Chapter 15 Acid-Base Equilibria - HCC Learning Web

Chapter 15

Acid-Base Equilibria

Section 15.1Solutions 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 15.1Solutions of Acids or Bases Containing a Common Ion

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 15.2Buffered Solutions

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.



Adding an Acid to a Buffer


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Buffers





buffers.html


Solving Problems with Buffered Solutions



Buffering: How Does It Work?



Buffering: How Does It Work?



Henderson–Hasselbalch Equation

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


a

ApH = p + log

HA

K


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


EXERCISE!




Buffered Solution Characteristics

Buffers contain relatively large concentrations of a weak acid and corresponding conjugate base.

Added H+ reacts to completion with the weak base.

Added OH- reacts to completion with the weak acid.



Buffered Solution Characteristics

The pH in the buffered solution is determined by the ratio of the concentrations of the weak acid and weak base. As long as this ratio remains virtually constant, the pH will remain virtually constant. This will be the case as long as the concentrations of the buffering materials (HA and A– or B and BH+) are large compared with amounts of H+ or OH– added.


Section 15.3Buffering Capacity

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.



Optimal buffering occurs when [HA] is equal to [A–].

It is for this condition that the ratio [A–] / [HA] is most resistant to change when H+ or OH– is added to the buffered solution.



Choosing a Buffer

pKa of the weak acid to be used in the buffer should be as close as possible to the desired pH.


Section 15.4Titrations and pH Curves

Titration Curve

Plotting the pH of the solution being analyzed as a function of the amount of titrant added.

Equivalence (Stoichiometric) Point – point in the titration when enough titrant has been added to react exactly with the substance in solution being titrated.



Neutralization of a Strong Acid with a Strong Base





neutralizationofacid.html


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

with 0.100 M NaOH



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



Weak Acid–Strong Base Titration

Step 1: A stoichiometry problem (reaction is assumed to run to completion) then

determine concentration of acid remaining and conjugate base formed.

Step 2: An equilibrium problem (determine position of weak acid equilibrium and

calculate pH).



Consider a solution made by mixing 0.10 mol of HCN (Ka = 6.2 × 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


CONCEPT CHECK!


Let’s Think About It…

Why isn’t NaOH a major species?

Why aren’t H+ and CN– major species?

List all possibilities for the dominant reaction.




The possibilities for the dominant reaction are:

1. H2O(l) + H2O(l) H3O+(aq) + OH–(aq)

2. HCN(aq) + H2O(l) H3O+(aq) + CN–(aq)

3. HCN(aq) + OH–(aq) CN–(aq) + H2O(l)

4. Na+(aq) + OH–(aq) NaOH

5. Na+(aq) + H2O(l) NaOH + H+(aq)




How do we decide which reaction controls the pH?

H2O(l) + H2O(l) H3O+(aq) + OH–(aq)

HCN(aq) + H2O(l) H3O+(aq) + CN–(aq)

HCN(aq) + OH–(aq) CN–(aq) + H2O(l)



HCN(aq) + OH–(aq) CN–(aq) + H2O(l)

What are the major species after this reaction occurs?

HCN, CN–, H2O, Na+




Now you can treat this situation as before.

List the possibilities for the dominant reaction.

Determine which controls the pH.



Calculate the pH of a solution made by mixing 0.20 mol HC2H3O2 (Ka = 1.8 × 10–5) with 0.030 mol NaOH in 1.0 L of aqueous solution.


CONCEPT CHECK!



What are the major species in solution?

Na+, OH–, HC2H3O2, H2O

Why isn’t NaOH a major species?

Why aren’t H+ and C2H3O2– major species?




What are the possibilities for the dominant reaction?

1. H2O(l) + H2O(l) H3O+(aq) + OH–(aq)

2. HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2–(aq)

3. HC2H3O2(aq) + OH–(aq) C2H3O2–(aq) + H2O(l)

4. Na+(aq) + OH–(aq) NaOH(aq)

5. Na+(aq) + H2O(l) NaOH + H+(aq)

Which of these reactions really occur?




Which reaction controls the pH?

H2O(l) + H2O(l) H3O+(aq) + OH–(aq)

HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2–(aq)

HC2H3O2(aq) + OH–(aq) C2H3O2–(aq) + H2O(l)

How do you know?




K = 1.8 × 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


Steps Toward Solving for pH

Ka = 1.8 × 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


Calculate the pH of a 100.0 mL solution of 0.100 M

acetic acid (HC2H3O2), which has a Ka value of 1.8 × 10–

5.

pH = 2.87


EXERCISE!


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 × 10–5, and 50.0 mL of a 0.10 M NaOH solution.

pH = 4.74


CONCEPT CHECK!


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

pH = 8.72


CONCEPT CHECK!


The pH Curve for the Titration of 50.0 mL of 0.100 MHC2H3O2 with 0.100 M NaOH



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



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

with 0.10 M HCl


Section 15.5Acid-Base Indicators

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).



The Acid and Base Forms of the Indicator Phenolphthalein



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



Useful pH Ranges for Several Common Indicators


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Chapter 15 Acid-Base Equilibria - HCC Learning Web