Arrhenius Theory of Electrolytic Dissociation: Concept, Examples & Applications

The Arrhenius Theory of Electrolytic Dissociation is a fundamental concept in JEE Main Chemistry that describes how electrolytes split into ions when dissolved in water, enabling electrical conduction in solutions. Arrhenius’ work introduced the foundational ideas about acids, bases, and the mechanism of ionic conduction, directly supporting both theoretical and numerical problem-solving for JEE aspirants. Mastery of this topic is crucial for tackling questions involving the dissociation of salts, calculation of the degree of dissociation, the Van’t Hoff factor, and the difference between strong and weak electrolytes.

Definition and Introduction

Arrhenius Theory of Electrolytic Dissociation states that when an electrolyte dissolves in water, it spontaneously dissociates into positively charged cations and negatively charged anions. This separation into free ions makes the solution capable of conducting electricity. For example, NaCl in water dissociates as NaCl (aq) → Na⁺(aq) + Cl^-(aq).

Dissociation refers to the splitting of ionic compounds into ions, while ionization can also occur for some covalent compounds in water. Only substances that split into ions (electrolytes) enable electrical conduction; non-electrolytes do not dissociate and therefore cannot conduct electricity in solution.

Arrhenius’s Postulates: Theory Explained

• Electrolytes produce ions when dissolved in water, enabling electricity flow.
• Acids release H⁺ ions in aqueous solution, while bases release OH^- ions.
• The process is spontaneous and can be complete (strong electrolyte) or partial (weak electrolyte).
• Ions exist freely in dilute solutions but are held together by electrostatic forces in crystals.
• The extent of dissociation depends on the nature of the solute, solvent, concentration, and temperature.

Arrhenius Acid and Base Concept

This definition is strictly limited to aqueous solutions. For JEE, you must know the difference between this concept and Bronsted-Lowry or Lewis theory. Arrhenius acids and bases are identified by their ability to increase [H⁺] or [OH^-] in water, respectively.

Process of Electrical Conduction in Electrolytic Solution

When an electric field is applied to an aqueous electrolyte, free ions move towards the electrodes: cations towards the cathode and anions towards the anode. The movement of these ions constitutes current in the solution. The number and mobility of these ions determine the solution’s conductivity.

• Strong electrolytes dissociate completely—high conductivity (e.g., HNO₃, NaCl).
• Weak electrolytes dissociate partially—low conductivity (e.g., CH₃COOH, NH₄OH).
• Nonelectrolytes (e.g., glucose, urea) do not dissociate; no conduction occurs.

Factors Affecting Degree of Dissociation

• Nature of solute: Ionic compounds dissociate easily; covalent compounds may undergo ionization (e.g., HCl in water).
• Nature of solvent: Higher dielectric constant (as in water) facilitates dissociation.
• Concentration: Degree of dissociation increases as solution becomes more dilute (Ostwald's dilution law).
• Temperature: Increase in temperature usually increases ionization.

Difference Between Dissociation and Ionization

• Dissociation: Separation of pre-existing ions in an ionic compound (e.g., NaCl in water).
• Ionization: Formation of ions from molecules, often in covalent compounds (e.g., HCl forms H⁺ and Cl^- in water).
• Dissociation is usually reversible; ionization can be irreversible.
• Difference between electrolyte and non-electrolyte is central to this distinction.

Numerical Application: Van’t Hoff Factor and Degree of Dissociation

For JEE, quantitative application of Arrhenius theory involves the Van’t Hoff factor (i) and degree of dissociation (α). If a salt AB dissociates as AB → A⁺ + B^- and α is the fraction dissociated, then:

• Initial moles of AB: 1
• Moles dissociated: α
• Moles of A⁺: α; B^-: α; undissociated AB: 1–α
• Total particles: (1–α) + α + α = 1 + α

Thus, Van’t Hoff factor i = (total particles after dissociation) / (initial particles) = 1 + α.

Arrhenius Theory: Limitations and Criticisms

• Applies only to aqueous solutions, not to non-aqueous or gaseous media.
• Cannot explain basicity/acidity in non-water solvents (e.g., NH₃ as a base in liquid NH₃).
• Does not include substances without H⁺ or OH^- groups acting as acids or bases.
• Some salts and bases do not dissociate fully in water yet exhibit conductivity above predicted.
• Equilibrium between undissociated and dissociated ions in weak electrolytes is not fully addressed.

Later theories like Bronsted-Lowry theory and Lewis acid-base concept resolved many of these issues.

Key Takeaways and Exam Tips

• Memorize definitions of Arrhenius acid and base; use Arrhenius concept of acid and base in MCQs.
• Strong electrolytes: nearly complete dissociation (α ≈ 1); weak electrolytes: partial (α << 1).
• Calculate Van’t Hoff factor for questions about colligative properties and abnormal molar mass.
• Conduction by electrolytic solutions depends on ion mobility and quantity; nonelectrolytes give zero current.
• Distinguish carefully between dissociation (ionic) and ionization (covalent).
• Review degree of dissociation numericals and how dilution and temperature impact α for weak electrolytes.
• Check ionic equilibrium and conductivity problems for further practice.

Fast Revision: Arrhenius Theory of Electrolytic Dissociation vs. Other Theories

For deep dives, review acids, bases, and salts and chemical equilibrium sections. Vedantu provides in-depth notes and JEE Main exam patterns for all key chemistry chapters.

Solved Example: JEE Main Numerical

A sample contains 0.20 mol of NaCl dissolved in water. If the degree of dissociation (α) is 0.98, calculate the total moles of particles present after dissociation and the Van’t Hoff factor.

• Moles of NaCl initially = 0.20
• Moles dissociated: α × 0.20 = 0.196
• Moles Na⁺ ions formed: 0.196; Cl^- ions: 0.196; undissociated NaCl: 0.20 – 0.196 = 0.004
• Total particles = 0.004 + 0.196 + 0.196 = 0.396
• Van’t Hoff factor i = final particles / initial moles = 0.396 / 0.20 = 1.98

This approach is essential for abnormal colligative property questions and for distinguishing between strong and weak electrolytes in JEE Main Chemistry.

For more practice and topic-wise chapter links, check redox reactions and electrochemistry and equilibrium questions on the Vedantu website.

The Arrhenius Theory of Electrolytic Dissociation remains a critical starting point for understanding all acid-base theories. A strong foundation in this topic enhances your ability to tackle related JEE Main problems, especially those on ionic equilibria and solution conductivity.