Grade 11

Grade 11Balance


Common ion effect


The common ion effect is an important concept in chemistry. It refers to the shift in equilibrium that occurs when a compound that has a common ion with the dissolved substance is added to a solution. It is a specific application of Le Chatelier's principle, which states that if a system in equilibrium is disturbed, it will adjust itself to counteract the disturbance and restore a new equilibrium.

Understanding the concept

The common ion effect can be explained through various examples and applications in chemistry. It is especially important when dealing with poorly soluble salts and weak acids or bases. Before diving into the examples, let's consider the equilibrium concepts involved.

Chemical equilibrium basics

Chemical equilibrium occurs when the forward and reverse reactions occur at the same rate. For a general reaction:

A + B ⇌ C + D

The equilibrium constant (K_c) for the reaction is defined as:

K_c = [C][D] / [A][B]

Now, when the common ion is introduced to the components of this equilibrium state, it affects the position of the equilibrium state by changing the concentration of the products or reactants.

Role of the common ion effect

When a salt, which shares a common ion with a solute already present in the solution, is added, the solubility of the solute decreases. This happens because the added common ion increases the ion concentration on one side of the equilibrium equation. As a result, the equilibrium will shift in favor of the formation of the reactant side to achieve a new equilibrium.

Example

Example 1: Common ion effect in solubility

Consider a saturated solution of calcium sulphate (CaSO_4):

CaSO_4 (s) ⇌ Ca^{2+} (aq) + SO_4^{2-} (aq)

Now, let's add another compound, such as sodium sulfate (Na_2SO_4), which provides an additional source of sulfate ions (SO_4^{2-}) in the solution. The equilibrium will shift to the left according to Le Chatelier's principle, resulting in a decreased solubility of CaSO_4.

CaSO₄ Na₂SO₄ Ca2+SO4

Example 2: Common ion effect with weak acids

Consider acetic acid (CH_3COOH) in water. The equilibrium expression is:

CH_3COOH (aq) ⇌ CH_3COO^- (aq) + H^+ (aq)

If sodium acetate (CH_3COONa) is added to the solution, it will dissociate and provide additional acetate ions (CH_3COO^-). This additional source of acetate ions will push the equilibrium to the left, resulting in fewer hydrogen ions being produced, which will decrease the acidity of the solution.

CH₃COOH CH₃COONa CH₃COO⁻ + H⁺

Mathematical representation

For a general equilibrium reaction involving a common ion, consider a solution of a weak electrolyte in water:

AB(s) ⇌ A^+ (aq) + B^- (aq)

Adding an electrolyte that shares a common ion, say B^−, shifts the equilibrium:

[B^-]_{initial} rightarrow [B^-]_{new} = [B^-]_{initial} + [common  ion]

With an increase in [B^-], the reaction quotient changes, and the system responds by shifting the equilibrium position to maintain K_c constant according to the equilibrium constant expression. This indicates a decrease in the solubility of the solute.

Applications of common ion effect

The common ion effect has many applications in chemistry, including:

  • Buffer solutions: Buffer solutions are able to resist changes in pH when small amounts of acid or base are added. This is because they contain a weak acid (or base) and its conjugate base (or acid), which produce a common ion effect.
  • Precipitation of salts: It is used to precipitate salts by adding a common ion to the solution, thereby decreasing the solubility of the salt in a controlled manner.
  • Industrial processes: The common ion effect is used in a variety of industrial processes where the control of solubility and precipitation is important, such as water treatment and purification.

Conclusion

The common ion effect is an interesting and highly relevant concept in the field of chemistry, especially for those who deal with equilibrium and solution chemistry. Understanding this effect provides insight into how the presence of a common ion affects the concentration of species in a solution, which ultimately affects solubility, pH, and the ability to control chemical processes.


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