by the change in the color of the
acid-base indicator.
Source: Ken O’Donoghue
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9a–37
9.9 Buffers
A buffer made of equimolar amounts of a weak
acid and its conjugate base will have a pH
equal to log Ka.
log Ka = pKa
Adding a little acid will shift the pH of the buffer
down, adding a little base will shift the pH of
the buffer up.
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9a–38
9.9 Buffers
What is the pH of a buffer made of 0.10 mole
of CH3COOH and 0.10 mole of NaCH3COO?
What is the pH of a buffer made with 1.0 mole
of ammonia and 1.0 mole of ammonium
chloride?
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9a–39
9.9 Buffers
A buffer is made with 1.0 mole of CH3COOH
and 0.10 mole of NaCH3COO.
What component of the buffer reacts with
added H3O1+?
What component of the buffer reacts with
added OH1?
(III-4)
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9a–40
9.9 Buffers
A buffer is made with 1.0 mole of NH3 and 1.0
mole of NH4Cl.
What component of the buffer reacts with
added H3O1+?
What component of the buffer reacts with
added OH1?
(III-4)
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9a–41
9.10 Acid-Base Titrations
In an acid-base titration, a measured
volume of an acid or base of known
concentration is reacted with a
measured volume of a base or acid of
unknown concentration.
The reaction is conducted in a way that
exactly equimolar amounts of H3O1+ and
OH1 are combined.
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9a–42
Figure 9.15
A schematic diagram
showing the setup used
for titration procedures.
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9a–43
Figure 9.16
An acid-base titration using an indicator
that is yellow in acidic solution and red
in basic solution.
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9a–44
9.10 Acid-Base Titrations
A student titrates 25.00 mL of vinegar
(acetic acid in water) with 15.85 mL of
0.1048 M NaOH. What is the concentration of acetic acid in the vinegar?
CH3COOH(aq) + NaOH(aq) 
CH3COONa(aq) + H2O
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9a–45