Acids and Bases Chapter 15 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display

Acids and Bases

Chapter 15

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

1) Compare and contrast Arrhenius, Brønsted, and Lewis acids

and bases.

2) Describe what is meant by conjugate acid-base pairs and

give several examples.

3) Use Kw to determine [H+] and [OH-] of solutions.

4) Discuss the pH scale and calculate pH and pOH given either

[H+] or [OH-].

5) Define strong and weak acids and bases and give several

examples of each.

6) Relate properties of conjugate acid-base pairs.

7) Determine Ka from experimental data.

8) Calculate pH, [H+] , weak acid concentration, and conjugate base

concentration given Ka and the initial concentration of the weak

acid using the quadratic equation or the method of successive

approximation as needed.

9) Calculate percent ionization for a weak acid needed.

10) Calculate pH, [OH-], weak base concentration, and conjugate acid

concentration given Kb and the initial concentration of the weak

base using the quadratic equation or the method of successive

approximation

11) Show the relationship between Ka,, Kb, and Kw.

12) Calculate concentrations of all species present at equilibrium for

diprotic and polyprotic acids.

13) Relate molecular structure and the strength of acids.

14) Predict the relative strengths of oxoacids.

15) Describe salt hydrolysis and explain how some salts produce

neutral solutions, some acidic solutions and others basic solutions.

16) Calculate the pH of salt solutions and determine the percent

hydrolysis.

17) Describe acid-base properties of oxides and hydroxides.

18) Give several examples of Lewis acid-base reactions.

Acids

Have a sour taste. Vinegar owes its taste to acetic acid. Citrusfruits contain citric acid.

React with certain metals to produce hydrogen gas.

React with carbonates and bicarbonates to produce carbon dioxide gas

Have a bitter taste.

Feel slippery. Many soaps contain bases.

Bases

4.3

Arrhenius acid is a substance that produces H+ (H3O+) in water

Arrhenius base is a substance that produces OH- in water

4.3

Copyright ©2001 by Harcourt, Inc. All rights reserved.

Brønsted-Lowry Model of Acids and Bases

Acid is proton donor, base is proton acceptor

HB(aq)+A-(aq) HA(aq)+B-(aq)

HB, HA are acids; A-, B-are bases. HB –B-and HA –A-are conjugate acid-base pairs.

Copyright © 2001 by Harcourt, Inc. All rights reserved.

Acidic and Basic Water Solutions

Aqueous solution

pH


Aqueous solution

In any aqueous solution, there is an equilibrium between H30+

(H+) ions and OH–ions:

H2O H+(aq) + OH–(aq)

KW = [H+] ́ [OH–] = 1.0 ́ 10-14 at 25°C

1. Pure water: [H+] = [OH–] = 1.0 ́ 10–7 M; neutral solution

2. Acidic solution: [H+] > 1.0 ́ 10–7 M> [OH–]

3. Basic solution: [OH–] > 1.0 ́ 10–7 M> [H+]In seawater, [H+] = 5 ́ 10-9 M; [OH–] = ?[OH–] = (1.0 ́ 10-14) / (5 ́ 10-9) = 2 ́ 10-6 M


pH

pH = –1og10[H+]

Neutral solution: pH = 7.0

Acidic solution: pH < 7.0

Basic solution: pH > 7.0Suppose [H+] = 2.4 ´10-6 M; calculate pH

pH = –1og10(2.4 ´10–6) = 5.62

Suppose pH = 8.68; calculate [H+]

[H+] = 10–8.68 = 2.1 ´10–9 M

pOH = –1og10[OH–]; pH + pOH = 14.00 at 25°C

A Brønsted acid is a proton donorA Brønsted base is a proton acceptor

acidbase acid base

15.1

acidconjugate

basebase conjugate

acid

O

H

H + O

H

H O

H

H H OH-+[ ] +

Acid-Base Properties of Water

H2O (l) H+ (aq) + OH- (aq)

H2O + H2O H3O+ + OH-

acid conjugate

base

base conjugate

acid

15.2

autoionization of water

H2O (l) H+ (aq) + OH- (aq)

The Ion Product of Water

Kc =[H+][OH-]

[H2O][H2O] = constant

Kc[H2O] = Kw = [H+][OH-]

The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature.

