Ionic Equilibria in Aqueous Systems

Chapter 19

Ionic Equilibria in Aqueous Systems

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The effect of addition of acid or base to an unbuffered solution

base added

acid added

The effect of addition of acid or base to a buffered solution

acid added

base added

2

An acid-base buffer is a solution that lessens the impact on pH from the addition of acid or base.

Most often, the components of a buffer are the conjugate acid-base pair of a weak acid (or base).

Buffers work through a phenomenon known as the common-ion effect .

CH 3 COOH( aq ) + H 2 O( l ) H 3 O + ( aq ) + CH 3 COO - ( aq )

If some CH 3 COO - ion is added, the equilibrium position shifts to the left; thus, [H 3 O + ] decreases, lowering the extent of acid dissociation.

Similarly, if acetic acid is dissolved in a sodium acetate solution, acetate ion and H 3 O + ion from the acid enter the solution. The acetate ion already present in the solution prevents the acid from dissociating as much as it would in pure water, thus lowering [H 3 O + ].

In this case, acetate ion is called the common ion .

The common-ion effect occurs when a given ion is added to an equilibrium mixture that already contains that ion, and the position of the equilibrium shifts away from forming more of it.

2

How a buffer works

2

The Henderson-Hasselbalch Equation

2 3 + -

3 + -

K a

3 + K a

-

-

3 + K a

[base]

pH = p K a + log

[acid]

6

Buffer Capacity and Buffer Range

Buffer capacity is the ability to resist pH change

A buffer has the highest capacity when the component concentrations are equal (in other terms, its pH ≈ p K a of its acid component)

Buffer range is the pH range over which the buffer acts effectively

Buffers have a usable range within ± 1 pH unit of the p K a of its acid component.

6

The relation between buffer capacity and pH change

6

Colors and approximate pH range of some common acid-base indicators

An acid-base indicator is a weak acid, with a different color than its conjugate base and with the color change occuring over a specific and narrow pH range.

6

The color change of the indicator bromthymol blue

Basic

Acidic

6

6

Curve for a weak acid-strong base titration.

6

Sample Problem 19.4

Calculate the pH during the titration of 40.00 mL of 0.1000 M propanoic acid (HPr; K a = 1.3 x 10 -5 ) after adding the following volumes of 0.1000 M NaOH:

(a) 0.00 mL

(b) 30.00 mL

(c) 40.00 mL

(d) 50.00 mL

The amounts of HPr and Pr - will be changing during the titration. Remember to adjust the total volume of solution after each addition.

Find the starting pH using the methods of Sample Problem 18.8.

[Pr - ][H 3 O + ]

[Pr - ] = x = [H 3 O + ]

K a =

[HPr]

x = 1.1 x 10 -3 ; pH = 2.96

HPr( aq ) + OH - ( aq ) Pr - ( aq ) + H 2 O ( l )

-

0.004000

0

-

Initial

0.003000

-

Change

-

-

Final

0

-

0.001000

0.003000

13

Sample Problem 19.4

continued (2 of 3)

0.001000 mol

[H 3 O + ] = 1.3 x 10 -5 x

= 4.3 x 10 -6 M

pH = 5.37

0.003000 mol

(c) When 40.00 mL of NaOH are added, all of the HPr will be reacted and the [Pr - ] will be

(0.004000 mol)

= 0.05000 M

(0.004000 L) + (0.004000 L)

1.0 x 10 -14

K w

K a x K b = K w

K b = = = 7.7 x 10 -10

1.3 x 10 -5

K a

K w

[H 3 O + ] = = 1.6 x 10 -9 M

pH = 8.80

14

Sample Problem 19.4

continued (3 of 3)

(d) 50.00 mL of NaOH will produce an excess of OH - .

mol excess base = (0.1000 M )(0.05000 L - 0.04000 L) = 0.00100 mol

1.0 x 10 -14

M = (0.00100 mol)

M = 0.01111

[H 3 O + ] = = 9.0 x 10 -13 M

(0.0900 L)

0.01111

pH = 12.05

15

Curve for a weak base-strong acid titration

15

Formation of acidic precipitation

15