CHM1 19 The Common Ion Effect Collection

Processing...

The Common Ion Effect

This page was also found in Chapter 19, but is pertinent here so is included in this chapter as well.

Le Châtelier's Principle states that if an equilibrium gets out of balance, the reaction will shift to restore the balance. If a common ion is added to a weak acid or weak base equilibrium, then the equilibrium will shift towards the reactants, in this case the weak acid or base.

Introduction

The common-ion effect is used to describe the effect on an equilibrium involving a substance that adds an ion that is a part of the equilibrium.

Example 5

Consider the lead(II) ion concentration in this saturated solution of PbCl₂. The balanced reaction is

PbCl₂(s) ? Pb²⁺(aq) + 2Cl^-(aq)

Assuming the concentration of dissolved lead(II) chloride is s, then:

[Pb²⁺] = s

[Cl^-] = 2s

Put these values into the solubility product expression, and do the sum.

/attachments/19f73a7a-f145-11e9-8682-bc764e2038f2/kspcalc1.gif

So the concentration of lead(II) ions in the solution is 1.62 x 10^-2 M. What happens if you add some sodium chloride to this saturated solution? Sodium chloride shares an ion with lead(II) chloride. The chloride ion is common to both of them; this is the origin of the term "common ion effect".

Look at the original equilibrium expression again:

PbCl₂(s) ? Pb²⁺(aq) + 2Cl^-(aq)

What would happen to that equilibrium if you added extra chloride ions? According to Le Châtelier, the position of equilibrium will shift to counter what you have just done. In this case, it would tend to remove the chloride ions by making extra solid lead(II) chloride.

Note: Actually, of course, the concentration of lead(II) ions in the solution is so small to start with, that only a tiny proportion of the extra chloride ions can be converted into solid lead(II) chloride.

The lead(II) chloride will become even less soluble - and, of course, the concentration of lead(II) ions in the solution will decrease. Something similar happens whenever you have a sparingly soluble substance. It will be less soluble in a solution which contains any ion which it has in common. This is the common ion effect.

A Simple Example

Suppose you tried to dissolve some lead(II) chloride in some 0.100M sodium chloride solution instead of in water. What would the concentration of the lead(II) ions be this time? As before, let's call the concentration of the lead(II) ions s.

/attachments/19f73a7a-f145-11e9-8682-bc764e2038f2/P1_CIONEf_1.png

Common Ion Effect on Weak Acids and Bases

Adding a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion. The common ion effect suppresses the ionization of a weak acid by adding more of an ion that is a product of this equilibrium.

Example 6
The common ion effect of H₃O⁺ on the ionization of acetic acid The common ion effect suppresses the ionization of a weak base by adding more of an ion that is a product of this equilibrium.

Example 7

Consider the common ion effect of OH^- on the ionization of ammonia

/attachments/19f73a7a-f145-11e9-8682-bc764e2038f2/common_ion_2.jpg

Therefore, adding the common ion of hydroxide, shifts the reaction towards the left to decrease the stress (LeChâtelier's Principle), forming more reactants. This decreases the reaction quotient, for the reaction is being pushed towards the left to reach equilibrium. The equilibrium constant, K_b=1.8*10^-5, does not change. The reaction is put out of balance, or equilibrium.

Q_A= [NH⁴⁺][OH^-]/[NH₃]

At first, when more hydroxide is added, the quotient is greater than the equilibrium constant. The reaction then shifts right, causing the denominator to increase, decreasing the reaction quotient and pulling towards equilibrium and causing Q to decrease towards K.

? Previous Section

Index Page Next Section ?