Acids and bases2

Section 14.1
The Nature of Acids and Bases
Acid-Base
Models
Arrhenius
Model
Lewis Model
Brønsted-Lowry
Model

Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
4.3Arrhenius Theory: Only in aqueous media

A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
acidbase acid base
15.1
acid
conjugate
base
base
conjugate
acid
Brønsted Theory: In all media (aqueous or non-aqueous)

Section 14.11
The Lewis Acid-Base Model
Lewis Acids and Bases
 Lewis acid: electron pair acceptor (Lewis acid has an
empty atomic orbital)
 Lewis base: electron pair donor (has a lone pair of
electrons)
Lewis acid Lewis base
Copyright © Cengage Learning. All rights reserved 7
O
H
H
O
H
H
AlAl3+
+ 6
6
3+

Lewis Acids and Bases
N H
••
H
H
acid base
F B
F
F
+ F B
F
F
N H
H
H
No protons donated or accepted!
(Arrhenius and Brønsted-Lowry models can not explain this reaction)

Section 14.11
The Lewis Acid-Base Model
Three Models for Acids and Bases

Section 14.1
Acid in Water
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
 Conjugate base is everything that remains of the acid
molecule after a proton is lost.
 Conjugate acid is formed when the proton is
transferred to the base.

Section 14.2
Acid Strength
 Strong acid:
 Ionization equilibrium lies far to the right.
 Yields a weak conjugate base.
HCl (aq) + H2O(l) H+(aq) + Cl-(aq)
base weak conjugate base
 Weak acid:
 Ionization equilibrium lies far to the left.
 Weaker the acid, stronger its conjugate base.
HF (aq) + H2O(l) H+(aq) + F-(aq)
11
Strong conjugate base

Section 14.2
Acid Strength
Various Ways to Describe Acid Strength

Section 14.2
Acid Strength
Water as an Acid and a Base
 Water is amphoteric:
 Behaves either as an acid or as a base.
 At 25°C:
Kw = [H+][OH–] = 1.0 × 10–14

Section 14.2
Acid Strength
Self-Ionization of Water
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Section 14.2
Acid Strength
acid base conjugate conjugate
acid base
What is the equilibrium constant expression for
an acid acting in water?
 
3H O A
=
HA
 
      K
CONCEPT CHECK!

Section 14.2
Acid Strength
If the equilibrium lies to the right, the value for Ka
is __________.
large (or >1)
If the equilibrium lies to the left, the value for Ka is
___________.
small (or <1)
CONCEPT CHECK!

Section 14.2
Acid Strength
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
–
18
CONCEPT CHECK!
1 23 4

Section 14.3
The pH Scale
pH = –log[H+]
Copyright © Cengage Learning. All rights reserved
19
What is pH?
pH is the measure of acidity.
[H+] = [OH-]
[H+] > [OH-]
[H+] < [OH-]
Solution Is
neutral
acidic
basic
[H+] = 1 x 10-7
[H+] > 1 x 10-7
[H+] < 1 x 10-7
pH = 7
pH < 7
pH > 7
At 250C
pH [H+]

Section 14.3
The pH Scale
The pH Scale and pH
Values of Some Common
Substances

Section 14.3
The pH Scale
Calculate the pH for each of the following solutions.
a) 1.0 × 10–4 M H+ b) 0.040 M OH–
EXERCISE!
a) pH = –log[H+] = –log(1.0 × 10–4 M) = 4.00
b) Kw = [H+][OH–] = 1.00 × 10–14 = [H+](0.040 M) = 2.5 × 10–13 M H+
pH = –log[H+] = –log(2.5 × 10–13 M) = 12.60

Section 14.3
The pH Scale
pH and pOH
 Recall:
Kw = [H+][OH–]
–log Kw = –log[H+] – log[OH–]
pKw = pH + pOH
14.00 = pH + pOH

