chapter 16 acids and bases. chapter 16 table of contents 2 16.1acids and bases 16.2acid strength...
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Chapter 16
Acids and Bases
Chapter 16
Table of Contents
2
16.1Acids and Bases
16.2Acid Strength
16.3Water as an Acid and a Base
16.4The pH Scale
16.5 Calculating the pH of Strong Acid Solutions
16.6Buffered Solutions
Section 16.1
Acids and Bases
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3
Models of Acids and Bases
• Arrhenius: Acids produce H+ ions in solution, bases produce OH– ions.
• Brønsted–Lowry: Acids are proton (H+) donors, bases are proton acceptors.
HCl + H2O → Cl– + H3O+
acid base
Section 16.1
Acids and Bases
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Acid in Water
HA(aq) + H2O(l) H3O+(aq) + A–(aq)
acid base conjugate conjugate acid base
• Conjugate base is everything that remains of the acid molecule after a proton is lost and is found by removing an H+.
• Conjugate acid is formed when the proton is transferred to the base and is found by adding an H+.
Section 16.1
Acids and Bases
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5
Practice – Conjugate Acids and Bases
• Identify the conjugate acid for each base.
a) F-
b) HCO3-
c) H2O
13-
HF
H2CO3
H3O+
• Identify the conjugate base for each acid.
a) HBr
b) HSO4-
c) H2O
Br-
SO4-2
OH-
Section 16.1
Acids and Bases
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• Water acts as a base accepting a proton from the acid.
• Forms hydronium ion (H3O+).
Section 16.1
Acids and Bases
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Concept Check
Which of the following represent conjugate acid–base pairs?
a) HCl, HNO3
b) H3O+, OH–
c) H2SO4, SO42–
d) HCN, CN–
Acid Strength
Section 16.2
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8
Strong Acid
• Completely ionized or completely dissociated
Acid Strength
Section 16.2
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9
Weak Acid
• Most of the acid molecules remain intact.
Acid Strength
Section 16.2
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Behavior of Acids of Different Strengths in Aqueous Solution
Acid Strength
Section 16.2
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• A strong acid contains a relatively weak conjugate base.
Acid Strength
Section 16.2
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12
Ways to Describe Acid Strength
Acid Strength
Section 16.2
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Common Strong Acids
• Sulfuric acid, H2SO4
• Hydrochloric acid, HCl
• Nitric acid, HNO3
• Perchloric acid, HClO4
Common Weak Acids
• Hydroflouric Acid HF
• Phosphoric Acid H3PO4
• Organic Acids
Acid Strength
Section 16.2
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14
Common Strong Bases
• Ammonia NH3
• Organic Ammines R-NH2
• Insoluble Hydroxides
Common Weak Bases
• Group I Hydroxides NaOH
Acid Strength
Section 16.2
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• Oxyacid – acidic proton is attached to an oxygen atom• Organic acid – have a carbon atom backbone and
commonly contain the carboxyl group:
Typically a weak acid
Acid Strength
Section 16.2
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Concept Check
Consider a 1.0 M solution of HCl.
Order the following from strongest to weakest base and explain:
H2O(l)
A–(aq) (from weak acid HA)
Cl–(aq)
Acid Strength
Section 16.2
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Let’s Think About It…
• How good is Cl–(aq) as a base?• Is A–(aq) a good base?
The bases from strongest to weakest are:
A–, H2O, Cl–
Acid Strength
Section 16.2
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18
Concept Check
Acetic acid (HC2H3O2) and HCN are both weak acids. Acetic acid is a stronger acid than HCN.
Arrange these bases from weakest to strongest and explain your answer:
H2O Cl– CN– C2H3O2
–
The bases from weakest to strongest are: Cl–, H2O, C2H3O2
–, CN–
Water as an Acid and a Base
Section 16.3
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19
Water as an Acid and a Base
• Water is amphoteric: H2O(l) + H2O(l) H3O+(aq) + OH–(aq)
Behaves either as an acid or as a base.• At 25°C:
Kw = [H+][OH–] = 1.0 × 10–14
[X] = molarity of X or moles X per Liter
• No matter what the solution contains, the product of [H+] and [OH–] must always equal 1.0 × 10–14 at 25°C.
Water as an Acid and a Base
Section 16.3
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Three Possible Situations
• [H+] = [OH–]; neutral solution• [H+] > [OH–]; acidic solution• [H+] < [OH–]; basic solution
In each case, however,
Kw = [H+][OH–] = 1.0 × 10–14.
Water as an Acid and a Base
Section 16.3
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Concept Check
In an acidic aqueous solution, which statement below is correct?
a) [H+] < 1.0 × 10–7 Mb) [H+] > 1.0 × 10–7 Mc) [OH–] > 1.0 × 10–7 Md) [H+] < [OH–]
Water as an Acid and a Base
Section 16.3
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22
Exercise
In an aqueous solution in which [OH–] = 2.0 x 10–10 M, the [H+] = ________ M, and the solution is ________.
a) 2.0 × 10–10 M ; basicb) 1.0 × 10–14 M ; acidicc) 5.0 × 10–5 M ; acidicd) 5.0 × 10–5 M ; basic
[H+] = Kw/[OH–] = 1.0 × 10–14/2.0 × 10–10 = 5.0 × 10–5 M.Since [H+] is greater than [OH–], the solution is acidic.
