Electrolytic Cells And Electrolysis

The conductivity of acetic acid (CH₃COOH) decreases on dilution because it is a weak electrolyte that ionizes only partially in solution. This behavior can be understood by looking at the properties of weak electrolytes.

1. Weak Electrolyte Behavior:

Acetic acid is a weak electrolyte, meaning that it does not fully ionize in solution. The ionization of acetic acid in water is represented as: This reaction is reversible, and only a small fraction of acetic acid molecules dissociate into ions.

2. Effect of Dilution:

When acetic acid is diluted with water, the concentration of both dissociated ions ( and ) decreases. Since conductivity ( ) depends on the number of ions in the solution, the conductivity decreases as the solution becomes more dilute.

In the case of weak electrolytes like acetic acid, dilution shifts the equilibrium of the ionization reaction to the right, meaning more acetic acid dissociates. However, despite the increase in ionization, the overall number of ions still decreases with dilution, which results in a net decrease in conductivity.

3. Explanation of Decreased Conductivity:

As the solution is diluted, the number of dissociated ions decreases, reducing the electrical conductivity. This is because the conductivity is directly proportional to the concentration of ions in solution. Therefore, even though the ionization of acetic acid might increase slightly upon dilution, the overall ion concentration is lower, leading to reduced conductivity.

Conclusion:

In summary, dilution of acetic acid decreases the conductivity because it is a weak electrolyte that ionizes partially. As the concentration of ions decreases, the overall conductivity of the solution decreases as well.

Electrolysis of Aqueous NaCl

1. Standard Electrode Potentials:

The standard electrode potential for the reduction of oxygen to water is: While for the reduction of chloride ions to chlorine gas:

2. Why Chlorine is Liberated at the Anode:

Although the standard electrode potential for chloride ions is slightly more positive than that for oxygen, this does not automatically mean that oxygen will be reduced to water at the anode in the electrolysis of aqueous NaCl.

The reason why chlorine gas ( ) is liberated at the anode during electrolysis is due to the **concentration of ions** in the solution. Chloride ions ( ) are present in **much higher concentrations** compared to oxygen in the solution.

The **concentration effect** plays a significant role in determining which ion is more readily oxidized. Even though oxygen has a more positive reduction potential, the abundance of chloride ions ensures that they are more readily oxidized to form chlorine gas at the anode.

3. Explanation of the Process:

During the electrolysis of aqueous NaCl, the **anode** is the site of oxidation, where electrons are released. At the anode, chloride ions ( ) are oxidized to chlorine gas ( ): At the **cathode**, water is reduced to form hydroxide ions ( ) and hydrogen gas ( ):

4. Conclusion:

Despite the standard electrode potential being more positive for oxygen reduction than for chloride ion reduction, chlorine is preferentially liberated at the anode during the electrolysis of aqueous NaCl. This is because the concentration of chloride ions in the solution is much higher, making them more readily available for oxidation.

Kohlrausch's Law of Independent Migration of Ions

The Kohlrausch law of independent migration of ions states that the molar conductivity of an electrolyte at infinite dilution is the sum of the individual contributions from its cation and anion.

This implies that each ion contributes independently to the total conductivity, and the total can be obtained by simply adding the limiting molar conductivities of each ion.

Mathematical Expression:

Where:

• : Molar conductivity at infinite dilution
• : Limiting molar conductivity of the cation
• : Limiting molar conductivity of the anion

This law is particularly useful for calculating unknown molar conductivities and analyzing electrolyte behavior in dilute solutions.