At 250CKw = [H+][OH-] = 1.0 x 10-14

[H+] = [OH-]

[H+] > [OH-]

[H+] < [OH-]

Solution Is

neutral

acidic

basic

15.2

What is the concentration of OH- ions in a HCl solution whose hydrogen ion concentration is 1.3 M?

Kw = [H+][OH-] = 1.0 x 10-14

[H+] = 1.3 M

[OH-] =Kw

[H+]

1 x 10-14

1.3= = 7.7 x 10-15 M

15.2

pH – A Measure of Acidity

pH = -log [H+]

[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+]

15.3

15.3

pOH = -log [OH-]

[H+][OH-] = Kw = 1.0 x 10-14

-log [H+] – log [OH-] = 14.00

pH + pOH = 14.00

The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater?

pH = -log [H+]

[H+] = 10-pH = 10-4.82 = 1.51 x 10-5 M

The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?

pH + pOH = 14.00

pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60

pH = 14.00 – pOH = 14.00 – 6.60 = 7.40

15.3

Strong Electrolyte – 100% dissociation

NaCl (s) Na+ (aq) + Cl- (aq)H2O

Weak Electrolyte – not completely dissociated

CH3COOH CH3COO- (aq) + H+ (aq)

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)

15.4

HF (aq) + H2O (l) H3O+ (aq) + F- (aq)

Weak Acids are weak electrolytes


HSO4- (aq) + H2O (l) H3O+ (aq) + SO4

2- (aq)

H2O (l) + H2O (l) H3O+ (aq) + OH- (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

15.4

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.

• H3O+ is the strongest acid that can exist in aqueous solution.

• The OH- ion is the strongest base that can exist in aqeous solution.

15.4

15.4

Strong Acid Weak Acid

15.4

What is the pH of a 2 x 10-3 M HNO3 solution?

HNO3 is a strong acid – 100% dissociation.


pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.70

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) = 12.5615.4

HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

Weak Acids (HA) and Acid Ionization Constants

HA (aq) H+ (aq) + A- (aq)

Ka =[H+][A-][HA]

Ka is the acid ionization constant

Ka

weak acidstrength

15.5


Calculation of [H+] in solution example

Find [H+] in 0.200 M HC2H3O2, Ka=1.8 ´10–5

Let [H+] = x, then [C2H3O2–] = x,

[HC2H3O2] = 0.200 – x

= 1.8 ´10–5

Assume 0.200 - x 0.200 ; x = 1.9 ´10–3 M

In this case, % ionization = ´100 = 1.0%

x2

0.200 – x

1.9 ´10–3

0.200

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 M15.5

When can I use the approximation?

0.50 – x 0.50Ka << 1

When x is less than 5% of the value from which it is subtracted.

x = 0.0190.019 M0.50 M

x 100% = 3.8%Less than 5%

Approximation ok.

What is the pH of a 0.05 M HF solution (at 250C)?

Ka x2

0.05= 7.1 x 10-4 x = 0.006 M

0.006 M0.05 M

x 100% = 12%More than 5%

Approximation not ok.

Must solve for x exactly using quadratic equation or method of successive approximation. 15.5

Solving weak acid ionization problems:

1. Identify the major species that can affect the pH.

• In most cases, you can ignore the autoionization of water.

• Ignore [OH-] because it is determined by [H+].

2. Use ICE to express the equilibrium concentrations in terms of single unknown x.

3. Write Ka in terms of equilibrium concentrations. Solve for x by the approximate method. If approximation is not valid, solve for x exactly.

4. Calculate concentration of all species and/or pH of the solution.

15.5

What is the pH of a 0.122 M monoprotic acid whose Ka is 5.7 x 10-4?


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

0.0083 M0.122 M

x 100% = 6.8%More than 5%

Approximation not ok.

15.5

Ka =x2

0.122 - x= 5.7 x 10-4 x2 + 0.00057x – 6.95 x 10-5 = 0

ax2 + bx + c =0-b ± b2 – 4ac

2ax =

x = 0.0081 x = - 0.0081


Initial (M)

Change (M)

Equilibrium (M)

0.122 0.00

-x +x

0.122 - x

0.00

+x

x x

[H+] = x = 0.0081 M pH = -log[H+] = 2.09

15.5

percent ionization = Ionized acid concentration at equilibrium

Initial concentration of acidx 100%

For a monoprotic acid HA

Percent ionization = [H+]

[HA]0

x 100% [HA]0 = initial concentration

15.5


Weak Bases

Molecules

Anions derived from weak acids


Molecules

NH3(aq) + H2O OH–(aq) + NH4+(aq)


Anions derived from weak acids

F–(aq) + H2O HF(aq) + OH–(aq)


Weak Bases

Equilibrium constant

•Expression for Kb

•Relation between Ka and Kb

•Calculation of [OH–] in solution of weak base

NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)

Weak Bases and Base Ionization Constants

Kb =[NH4

+][OH-][NH3]

Kb is the base ionization constant

Kb

weak basestrength

15.6

Solve weak base problems like weak acids except solve for [OH-] instead of [H+].