Section 14.3
The pH Scale
The pH of a solution is 5.85. What are the [H+] and [OH-]
for this solution?
EXERCISE!
[H+] = 10–5.85 = 1.4 × 10–6 M
pH + pOH=14
pOH=8.15
[OH-]= 10-8.15= 7.08×10-9

Section 14.3
The pH Scale
Strong Electrolyte
100% Dissociation

Strong Acids are strong electrolytes
HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)
HNO3 (aq) + H2O (l) H3O+ (aq) + NO3
- (aq)
HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4
- (aq)
H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4
- (aq)
Strong Bases are strong electrolytes
NaOH (s) Na+ (aq) + OH- (aq)
H2O
KOH (s) K+ (aq) + OH- (aq)
H2O
Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq)
H2O

Weak Electrolyte
Not completely dissociated

CH3COOH CH3COO- (aq) + H+ (aq)
HF (aq) + H2O (l) H3O+ (aq) + F- (aq)
Weak Acids are weak electrolytes
- (aq)
HSO4
- (aq) + H2O (l) H3O+ (aq) + SO4
2- (aq)
H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)
F- (aq) + H2O (l) OH- (aq) + HF (aq)
Weak Bases are weak electrolytes
NO2
- (aq) + H2O (l) OH- (aq) + HNO2 (aq)

Conjugate acid-base pairs
• The conjugate base of a strong acid has no measurable
strength.
HZ + H2O H3O+ + Z-
Acid Conjugate Base
Base Conjugate Acid

What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation.
- (aq)
pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7
Start
End
0.002 M
0.002 M 0.002 M0.0 M
0.0 M 0.0 M
What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution?
Ba(OH)2 is a strong base – 100% dissociation.
Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq)
Start
End
0.018 M
0.018 M 0.036 M0.0 M
0.0 M 0.0 M
pH = 14.00 – pOH = 14.00 + log(0.036) = 12.6

Weak Acids (HA) and Acid Ionization Constants
Not completely dissociated

HA (aq) + H2O (l) H3O+ (aq) + A- (aq)
HA (aq) H+ (aq) + A- (aq)
Ka =
[H+][A-]
[HA]
Ka is the acid ionization constant
Ka
weak acid
strength
15.5

What is the pH of a 0.5 M HF solution (at 250C)?
HF (aq) H+ (aq) + F- (aq) Ka =
[H+][F-]
[HF]
= 7.1 x 10-4
HF (aq) H+ (aq) + F- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.50 0.00
-x +x
0.50 - x
0.00
+x
x x
Ka =
x2
0.50 - x
= 7.1 x 10-4
Ka 
x2
0.50
= 7.1 x 10-4
0.50 – x  0.50Ka << 1
x2 = 3.55 x 10-4 x = 0.019 M
[H+] = [F-] = 0.019 M pH = -log [H+] = 1.72
[HF] = 0.50 – x = 0.48 M

What is the pH of a 0.122 M monoprotic acid whose
Ka is 5.7 x 10-4?
HA (aq) H+ (aq) + A- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.122 0.00
-x +x
0.122 - x
0.00
+x
x x
Ka =
x2
0.122 - x
= 5.7 x 10-4
Ka 
x2
0.122
= 5.7 x 10-4
0.122 – x  0.122Ka << 1
x2 = 6.95 x 10-5 x = 0.0083 M
pH= 2.08

Percent Dissociation (Ionization) of weak acids

Section 14.5
Calculating the pH of Weak Acid Solutions
 For a given weak acid, the percent dissociation increases as the
acid becomes more dilute.
 For example:
Percent ionization: HF(0.01 M)>HF(0.1 M)
37
amount dissociated (mol/L)
Percent dissociation = 100%
initial concentration (mol/L)

For a monoprotic acid HA
Percent ionization =
[H+]
[HA]0
x 100% [HA]0 = initial concentration
Weak Acids (HA) not completely dissociate

