13-1
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Questions for Consideration
1.
2.
3.
4.
5.
6.
13-2
How do acids and bases differ from other
substances?
How do strong and weak acids differ?
How can we compare the different strengths
of weak acids?
What causes aqueous solutions to be acidic
or basic?
How is pH related to acidity and basicity?
What is a buffered solution?
Topic Outline
1.
2.
3.
4.
5.
6.
13-3
Properties and Definitions of Acids and Bases
Strong and Weak Acids and Bases
Acid Ionization Constants
Acidic, Basic, and Neutral Solutions
The pH and pOH Scale
Buffered Solutions
Acids and Bases
13-4
Acids and Bases: Properties
Acids and bases have properties that differ from other
substances:
 Acids taste sour.
 Bases taste bitter and feel slippery.
 Both change the color of some dyes.
 Acids cause many metals to corrode.
 Acids and bases combine to neutralize each other.
13-5
CAUTION: Do not taste laboratory chemicals.
Acids and Bases: Arrhenius
 1800’s Arrhenius Model
 An acid in water produces hydrogen (H+) ions.
 A base in water produces hydroxide (OH) ions.
HCl(g)  H+(aq) + Cl(aq)
NaOH(s)  Na+(aq) + OH(aq)
 Arrhenius earned the Nobel prize for his work
that showed that H+(aq) and OH(aq) ions are
important in acid-base chemistry.
13-6
Acids and Bases: Arrhenius
 Problem with Arrhenius
Model:
 H+ does not exist completely
free in aqueous solution. It
associates strongly with
other water molecules.
 Chemists recognize this by
representing an aqueous H+
ion as H3O+(aq), the
hydronium ion.
13-7
Figure 13.4
Acids and Bases: Bronsted Lowry
 1923 Brønsted-Lowry definitions:
 An acid donates an H+ ion to another substance.
 A base accepts an H+ from another substance.
HCl(g) + H2O(l)  H3O+(aq) + Cl(aq)
NH3(aq) + H2O(l)  NH4+(aq) + OH(aq)
 In an acid-base reaction, the acid donates (transfers)
an H+ to a base.
13-8
Acids and Bases: Bronsted Lowry
HCl(g) + H2O(l)  H3O+(aq) + Cl(aq)
Acid
base
NH3(aq) + H2O(l)  NH4+(aq) + OH(aq)
base
acid
13-9
Sample Problem
For each reaction, identify the Brønsted-Lowry
acid and base reactants.
1. OCl-(aq) + H2O(l)  HOCl(aq) + OH-(aq)
2. H2SO4(aq) + F-(aq) → HSO4-(aq) + HF(aq)
3. NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq)
13-10
Sample Problem
For each reaction, identify the Brønsted-Lowry
acid and base reactants.
1. OCl-(aq) + H2O(l)  HOCl(aq) + OH-(aq)
2. H2SO4(aq) + F-(aq) → HSO4-(aq) + HF(aq)
3. NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq)
 in red bold letters – acid
 Other reactant - base
13-11
Acids and Bases: Bronsted Lowry
 Many anions from dissolved ionic compounds act
as bases:
Na2CO3(s)  2Na+(aq) + CO32(aq)
 The carbonate ion acts as a base in water:
CO32(aq) + H2O(l)  HCO3(aq) + OH(aq)
base
acid
 and with other acids:
HF(aq) + CO32(aq)  F(aq) + HCO3(aq)
acid
base
13-12
Acids and Bases: Bronsted Lowry
NH3(g) + HCl(g)  NH4+(g) + Cl(g)
Base
acid
Reverse reaction
NH4+(g) + Cl(g)  NH3 (g) + HCl(g)
acid
base
The reverse reaction is
also an acid base reaction!
13-13
Acids and Bases: Bronsted Lowry
Imagine the following reaction going in the
reverse direction. What would be the
acid and what would be the base?
HF(aq) + CO32(aq)  F(aq) + HCO3(aq)
Acid1
base1
base2 acid2
13-14
We call the acid and base products the
conjugates of the base and acid that formed
them.
Acids and Bases: Bronsted Lowry
HF(aq) + CO32(aq)  F(aq) + HCO3(aq)
acid
base
conjugate
base
conjugate
acid
A conjugate acid-base pair differs only by
one proton.
13-15
Sample Problem
Determine the formula of the conjugate base of
each acid.
Acid
Conjugate Base
HCO3H2CO3
H2PO4H3PO4
HPO42PO43NH4+
NH3
H2O
OH13-16
Sample Problem
Determine the formula of the conjugate acid of
each base.
Base
Conjugate Acid
HSO4
SO42
H2CO3
HCO3
NH2
NH3
ClO2
HClO2
H2PO4 
H3PO4
13-17
Sample Problem
Identify the acid and base reactants and
their conjugate acid and base:
HCO3- + H3PO4  H2CO3 + H2PO4base
acid
Conjugate acid-base pairs
= H3PO4 and H2PO4= HCO3- and H2CO3
13-18
conjugate conjugate
acid
base
Acids and Bases: Bronsted Lowry
Amphoteric substance