15.6

15.7

Ionization Constants of Conjugate Acid-Base Pairs


A- (aq) + H2O (l) OH- (aq) + HA (aq)

Ka

Kb

H2O (l) H+ (aq) + OH- (aq) Kw

KaKb = Kw

Weak Acid and Its Conjugate Base

Ka = Kw

Kb

Kb = Kw

Ka


Relation between Ka and Kb

Ka ´Kb = KW = 1.0 ´10–14

Ka Kb

HF 6.9 ´10–4 F– 1.4 ´10–11

HAc 1.8 ´10–5 Ac– 5.6 ´10–10

NH4+ 5.6 ´10–10 NH3 1.8 ´10–5

Strength of base is inversely related to that of conjugate weak acid.


Calculation of [OH–] in solution of weak base

= 1.4 ´10–11 ; = 1.4 ´10–11

[OH–] = 1.2 ´10–6 M ; pH = 8.08

[HF][OH–][F–]

[OH–]2

0.10

15.8

Molecular Structure and Acid Strength

H X H+ + X-

The stronger the bond

The weaker the acid

HF << HCl < HBr < HI

15.9


Z O H Z O- + H+- +

The O-H bond will be more polar and easier to break if:

• Z is very electronegative or

• Z is in a high oxidation state

15.9


1. Oxoacids having different central atoms (Z) that are from the same group and that have the same oxidation number.

Acid strength increases with increasing electronegativity of Z

H O Cl O

O••

••••••

••

••••

••••

H O Br O

O••

••••••

••

••••

••••Cl is more electronegative than Br

HClO3 > HBrO3

15.9


2. Oxoacids having the same central atom (Z) but different numbers of attached groups.

Acid strength increases as the oxidation number of Z increases.

HClO4 > HClO3 > HClO2 > HClO

15.9

Acid-Base Properties of SaltsNeutral Solutions:

Salts containing an alkali metal or alkaline earth metal ion (except Be2+) and the conjugate base of a strong acid (e.g. Cl-, Br-, and NO3

-).

NaCl (s) Na+ (aq) + Cl- (aq)H2O

Basic Solutions:

Salts derived from a strong base and a weak acid.

NaCH3COOH (s) Na+ (aq) + CH3COO- (aq)H2O

CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH- (aq)

15.10

Acid-Base Properties of Salts

Acid Solutions:

Salts derived from a strong acid and a weak base.

NH4Cl (s) NH4+ (aq) + Cl- (aq)

H2O

NH4+ (aq) NH3 (aq) + H+ (aq)

Salts with small, highly charged metal cations (e.g. Al3+, Cr3+, and Be2+) and the conjugate base of a strong acid.

Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq)3+ 2+

15.10

Acid Hydrolysis of Al3+

15.10

Acid-Base Properties of Salts

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

15.10

Oxides of the Representative ElementsIn Their Highest Oxidation States

15.11

Arrhenius acid is a substance that produces H+ (H3O+) in water

A Brønsted acid is a proton donor

A Lewis acid is a substance that can accept a pair of electrons

A Lewis base is a substance that can donate a pair of electrons

Definition of An Acid

H+ H O H••••

+ OH-••••••

acid base

N H••

H

H

H+ + N H

H

H

H

acid base15.12

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!

15.12


Overall Results

Salt Cation Anion

NaNO3 Na+(Sp.) NO3–(Sp.) neutral

KF K+(Sp.) F–(Ba.) basic

FeCl2 Fe2+(Ac.) Cl–(Sp.) acidic

If cation is acidic, anion basic, compare Ka, Kb values

NH4F: Ka = 5.6 ´10–10, Kb = 1.4 ´10–11; acidic

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Acids and Bases Chapter 15 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display