Section 14.5
Calculating the pH of Weak Acid Solutions
A solution of 8.00 M formic acid (HCHO2) is 0.47%
ionized in water. Calculate the Ka value for formic acid.
38
EXERCISE!
Ka = = 1.8 × 10–4
[H+][HCO2-]
[HCHO2]
If 8.00 M of the acid is 0.47% ionized, then 0.038 M dissociates.
Ionized concentration: [HCHO2] = 0.47×0.08= 0.038
HCHO2(aq) + H2O  H3O+(aq) + CHO2
-(aq)
I 8.00 0 0
C -0.038 +0.038 +0.038
E 7.96 0.038 0.038

Section 14.6
Bases
Bases
- Arrhenius theory: bases produce OH– ions.
-Brønsted–Lowry theory: bases are proton acceptors.
In a basic solution at 25°C, pH > 7.

Section 14.6
Bases
 Ionic compounds containing OH- are generally
considered strong bases.
 LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
 pOH = –log[OH–]
 pH = 14.00 – pOH
Strong bases-100% dissociation

Section 14.6
Bases
Weak bases
Not completely dissociate

Section 14.6
Bases
 Equilibrium expression for weak bases uses Kb.
CN–(aq) + H2O(l) HCN(aq) + OH–(aq)
 
b
HCN OH
=
CN


  
  
K
Base

NH3 (aq) + H2O (l) NH4
+ (aq) + OH- (aq)
Kb =
[NH4
+][OH-]
[NH3]
Kb is the base ionization constant
Kb
weak base
strength
Solve weak base problems like weak acids
except solve for [OH-] instead of [H+].

Section 14.6
Bases
 pH calculations for solutions of weak bases are very
similar to those for weak acids.
 Kw = [H+][OH–] = 1.0 × 10–14
 pOH = –log[OH–]
 pH = 14.00 – pOH

Section 14.6
Bases
Polyprotic acid

Section 14.7
Polyprotic Acids
 Acids that can furnish more than one proton. H2SO4, H3PO4
 Always dissociates in a stepwise manner, one proton at a time.
H2SO4(aq) HSO4
- (aq) + H+
 SO4
2- (aq) + H+
 The conjugate base of the first dissociation equilibrium becomes
the acid in the second step.
 For a typical weak polyprotic acid:
Ka1 > Ka2 > Ka3
 For a typical polyprotic acid in water, only the first dissociation
step is important to pH.

Section 14.7
Polyprotic Acids
Calculate the pH of a 5.00 M solution of H3PO4.
Ka1 = 7.5 × 10-3
Ka2 = 6.2 × 10-8
Ka3 = 4.8 × 10-13
48
EXERCISE!

Section 14.7
Polyprotic Acids
Salts

Section 14.8
Acid-Base Properties of Salts
 Salts are ionic compounds.
 When dissolved in water, break up into its ions (cation
and anion)
 In some salts cation and anion can behave as acids or
bases.

Section 14.8

Section 14.1
Salts
Neutral
salts
Basic saltsAcidic salts

Section 14.8
 The salt of a strong acid and a strong base gives a
neutral solution.
 KCl (From HCl and KOH), NaNO3 (From NaCl and HNO3)
 CaSO4
 BaBr2
1) Neutral Salts:

Section 14.8
Salts derived from a strong base and a weak acid.
 Other examples:
 NaF (F- is conjugate base of HF), KCH3COO(CH3COO- is
conjugate base of CH3COOH)
2) Basic Salts:
NaCH3COOH (s) Na+ (aq) + CH3COO- (aq)
H2O
CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH- (aq)

Section 14.8
- Salts derived from a strong acid and a weak base.
 NH4Cl (NH4
+ is conjugate base of NH3)
- Salts with small, highly charged metal cations (e.g. Al3+,
Cr3+, and Be2+)
3) Acidic Salts:
NH4Cl (s) NH4
+ (aq) + Cl- (aq)
H2O
NH4
+ (aq) NH3 (aq) + H+ (aq)
Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq)
3+ 2+

Section 14.8
What about the salts derived from a
Weak base and a weak acid?

