chapter 15 acid–base equilibria. chapter 15 table of contents copyright © cengage learning. all...

Chapter 15

Acid–Base Equilibria

Chapter 15

Table of Contents

Copyright © Cengage Learning. All rights reserved 2

15.1Solutions of Acids or Bases Containing a Common Ion

15.2 Buffered Solutions

15.3 Buffering Capacity

15.4 Titrations and pH Curves

15.5 Acid–Base Indicators

Section 15.1

Solutions of Acids or Bases Containing a Common Ion

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Common Ion Effect

• Shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction.

• An application of Le Châtelier’s principle.

Section 15.1

Solutions of Acids or Bases Containing a Common Ion

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Example

HCN(aq) + H2O(l) H3O+(aq) + CN-

(aq)

• Addition of NaCN will shift the equilibrium to the left because of the addition of CN-, which is already involved in the equilibrium reaction.

• A solution of HCN and NaCN is less acidic than a solution of HCN alone.

Section 15.2

Atomic MassesBuffered Solutions

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Key Points about Buffered Solutions

• Buffered Solution – resists a change in pH.• They are weak acids or bases containing a

common ion.• After addition of strong acid or base, deal with

stoichiometry first, then the equilibrium.

Section 15.2


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Henderson–Hasselbalch Equation

a

ApH = p + log

HA

K

• For a particular buffering system (conjugate acid–base pair), all solutions that have the same ratio [A–] / [HA] will have the same pH.

Section 15.2


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Exercise

What is the pH of a buffer solution that is 0.45 M acetic acid (HC2H3O2) and 0.85 M sodium acetate (NaC2H3O2)? The Ka for acetic acid is 1.8 × 10–5.

pH = 5.02

Section 15.2


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Section 15.3

The Mole Buffering Capacity

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• The amount of protons or hydroxide ions the buffer can absorb without a significant change in pH.

• Determined by the magnitudes of [HA] and [A–].• A buffer with large capacity contains large

concentrations of the buffering components.

Section 15.3


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The amount of H+ or OH- that buffered solution can absorb without a significant change in pH

Buffer capacity measures how well a solution resists changes in pH when acid or base is added. The greater the buffer capacity, the less the pH changes.

Section 15.3


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0.67

0.06

0.04

HA

A1

0.05

0.05

HA

A

:B

0.9965.01

4.99

HA

A1

5.00

5.00

HA

A

:A

H mol 0.01

H mol 0.01

2) Magnitudes of [HA] and [A-] the capacity of a buffered soln.

Ex : soln A : 5.00 M HOAc + 5.00 M NaOAc soln B : 0.05 M HOAc + 0.05 M NaOAc pH change when 0.01 mol of HCl(g) is added

Section 15.3


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3) [A-] / [HA] ratio the pH of a buffered soln.

soln new 50.5% soln original

49.5HA

A100

0.01

1.00

HA

A

0.98HA

A1.00

1.00

1.00

HA

A

2% :C

H mol 0.01

H mol 0.01

Section 15.4

Titrations and pH Curves

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The pH Curve for the Titration of 50.0 mL of 0.200 M HNO3 with 0.100 M NaOH

Section 15.4


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The pH Curve for the Titration of 100.0 mL of 0.50 M NaOH with 1.0 M HCI

Section 15.4


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Concept Check

Consider a solution made by mixing 0.10 mol of HCN (Ka = 6.2 x 10–10) with 0.040 mol NaOH in 1.0 L of aqueous solution.

What are the major species immediately upon mixing (that is, before a reaction)?

HCN, Na+, OH–, H2O

Section 15.4


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Let’s Think About It…

K = 1.8 x 109

HC2H3O2(aq) + OH– C2H3O2–(aq) + H2O

Before 0.20 mol 0.030 mol 0

Change –0.030 mol –0.030 mol +0.030 mol

After 0.17 mol 0 0.030 mol

Section 15.4


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Steps Toward Solving for pH

Ka = 1.8 x 10–5

pH = 3.99

HC2H3O2(aq) + H2O H3O+ + C2H3O2-(aq)

Initial 0.170 M ~0 0.030 M

Change –x +x +x

Equilibrium 0.170 – x x 0.030 + x

Section 15.4


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Exercise

Calculate the pH of a 100.0 mL solution of 0.100 M acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5.

pH = 2.87

Section 15.4


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Concept Check

Calculate the pH of a solution made by mixing 100.0 mL of a 0.100 M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, and 50.0 mL of a 0.10 M NaOH solution.

pH = 4.74

Section 15.4


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Concept Check

Calculate the pH of a solution at the equivalence point when 100.0 mL of a 0.100 M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, is titrated with a 0.10 M NaOH solution.

pH = 8.72

Section 15.4


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The pH Curve for the Titration of 50.0 mL of 0.100 M HC2H3O2 with 0.100 M NaOH

Section 15.4


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The pH Curves for the Titrations of 50.0-mL Samples of 0.10 M Acids with Various Ka Values with 0.10 M NaOH

Section 15.4


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The pH Curve for the Titration of 100.0mL of 0.050 M NH3 with 0.10 M HCl

Section 15.4


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P.228

Kjeldahl Nitrogen Analysis :

developed in 1883, the analysis remains one of the most widely used methods for determining nitrogen in organic substances such as protein, cereal, and flour.

Section 15.4


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Melamine contaminated milk

Section 15.5

Acid–Base Indicators

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• Marks the end point of a titration by changing color.

• The equivalence point is not necessarily the same as the end point (but they are ideally as close as possible).

Section 15.5


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The Acid and Base Forms of the Indicator Phenolphthalein

Section 15.5


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The Methyl Orange Indicator is Yellow in Basic Solution and Red in Acidic Solution

Section 15.5


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1) Usually a weak organic acid or base that has distinctly different colors in its nonionized & ionized forms.

HIn(aq) H+(aq) + In-

(aq) pKHIn

nonionized ionized

form form

Section 15.5


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1)

2)

3)

-

HIn

-

HIn

In & HIn of color the of ncombinatio

InHIn

1pKpH

In of color theshow 10

HIn

In

1-pKpH

HIn of color theshow 10

In

HIn

Section 15.5


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Useful pH Ranges for Several Common Indicators

chapter 15 acid–base equilibria. chapter 15 table of contents copyright © cengage learning. all...

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