Section 16.4
The pH Scale
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• pH = –log[H+]• A compact way to
represent solution acidity.• pH decreases as [H+]
increases.• Significant figures:
The number of decimal places in the log is equal to the number of significant figures in the original number.
Section 16.4
The pH Scale
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pH Range
• pH = 7; neutral• pH > 7; basic
Higher the pH, more basic.• pH < 7; acidic
Lower the pH, more acidic.
Section 16.4
The pH Scale
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25
The pH Scale and pH Values of Some Common Substances
Blood cells can exist only over a narrow pH range.
Section 16.4
The pH Scale
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Exercise
Calculate the pH for each of the following solutions.
a) 1.0 × 10–4 M H+
pH = –log[H+] = –log(1.0 × 10–4 M) = 4.00
b) 0.040 M OH–
Kw = [H+][OH–] = 1.00 × 10–14 = [H+](0.040 M)
[H+] = 2.5 × 10–13 M H+
pH = –log[H+] = –log(2.5 × 10–13 M) = 12.60
Section 16.4
The pH Scale
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Exercise
The pH of a solution is 5.85. What is the [H+] for this solution?
[H+] = 1.4 × 10–6 M
[H+] = 10–5.85 = 1.4 × 10–6 M
Section 16.4
The pH Scale
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pH and pOH
• Recall:
Kw = [H+][OH–]
–log Kw = –log[H+] – log[OH–]
pKw = pH + pOH
14.00 = pH + pOH
Section 16.4
The pH Scale
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Exercise
Calculate the pOH for each of the following solutions.
a) 1.0 × 10–4 M H+
pH = –log[H+] = –log(1.0 × 10–4 M) = 4.00;
So, 14.00 = 4.00 + pOH; pOH = 10.00
b) 0.040 M OH–
pOH = –log[OH–] = –log(0.040 M) = 1.40
Section 16.4
The pH Scale
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30
Exercise
The pH of a solution is 5.85. What is the [OH–] for this solution?
[OH–] = 7.1 × 10–9 M
14.00 = 5.85 + pOH; pOH = 8.15
[OH–] = 10–8.15 = 7.1 × 10–9 M
Section 16.5
Calculating the pH of Strong Acid Solutions
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Concept Check
Consider an aqueous solution of 2.0 × 10–3 M HCl. What is the pH?
pH = 2.70
2.0 × 10–3 M HCl → 2.0 × 10–3 M H+ and 2.0 × 10–3 M Cl–
pH = –log[H+] pH = –log(2.0 × 10–3 M)pH = 2.70
Section 16.5
Calculating the pH of Strong Acid Solutions
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Concept Check
Calculate the pH of a 1.5 × 10–11 M solution of HCl.
pH = 10.82
Section 16.5
Calculating the pH of Strong Acid Solutions
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Concept Check
Calculate the pH of a 1.5 × 10–2 M solution of HNO3.
pH = 1.82
pH = –log[H+] pH = –log(1.5 × 10–2 M)pH = 1.82
Section 16.4
The pH Scale
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Practice Calculating ph, pOH, [H+] and [OH-]
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pH [H+] [OH-] pOH
_____ ______ 6.8x10-7 ______
9.2 ______ ______ ______
6.177.83
4.8
1.5x10-8
6 x10-10 2x10-5
Section 16.6
Buffered Solutions
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• Buffered solution – resists a change in its pH when either an acid or a base has been added. Presence of a weak acid and its conjugate
base buffers the solution.
Section 16.6
Buffered Solutions
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The Characteristics of a Buffer
1. The solution contains a weak acid HA and its conjugate base A–.
2. The buffer resists changes in pH by reacting with any added H+ or OH– so that these ions do not accumulate.
3. Any added H+ reacts with the base A–.
H+(aq) + A–(aq) → HA(aq)
4. Any added OH– reacts with the weak acid HA.
OH–(aq) + HA(aq) → H2O(l) + A–(aq)
Section 16.6
Buffered Solutions
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How Buffers Really Work
HA H+ + A-
A- + H2O HA + OH-
Consider the Equilibrium and La Chatelier’s Principle.
As you add H+ to a buffered system containing equal amounts of weak acid and its conjugate base, it will go through this equilibrium treadmill and keep the H+ concentration in check.
Section 16.6
Buffered Solutions
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Concept Check
If a solution is buffered with NH3 to which has been added NH4Cl, what reaction will occur if a strong base such as NaOH is added?
a) NaOH + NH3 → NaNH4Ob) NaOH + NH4
+ → Na+ + NH3 + H2Oc) NaOH + Cl– → NaCl + OH–
d) NaOH + H2O → NaH3O2