A substance that can act as either an acid or a base
Water is the most common amphoteric substance.
Another common amphoteric substance is the
bicarbonate ion, HCO3-:
HCO3-(aq) + OH-(aq) → CO32-(aq) + H2O(l)
acid
base
conjugate conjugate
base
acid
HCO3-(aq) + H3O+(aq) → H2CO3(aq) + H2O(l)
base
13-19
acid
conjugate conjugate
acid
base
Worksheet # 6- 5
Identify the conjugate acid – base pairs in the
following reactions.
1. H2SO4 + NH3 → NH4+ + HSO4-
2. SO42- + HClO3 → ClO3- + HSO43. OH- + H3PO4 + → H2PO4- + H2O
4. Nitric acid + hydroxide ion → nitrate ion + water
5. Carbonic acid + flouride ion → bicarbonate ion +
hydroflouric acid
13-20
Worksheet # 6- 5: Answers
Identify the conjugate acid – base pairs in the
following reactions.
1. H2SO4 + NH3 → NH4+ + HSO4-
2.
3.
4.
5.
SO42- + HClO3 → HSO4- + ClO3OH- + H3PO4 → H2PO4- + H2O
HNO3 + OH- → NO3- + H2O
H2CO3 + F- → HF + HCO3-
In green – acid-conjugate base
In red = base – conjugate acid
13-21
Strong and Weak Acids and Bases
Strong and weak acids and bases differ in
the extent of ionization.
 Strong acids ionize completely.
 Weak acids and bases ionize to only a small
extent – a small fraction of the molecules
ionize.
13-22
Strong and Weak Acids and Bases
Strong acids are strong electrolytes.
13-23
Figure 13.6
Strong and Weak Acids and Bases
13-24
Strong and Weak Acids and Bases
Strong bases are strong electrolytes.
13-25
Figure 13.7
Strong and Weak Acids and Bases
13-26
Strong and Weak Acids and Bases
Weak acids are weak electrolytes
13-27
Strong and Weak Acids and Bases
13-28
Strong and Weak Acids and Bases
Weak bases are weak electrolytes
13-29
Strong and Weak Acids and Bases
13-30
Strong and Weak Acids and Bases
HCl(aq)
13-31
HF(aq)
Sample Problem
Which of the following are weak acids and weak bases?
A. weak base
B. weak acid
C. strong acid
13-32
Sample Problem
One of the diagrams below represents HClO4, and
the other represents an aqueous solution of
HSO4-. Which is which? Explain your reasoning.
Figure from p. 513
13-33
HClO4(aq)
HSO4-(aq)
Sample Problem
Identify each of the following as a strong acid,
weak acid, strong base, or weak base.
1. HI(aq)
2. NaCH3CO2(aq)
3. NH4+(aq)
4. NH3(aq)
13-34
Strong acid
Weak base
Weak acid
Weak base
Worksheets # 6 - 6
Classify each of the following as strong acids and bases
and weak acids and bases.
1. HI
2. Acetic acid
3. Nitrite ion
4. H2S
5. Magnesium hydroxide
6. Perchloric acid
7. Sodium hydroxide
8. Citric acid
9. Sulfate ion
10. bisulfate ion
13-35
Worksheets # 6 - 6 : Answers
Classify each of the following as strong acids and bases
and weak acids and bases.
1. HI - strong acid
2. Acetic acid – weak acid
3. Nitrite ion – weak base
4. H2S – weak acid
5. Magnesium hydroxide – strong base
6. Perchloric acid – strong acid
7. Sodium hydroxide – strong base
8. Citric acid – weak acid
9. Sulfate ion – weak base
10. bisulfate ion - amphoteric
13-36
Acid Ionization Constants
 Acid strength depends on the relative number
of acid molecules that ionize when dissolved in
water – the degree of ionization.
HX →
+
H
+
X
Keq = [products] = [H+] [X-]

