Chapter 16

1

B U F F E R S S U R O V I E C S P R I N G 2 0 1 4

I. Buffer Solutions 2 A. Buffer is a solution that resists a change in pH with the addition of small amounts of acid or base  Examples of buffers:   Phosphate Acetic

Example 3  What is the pH of the buffer solution that contains 2.2 grams of NH 4 Cl in 250mL of 0.12M NH 3 ?

B. General expression for buffer solutions 4  In the previous example we found the [H 3 O + ] by solving for x in:

B. General expression for buffer solutions 5  Henderson – Hasselbalch equation  The equation shows that pH is controlled by 2 factors 1.

2.

Strength of acid (K a or pK a ) The two species are present in the same solution, the ratio of their concentrations is also their mole ratio.

C. Using the H – H equation 6  Calculate the pH of a 1.0L solution that has an 0.050M acetic acid and 0.075M sodium acetate. The K a of acetic acid is 1.8 X 10 -5

D. Preparing a buffer solution 7  1.

To be a useful a buffer must have 2 characteristics Be able to control the pH at the desired value 2.

The buffer should have the capacity to control the pH after the addition of small amounts of acid/base

Example 8 I want a buffer of pH 2.50, which of the following can I use to make the buffer and in what ratio?

1.

NaCl and HCl 2.

3.

CH 3 CO 2 H and NaCH 3 CO 2 H 3 PO 4 and NaH 2 PO 4

Example 9  Given a 25.00mL of 0.100M NaOH and 35.00mL of 0.125M HC 2 H 3 O 2 will this make a buffer?

E. Effect of added H 3 O + or OH on Buffer Systems 10  1.

A buffer solution was prepared by adding 4.95 g of NaCH 3 CO 2 H to 250mL of 0.150M acetic acid. What is the pH of the buffer initially?

2.

What is the new pH when 82mg of NaOH is added to a 100 mL aliquot of the buffer?

II. Acid/Base Titrations 11  A titration is one of the most powerful and accurate ways we have to determine the quantity of an acid/base in solution.

A. Strong acid – strong base 12  What happens to the pH as 0.10M HCl is slowly added to 50.0mL of 0.10M?

A. Strong acid – strong base 13

Volume of HCl Added (mL)

0 25 50

pH

B. Titration of Weak Acid and Strong Base 14  There are 3 important points to a titration: 1.

pH before base added is calculated from the weak acid K a and the acid concentration 2.

pH at the equil. point can be calculated from the conjugate base as the conjugate acid and the strong base has been consumed 3.

pH at the ½ equil. point is equal to the pK a weak acid of the

C. Titration of weak base with strong acid 15 Given a 25.0 mL sample of 0.10M NH 3 a.

being titrated with 0.10M HCl What is the pH of the solution before the titration begins?

b.

c.

What is the pH at the equilivance point?

What is the pH at the midpoint?

III. pH Indicators 16   An indicator is an organic compound that is itself a weak acid or weak base In aqueous solution the acid from is in equilibrium with its conjugate base.

IV. Solubility of Salts  Precipitation reactions are reactions in which one of the products are water soluble

A. The Solubility Product Constant, K sp  K sp reflects the solubility of a compound and is usually called the solubility product constant

B. Solubility and K sp  K sp is measured by experiments in the lab determining concentrations of ions in solution

C. Soluability and the Common Ion  What if we had a test tube of saturated AgNO 3 added more KNO 3 ?

and I

Example  What is the final molarity of AgSCN (s) if it is placed in   Water 0.010 M NaSCN

D. Common Ion Concepts  The solubility of a salt will always be reduced by the presence of a common ion  We made the approximation that the amount of common ion added to the solution will also be large compared with the amount of ion coming from the insoluble salt.

E. Precipitation Reactions  How do we know when a precipitate will form?

 Look at Q and how is Q related to K sp