Problems with Buffer Solutions

Predict whether a solution is a buffer solution

Have you ever noticed, when you drink a large amount of lemonade or orange juice, your blood doesn’t become acidic?  Your body is controlling the pH of your blood, because your blood is a buffer.

The definition of a buffer:
  • A solution that contains a weak acid/conjugate base mixture or a weak base/conjugate acid mixture
  • Buffer solutions resist changes in pH when a moderate amount of an acid or base is added

Buffer solutions can be made in two different ways:1.Weak Acid and one of its salts2.Weak Base and one of its salts Notice that it must be a weak acid or a weak base.
Acids and bases are removed by the weak acid/base and its saltsBuffer Capacity – the amount of acid or base that can be added before considerable change occurs to the pH

Example:

Which of the following are buffers?

  1. CH3COOH and CH3COONa
  2. HCl and NaCl
  3.  NH3 and NH4Cl
  4. NH3 and KCl

Solution:

Only 1 & 3 (CH3COOH and CH3COONa and NH3 and NH4Cl) because for 1) acetic acid is a weak acid and CH3COONa contains its weak base – CH3COO– )

For 3, NH3 is a weak base, and NH4Cl contains its conjugate weak acid, NH4+

Calculate the pH of a buffer solution

A common ion effect of an acid can be introduced when the salt of a weak acid is in the same solution as a weak acid.   You must account for the present of the reactant and product in your initial concentrations.

Conveniently for weak acids, these calculations can be worked out more simply using the Henderson Hasselbach equation:

where pKa = – log (Ka)

[A–] = the concentration of the salt 

[HA] = the initial concentration of the weak acid

Example: 

Find [H+] in a solution 1.0M HNO2 and 0.225M NaNO2. The Ka for HNO2 is 5.6 x 10-4.

Solution:

This can be set up as an ICE table, but more easily you can set this up as a Henderson Hasselback equation:

Prepare a buffer solution of a desired pH

If you would like to prepare a buffer of a desired pH, you must find an acid or a base with a pKa that is close to it.  When you are close, you can then use the Henderston-Hasselback equation to find the correct ratio for the pH.

Example:

Prepare a buffer with a pH of 5.0

Solution:

Examination of some common acids, you will find that acetic acid (CH3COOH) has a pKa of 4.74. (Ka of acetic acid is 1.75 x 10-5)

Next:  find the ratio of [A-]/[HA] will create an acetic acid buffer of pH 5.0 

That means a ratio of 1.74 is needed, which could mean to use 1.74 M of  CH3COONa and 1.0 M of CH3COOH

Apply the Henderson-Hasselbalch equation

If a  buffer is working, it should resist change when a small amount of acid or base is added to a solution.  This can be calculated conveniently using the Henderson Hasselback equation.

To solve these problems:

  • Assume that the reaction goes to completion and carry out the stoichiometric calculations
  • Carry out the equilibrium calculations

Example:

Calculate the change in pH that occurs when a 0.010 mole of solid NaOH is added to 1.0 L of the buffered solution that contains 0.50 M acetic acid (CH3COOH or HC2H3O2, Ka = 1.8×10–5) and 0.50 M sodium acetate (CH3COONa or NaC2H3O2)

Solution:

**Notice  that 0.010 mole of HC2H3O2 has been converted to 0.010 mole of C2H3O2– by the added OH–
The addition of 0.010 mole of OH– has consumed HC2H3O2 and produced C2H3O2–

0.49/1.8X10-5 > 500, and the assumption can be used.

[H+] = 1.7 X 10-5pH = 4.76
This is only slightly more basic than the pKa of 4.74, and evidence of the capabilities of a buffer.