13-37
[reactants]
[HX]
Acid Ionization Constants
• If the value of Keq is larger, then more
products exist at equilibrium and the acid has
a larger percentage of molecules which have
been ionized.
Keq = [products]
[reactants]
• If we are dealing with acids, we call the Keq as
Ka
Keq = Ka
13-38
Acid Ionization Constants
The strengths of weak acids can be compared by
examining the equilibrium constants for their
reactions with water, which we label Ka.
HCN(aq) + H2O(l) → H3O+(aq) + CN-(aq)
+
-
[H 3O ][CN ]
-10
Ka =
= 6.2  10
[HCN]
13-39
Acid Ionization Constants
CH3CO2H(aq) + H2O(l) → H3O+(aq) + CH3CO2-(aq)
[H 3O+ ][CH 3CO2- ]
-5
Ka =
= 1.8  10
[CH 3CO2 H]
13-40
Acid Ionization Constants
 The acid ionization constant, Ka, describes
the equilibrium that forms when an acid
reacts with water.
 The larger the Ka value, the stronger the
acid.
13-41
Acid Ionization Constants
• When using Ka values to determine
the strengths of conjugate acids and
bases, use this rule of thumb:
The stronger the acid, the weaker the
conjugate base.
13-42
Acid Ionization Constants
 Which is the larger value, 4.0  10-10 or 1.0  10-6? Which
would be found for the weaker acid?
13-43
Acid
Ionization
Constants
13-44
Sample Problem
Which solution has the greater
concentration of H3O+?
• 0.10 M HOCl Ka = 4.0 × 10-8
• 0.10 M HCN Ka = 6.2 × 10-10
We have equal concentrations of each acid, so we just
need to determine which acid is stronger.
Since HOCl has the larger Ka, it is the stronger acid
and will therefore have the greater concentration
of H3O+ ion.
13-45
Acid Ionization Constants
Polyprotic acids
 An acid that contains more than
one acidic hydrogen and can
thus donate more than one H+
ion
 The acid donates one H+ ion at a
time in steps
 The Ka values for polyprotic acids
are often labeled to indicate the
particular step in the overall
ionization process (Ka1, Ka2, Ka3,
etc.)
13-46
Acid Ionization Constants
H3PO4 + H2O → H3O+ + H2PO4 -
Ka1
H2PO4- + H2O → H3O+ + HPO4 -2
Ka2
HPO4 -2 + H2O → H3O+ + PO4 -3
13-47
Ka3
Acid Ionization Constants
 If an acid has more
than one hydrogen
atom, we need to
determine which
hydrogen atoms are
acidic.
 In oxoacids, the
acidic hydrogen
atoms are bonded to
oxygen atoms:
13-48
Figure 13.5
Acid Ionization Constants
13-49
Sample Problem
Oxalic acid, H2C2O4, occurs in plants and foods
such as parsley, rhubarb, almonds, and
green beans.
1. Write equations that show the ionization of
oxalic acid in water.
2. Besides water, which ion or molecule has the
highest concentration in solution when oxalic
acid is added to water?
13-50
Sample Problem
Write equations that show the ionization of oxalic acid in water.
H2C2O4(aq) + H2O(l) → H3O+(aq) + HC2O4-(aq)
Ka1 = 5.6 × 10-2
HC2O4- (aq) + H2O(l) → H3O+(aq) + C2O42-(aq)
Ka2 = 1.5 × 10-4
2. Besides water, which ion or molecule has the highest
concentration in solution when oxalic acid is added to water?
The Ka1 value for H2C2O4 shows that it is a weak acid and
ionizes only to a small extent. That means that most of the
oxalic acid is present in water as H2C2O4 and it is present in the
highest concentration.
1.
13-51
Worksheet # 6 - 7
Which of the following are stronger acids/bases based
on the acid ionization constants.
1) 2.5 M HCl or 2.5 M acetic acid
2) 0.67 M HNO2 and 0.67 M HCN
3) 0.004 M HF and 0.004 M HOCl
4) 1.15 M sulfurous acid and 1.15 M hydrosulfuric acid
(first ionization)
5) 0.43 M carbonic acid and 0.43 M oxalic acid (first
ionization)
6) Acetate ion and flouride ion
7) Nitrite ion and sulfite ion
8) Formate ion and benzoate ion
13-52
Worksheet # 6 - 7: Answers
Which of the following are stronger acids/bases based
on the acid ionization constants.
1) 2.5 M HCl or 2.5 M acetic acid
2) 0.67 M HNO2 and 0.67 M HCN
3) 0.004 M HF and 0.004 M HOCl
4) 1.15 M sulfurous acid and 1.15 M hydrosulfuric acid
(first ionization)
5) 0.43 M carbonic acid and 0.43 M oxalic acid (first
ionization)
6) Acetate ion and flouride ion
7) Nitrite ion and sulfite ion
8) Formate ion and benzoate ion
13-53
Acidic, Basic, and Neutral Solutions
Water self-ionizes to a very small extent:
H2O(l) + H2O(l) → H3O+(aq) + OH(aq)
The product of [H3O+] and [OH] is a constant
at a given temperature:
Kw = [H3O+][OH]
13-54
Acidic, Basic, and Neutral Solutions
At 25C, the value of Kw is 1.0 × 1014, so in
a neutral solution (as well as in pure
water):
[H3O+] = [OH] = 1.0 × 107
Kw = [H3O+][OH] = 1.0 × 1014
13-55
Acidic, Basic, and Neutral Solutions
Neutral Solution: [H3O+] = [OH]
Acidic Solution: [H3O+] > [OH]
Basic Solution: [OH] > [H3O+]
 Why would there be hydronium ions in a basic
solution?
 Why would there be hydroxide ions in an acidic
solution?
 Why would either be present in a neutral solution?
13-56
Acidic, Basic, and Neutral Solutions
 In an acidic solution, there is excess
H3O+, so the OH us very low.
 In a basic solution, there is excess OH,
so the H3O+ concentration is very low.
Neutral Solution: [H3O+] = [OH]
Acidic Solution: [H3O+] > [OH]
Basic Solution: [OH] > [H3O+]
13-57
Acidic, Basic, and Neutral Solutions
13-58
Acidic, Basic, and Neutral Solutions
 When the H3O+ concentration is known:
14
K
1.0  10