Solutions in which both the cation and the anion hydrolyze:
• Kb for the anion > Ka for the cation, solution will be basic
• Kb for the anion < Ka for the cation, solution will be acidic
• Kb for the anion  Ka for the cation, solution will be neutral

Section 14.8
Arrange the following 1.0 M solutions from lowest to highest pH.
HBr NaOH NH4Cl
NaCN NH3 HCN
NaCl HF
Justify your answer.
HBr, HF, HCN, NH4Cl, NaCl, NaCN, NH3, NaOH
CONCEPT CHECK!
The order is: HBr (strong acid), HF (Ka = 7.2 x 10-4), HCN (Ka = 6.2 x 10-10), NH4Cl (Ka =
5.6 x 10-10), NaCl (neutral), NaCN (Kb = 1.6 x 10-5), NH3 (Kb = 1.8 x 10-5), NaOH (strong
base).

Section 14.8
Molecular Structure and Acid Strength
-HX
-HOX
-HOXO
-HOXO2
-HOXO3
-(HO)nXOm

H X H+ + X-
The
stronger
the bond
The
weaker
the acid
HF << HCl < HBr < HI
15.9
across a group
- Group 7A
HX

What about 6A group?
H2O < H2S < H2Se < H2Te

What about 5A group?
NH3 < PH3 < AsH3 < SbH3

What about a period?
NH3 < H2O < HF
PH3 < H2S << HCl
AsH3 < H2Se << HBr

Section 14.9
The Effect of Structure on Acid-Base Properties
Oxyacids: Acids containing O
1) H–O–X (HOCl, HOBr, HOI)
These acids are very weak.
HOCl> HOBr> HOI

Section 14.9
Comparison of Electronegativity of X and Ka Value

Section 14.9
2) H–O–XO (HOClO, HOBrO, HOIO)
These acids are weak.
HOClO> HOBrO> HOIO
Other example: HNO2 (HONO)

Section 14.9
3) H–O–XO2 (HOClO2, HOBrO2, HOIO2)
These acids are Strong.
HOClO2> HOBrO2> HOIO2
Other example: HNO3 (HONO2)

Section 14.9
3) H–O–XO3 (HOClO3, HOBrO3, HOIO3)
These acids are Very Strong.
HOClO3> HOBrO3> HOIO3
HClO4 > HClO3 > HClO2 > HClO

Section 14.9
4) (HO)nXO (H3PO4, H3PO3, H3PO2)
These acids are weak.
H3PO4≈ H3PO3≈ H3PO2
Other example: H2SO3 (HO)2SO

Section 14.9
Several Series of
Oxyacids and Their
Ka Values
3 single O
2 single O
1 single O
No single O
2 single O
1 single O
1 single O
2 single O

Section 14.9
Acid-Base Properties of Oxides

Section 14.10
1) Acidic Oxides (Acid Anhydrides):
 O—X bond is strong and covalent.
SO2, NO2, CO2
 When H—O—X grouping is dissolved in water, the O—X
bond will remain intact. It will be the polar and
relatively weak H—O bond that will tend to break,
releasing a proton.

Section 14.10
2) Basic Oxides (Acid Anhydrides):

Oxides of the Representative Elements
In Their Highest Oxidation States
CO2 (g) + H2O (l) H2CO3 (aq)
N2O5 (g) + H2O (l) 2HNO3 (aq)
Acidic Oxide:
Acidic Oxide:
Basic Oxide: Na2O (s) + H2O (l) 2NaOH (aq)

What is amphoteric oxides
an amphoteric compound is a molecule or ion that can react
both as an acid and as a base. Al2O3 is an example of an
amphoteric oxide.
•In acid: Al2O3 + 6 HCl→ 2 AlCl3 + 3 H2O
•In base: Al2O3 + 2 NaOH + 3 H2O → 2 Na[Al(OH)4]

Acids and bases2

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