# CHAPTER 15: ACID-BASE - COMMON IONS, BUFFERS AND

3

+

-

sp

### problems)

A “common ion” is an ion present in an equilibrium system in high enough concentration that it affects the pH of the weak acid or weak base.

### Exercise 14.19 Common Ion Problem

What is the pH of a mixture containing 1.00

M

of HF and 1.00

M

of NaF?

Explain these results in terms of Le Chatelier’s Principle.

### Objective: Relate the predominant form of chemical species to pH and pKa.

Weak Acid and Weak Base Solutions

Not a Buffer Buffer

### 2

Observations of a buffer solution

(website: http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf

)

### Protonation state

:

In summary: A buffer – a) is a mixture of -

b) contains a ratio of acid and base forms such that -

c) can maintain a pH close to -

c) has a -

### Buffer capacity

is the ability of a buffer to buffer (

ha

) small changes in pH. Adding small amounts of strong acid or base should not affect buffer pH much IF the buffer capacity is high/strong/good.

Why is that ratio necessary? Use an equilibrium constant expression to explain your answer.

### Exercises

A solution contains 0.50

M

acetic acid (K a

= 1.8 x 10

-5

) and 0.50

M

sodium acetate. a) Explain why this solution could act as a buffer. b) Around what pH will this act as an effective buffer? _______________ Use a hydrolysis equation and equilibrium-constant expression to justify your answer. c) Examine the equilibrium-constant expression you wrote in b). When is the concentration of the acid form higher than the base form? (i.e. pH < pKa, pH = pKa, pH > pKa?) d) Examine the equilibrium-constant expression you wrote in b). When is the concentration of the base form higher than the acid form? (i.e. pH < pKa, pH = pKa, pH > pKa?) e) Write a net ionic equation for the reaction between the buffer and strong acid (i.e. HCl) f) Write a net ionic equation for the reaction between the buffer and strong base (i.e. NaOH)

### 3

The pH of a sodium acetate-acetic acid buffer is 4.50. Which form is more predominant in solution, acetate or acetic acid? What is the ratio of acetate to acetic acid in this buffer? K a

for acetic acid = 1.8 x 10

-5

.

The pH of an ammonia/ammonium chloride buffer is 9.50. Which form is more predominant in solution, ammonia or ammonium? What is the ratio of ammonia to ammonium in this buffer? Kb for ammonia =

1.8 x 10

-5

.

Write a net ionic equation for the reaction between the ammonia/ammonium buffer and strong acid (i.e.

HCl)

Write a net ionic equation for the reaction between the ammonia/ammonium buffer and strong base (i.e.

NaOH)

2.

1.

3.

Objective: Design a buffer solution with a target pH and buffer capacity.

Three routes to making an effective buffer solution:

### For acidic buffers, the pKa of the weak acid to be used in the buffer should be as close as possible to the desired pH. (Same goes for basic buffers.)

For example, if you need a buffer with a pH of 4.00, the best choice of a weak acid for this buffer should have a pKa as close to 4.00 as possible.

### Exercise

A chemist needs a solution buffered at pH 4.30 and can choose from the following acids: a. propanoic acid (K a

= 1.3 x 10

-5

) b. benzoic acid (K a

= 6.4 x 10

-5)

Which weak acid will work best? Use a calculation as justification.

### Exercise

A buffered solution is made by adding 75.0 g of sodium acetate to 500.0 mL of a 0.64 M solution of acetic acid. What is the pH of the final solution? Ka of acetic acid = 1.8 x 10

-5

You want to make a buffer out of 1.00 L of a 0.75 M solution of NH

3

(K b

= 1.8 x 10

-5

). What would you have to add, and how many moles of it would you have to add? Assume no volume change will occur.

How many moles of sodium hydroxide would have to be added to 1.0 L of 2.0 M acetic acid to produce a solution buffered at pH = 5.00?

### 5

AP Problems – Buffers

1992 D

The equations and constants for the dissociation of three different acids are given below.

HCO

H

2

PO

3

4

↔ H

↔ H

+

+

+ CO

3

2–

+ HPO

4

2–

K a

= 4.2×10

–7

K a

= 6.2×10

–8

HSO

4

↔ H

+

+ SO

4

2–

K a

= 1.3×10

–2

(a) From the systems above, identify the conjugate pair that is best for preparing a buffer with a pH of

(b) Explain briefly how you would prepare the buffer solution described in (a) with the conjugate pair you have chosen.

(c) If the concentrations of both the acid and the conjugate base you have chosen were doubled, how would the pH be affected? Explain how the capacity of the buffer is affected by this change in concentrations of acid and base.

(d) Explain briefly how you could prepare the buffer solution in (a) if you had available the solid salt of the only one member of the conjugate pair and solution of a strong acid and a strong base.

1993 A

### 6

CH

3

NH

2

+ H

2

O ↔ CH

3

NH

3

+

+ OH

Methylamine, CH

3

NH

2

, is a weak base that reacts according to the equation above. The value of the ionization constant, K b

, is 5.25

10

–4

. Methylamine forms salts such as methylammonium nitrate,

(CH

3

NH

3

+

)(NO

3

).

(a) Calculate the hydroxide ion concentration, [OH

], of a 0.225–molar aqueous solution of methylamine.

(b) Calculate the pH of a solution made by adding 0.0100 mole of solid methylammonium nitrate to 120.0 milliliters of a 0.225–molar solution of methylamine. Assume no volume change occurs.

(c) omitted

(d) A volume of 100. milliliters of distilled water is added to the solution in (c). How is the pH of the solution affected? Explain.

2009 part A, question #1

Answer the following questions that relate to the chemistry of halogen oxoacids.

Acid

HOCl

K a

at 298 K

2.9

10

-8

(d) A buffer solution is prepared by dissolving some solid NaOCl in a solution of HOCl at 298 K. The pH of the buffer solution is determined to be 6.48.

(i) Calculate the value of [H

3

O

+

] in the buffer solution.

(ii) Indicate which of HOCl

(aq)

or OCl

-

(aq)

is present at the higher concentration in the buffer solution. Support your answer with a calculation

2011 A

### 7

Each of three beakers contains 25.0 mL of a 0.100

M

solution of HCl, NH

3

, or NH

4

Cl, as shown above.

Each solution is at 25°C.

(a) Determine the pH of the solution in beaker 1. Justify your answer.

(b) In beaker 2, the reaction NH

3

NH

3

(aq)

is 1.8

10 -5 at 25°C.

(aq)

+ H

2

O

(l)



NH

4

+

(aq)

+ OH –

(aq)

occurs. The value of

K b

for

(i) Write the

K b

expression for the reaction of NH

3

(aq)

with H

2

O

(l)

.

(ii) Calculate the [OH – ] in the solution in beaker 2.

(c) In beaker 3, the reaction NH

4

+

(aq)

+ H

2

O

(l)



NH

3

(aq)

+ H

3

O +

(aq)

occurs.

(i) Calculate the value of

K a

for NH

4

+

(aq)

at 25°C.

(ii) The contents of beaker 2 are poured into beaker 3 and the resulting solution is stirred. Assume that volumes are additive. Calculate the pH of the resulting solution.

(d) The contents of beaker 1 are poured into the solution made in part (c)(ii). The resulting solution is stirred. Assume that volumes are additive.