w
[OH ] =
=
+
+
[H3O ]
[H3O ]
 When the OH concentration is known:
14
Kw
1.0  10
[H3O ] =
=

[OH ]
[OH  ]
+
13-59
Sample Problem
Fill in the blanks.
[H3O+]
[OH-]
10-5 M
10-12 M
10-6 M
10-12 M
13-60
Sample Problem
Fill in the blanks.
[H3O+]
13-61
[OH-]
10-5 M
10-9 M
10-12 M
10-2 M
-8
-6
10 M
10 M
10-2 M
10-12 M
Sample Problem
 Given the concentration of OH- ion in each
solution, calculate the concentration of
H3O+ in that solution. Identify each
solution as acidic, basic, or neutral.
1. [OH-] = 1.0 × 10-8 M
2. [OH-] = 0.010 M
13-62
Sample Problem
-14
K
1.0

10
-6
+
w
[H 3O+ ] =
=
=
1.0

10
M
H
O
3
[OH - ] 1.0  10-8
Acidic
-14
K
1.0

10
-12
+
w
[H 3O+ ] =
=
=
1.0

10
M
H
O
3
[OH - ] 1.0  10-2
Basic
13-63
Sample Problem
What is the [H3O+] and [OH] for each
solution?
 0.10 M HCl
[H3O+] = 0.10 M, [OH] = 1.0 × 10-13 M
 0.10 M NaOH
[H3O+] = 1.0 × 10-13 M, [OH] = 0.10 M
 0.00010 M HCl
[H3O+] = 1.0 × 10-4 M, [OH] = 1.0 × 10-10 M
13-64
Worksheet # 6 - 8
Complete the following table.
Solution
[H3O+]
[OH-]
0.25 M HI
2.55 x 10 – 3 M
9.58 x 10 – 10 M
0.071 M LiOH
1.0 x 10 – 7 M
HOH
13-65
Acidic/basic/
neutral
Worksheet # 6 – 8: Answers
Complete the following table.
Solution
[H3O+]
[OH-]
Acidic/basic/
neutral
0.25 M HI
0.25 M
4.0 x 10-14 M
acidic
2.55 x 10 – 3 M
3.92 x 10-12 M
acidic
1.04 x 10-5 M
9.58 x 10 – 10 M
acidic
1.41 x 10-13 M
0.071 M
basic
1.0 x 10 – 7 M
HOH
1.0 x 10-7 M
neutral
0.071 M LiOH
13-66
pH
The pH of a solution is defined as the
negative logarithm (base 10) of the
H3O+ concentration:
pH = log[H3O+]
13-67
pH
What is the pH of solutions with the following
hydronium ion concentrations?
1.
2.
3.
4.
[H3O+] = 101 M
[H3O+] = 105 M
[H3O+] = 107 M
[H3O+] = 1011 M
pH = 1
pH = 5
pH = 7
pH = 11
Which are acidic and which are basic?
13-68
The pH Scale
 Most aqueous solutions
have pH values that
range from 1-14.
 However, some
concentrated solutions
fall outside of this range.
 What is the pH of a 1.0 M HCl
solution?
13-69
Figure 13.13
The pH Scale
By what factor does [H3O+]
change when the pH
changes
 from 8.0 to 9.0?
 from 4.0 to 2.0?
13-70
Figure 13.13
Sample Problem
The hydronium ion concentration in a soil
sample was determined to be 5.0 × 107 M.
1. What is the pH of the soil? pH = 6.30
2.
13-71
Is the soil acidic or basic?
acidic
Sample Problem
What is the pH of each of the following
solutions? Once you’ve done the
calculation, check your answer to make
sure it makes sense.
pH = 3.07
1. 0.00085 M HCl
2. 0.010 M NaOH
pH = 12.00
3. 1.0 M HNO3
pH = 0.00
13-72
pOH
 pOH is defined as the negative logarithm
(base 10) of hydroxide ion concentration,
[OH-]:
pOH = - log [OH-]
The relationship between pH and pOH is:
pH + pOH = 14
13-73
pOH
[H3O+] = 10pH
[OH] = 10pOH
 What is the H3O+
concentration in a soil
sample if the pH is
measured to be 6.20?
6.3 × 10-7 M
13-74
Figure 13.14
pOH
13-75
Figure 13.15
pH and pOH
 Examine how
these quantities
are related.
13-76
Sample Problem
When the pH of water in a lake falls below about
4.5, the lake may be considered dead because
few organisms can survive in such an acidic
environment. What is the pH of a lake that has
an OH- concentration of 1.0 × 10-9 M? Would
this lake be considered dead?
If the pH falls below 4.5, then the lake is dead.
pH = -log [H3O+] = -log (1.0 × 10-5) = 5.00
Since the pH of the lake is 5.00, the lake is NOT dead.
13-77
pH and pOH
 pH meters and pH
indicators are often
used to determine the
pH of a solution.
13-78
pH and pOH
 Indicators are brightly
colored organic dyes
that are weak acids or
bases.
 Phenolphthalein has an
acid form that is
colorless and a
conjugate base form
that is bright pink.
13-79
pH and pOH
13-80
Figure 13.18
Worksheet # 6 – 9
Complete the following table.
Solution
[H3O+]
[OH-]
pH
pOH
0.025 M
HCl
2.05 x10 -12 M
8.12 x10 -5 M
9.44 pH
13.52 pOH
13-81
Acidic/
basic/
neutral
Worksheet # 6 – 9 :Answers
Complete the following table.
Solution
[H3O+]
[OH-]
pH
pOH
Acidic/
basic/
neutral
0.025 M
HCl
0.025
4.0 x10 -13
1.60
12.40
Acidic
2.05 x10 -12 M 4.88 x10 -3
11.69
2.31
Basic
9.91
4.09
Basic
4.56
Basic
1.23 x10 -10
8.12 x10 -5 M
3.63 x10 -10 2.75 x10 -5 9.44 pH
13-82
0.33
3.02 x10 -14
0.48
13.52 pOH Acidic
Buffer Solutions
A buffer (also known as a buffer system) is a
combination of a weak acid and its conjugate base (or a
weak base and its conjugate acid) in about equal
concentrations.
The main buffer system in blood is made of
H2CO3/HCO3- :


Figure 13.18
13-83
Buffer Solutions
 Which of the following systems, when
added to water, can act as a buffer system?
For each buffer system, write a balanced
equation.
1. HCl and NaOH
not a buffer system
2. CH3CO2H and NaCH3CO2 a buffer system
3. HBr and KBr
13-84
not a buffer system