Arrhenius Theory of Electrolytic Dissociation for IIT JEE Chemistry
Characteristics of Electrolytic dissociation IIT JEE Mains
In 1880s, Svante Arrhenius set the foundation for the theory of electrolytic dissociation. He was awarded the Noble Prize in 1903 for this theory, after which the theory gained importance. Based on the original theory, if fraction α mole of an electrolyte dissociates in water, it forms into 2α moles of ions, rest 1- α being the undissociated form. The assumption was based on the fact that every mole of salt, on dissolving in water, forms
(1- α) + 2α = (1+ α) = I, where I is the Vant Hoff factor.
The measure of degree of dissociation (α) is Arrhenius conductivity ratio, which is the ratio of equivalent conductivity at any given concentration t at infinite dilution. This theory proved to be successful and the supporters of Arrhenius – Vant Hoff and Oswald and were later on known as ‘Ionists’.
The modifications and applicability of theory was tried to extend to concentrated solutions and allowance and idea of free water was developed. But eventually Debye-Huckel theory of interionic interaction of complete dissociation, explained the concentration dependence of activity coefficients for dilute solutions, and later on extended to higher concentrated electrolytes. This resulted into complex equations with unknown parameters and no explanation of non-idealistic over the entire concentration range.
Gradually after years it was confirmed that the degree of dissociation and hydration numbers evaluated from vapour pressure data, instead of conductivity ratio. This explained and proved the non-ideal properties of electrolytes over large concentration range. Further it was found that the modalities and hydration number of free waters were different at surface and in the bulk of solution. This led to the application of the theory of electrolytes to the whole range of concentration from zero to saturation based on the idea of partial dissociation and free water.
Postulate states that; “in aqueous solution, the molecules of an electrolyte undergo spontaneous dissociation to form positive and negative ions.” Best example being NaCl, being dissociated into Na+ and Cl-
Dissociation simply means breaking up of compound into simpler constituents, which can recombine again under other conditions. In ionic or electrolytic dissociation, the addition of electrolyte or a solvent causes the molecules of the compound to break-up into ions (electrically charged particles). The dissociation property is used to explain electrical conductivity of the electrolyte and the compound.
It is assumed, according to the modern theory that solid electrolytes consist of two types of charged particles – positive and negative, which are held together by the electrostatic force of attraction. When these solid electrolytes are dissolved in the appropriate solvent, electrostatic force between charged particles is weakened, leading into the separation or dissociation into single charged entity. This is known as electrolytic dissociation or ion solvation. Based on the capability of electrolytes, its types are as follows:
1) Strong Electrolytes – Those electrolytes which dissociate completely into their respective ions, even at moderate conditions are called strong electrolytes. Their degree of dissociation is high, and their dissociation constant is simultaneously high. This type of electrolytes have high conductivity. Law of mass action is inapplicable as the dissociation is irreversible.
2) Weak electrolytes – Those electrolytes which dissociate to a limited extent are called weak electrolyte. These electrolytes have low degree of ionisation and lower dissociation constant value. They have low electrical conductivity. The dissociation is reversible; hence law of mass action is applicable. Example: acetic acid, formic acid, weak base like ammonium hydroxide and salts like ammonium acetate and silver acetate.
Characteristics of Electrolytic dissociation:
1. Dissociation is the process of separation of charged particles which already exist in a compound.
2. Dissociation involves ionic compounds.
3. Dissociation will either produce charged particles or electrically neutral particles.
4. Dissociation is reversible.
5. Dissociation is possible only when there are ionic bonds in a compound.
Difference between concepts of Ionisation and Dissociation:
The major difference between the two is the type of compounds involved.
Ionisation – The process of formation of ions from compounds which are not ionic in nature. It involves covalent compounds. It is irreversible in nature. Appropriate solvent is required to start the process of ionisation and it is also called ion solvation.
Example: In case of HCl molecule, H and Cl atoms are covalently bonded. However, upon dissolving it in water, it forms two ions, namely H+ and Cl– ions. HCl(aq) → H+ (aq) + Cl–(aq)
Dissociation – It is the process of spontaneous splitting of substance into constituent charged particles. The compound required must be ionic in nature. They are reversible in nature.
Example: In case of sodium chloride (NaCl) molecule NaCl(aq) → Na+ (aq) + Cl–(aq)
Passage of current through electrolytic solution causes asymmetry in the ionic atmosphere. Central ion moves towards the electrode and solvent molecule in the opposite direction. Due to the large number of either of the charge, the charge density increases at one end. Thiscauses decreased conductance, however, symmetry is achieved after a short time. The effect is represented: B=8.2x105Λ0/(DT)3/2
D is dielectric constant, η is viscosity in poises, T is absolute temperature
A single ion is surrounded by solvent molecule and other ions, thus ionic atmosphere of central ion involves forces of both. The movement of central ion in a direction opposite to that of ionic atmosphere causes the withdrawal force by solvent molecule on the movement of the central ion. This new retarding force on central ion due to friction between ion and solvent is known as electrophoretic effect. This causes decrease in the equivalent existing conductance. The electrophoretic force can be mathematically represented in the following equation:
A=82.4/(DT)1/2η D is dielectric constant, η is viscosity in poises, T is absolute temperature.
Degree of Dissociation:
The fraction of the total number of moles of weak electrolyte which ionises into respective ions in an aqueous solution at equilibrium state is called as the degree of dissociation. It is denoted by ‘α’. It can be represented in equation as:
Degree of dissociation and its value is found to be dependent on following factors:
a) Nature of Solute: If the ionisable part of the molecule is bonded with covalent bonds, less ions are produced. And if the ionisable part of molecule is held by electro-covalent bond, more ions are produced.
b) Nature of Solvent: The solvent is solely responsible to reduce electrostatic attraction force between two charged particles (ions). It is well known by Coulomb’s law, that forces between two charged particles are inversely proportional to the dielectric constant of the medium between them. Thus, more the dielectric constant, the more is the capacity of the solvent to separate the ions. Water has the highest dielectric constant, thus being the best solvent for dissociation of charged particles.
c) Concentration of Solution: By Ostwald’s dilution law “The degree of ionisation of any weak electrolyte is inversely proportional to the square root of concentration and directly proportional to the square root of dilution”. Thus, it implies that if the dilution of particular substance increases, logically it means more addition of solvent (concentration decreases). The degree of ionisation increases, as more molecules of the solvent cause more formation of ions.
d) Temperature: The temperature is directly proportional to the degree of dissociation. As the temperature increases, the kinetic energy of molecules increases leading to a decrease in the inter-particle attraction force and results into mores dissociation of ions.
Evidences in support of Arrhenius theory:
• X-ray diffraction studies show the presence of ions in electrolytes. It also shows that they conduct electricity in fused state.
• Electrolytic solutions obey Ohm’s law. This is possible especially if ions are present in the solution already.
• Some reactions are possible due to the presence of ions and ionic compounds:
• Based on Arrhenius theory, undissociated water is obtained in some system, which leads to the change in the enthalpy of the system. The phenomenon is known as enthalpy of neutralisation.
• The colour of electrolyte is due to the presence of ion.
• This theory forms the basis of solubility product, hydrolysis, common-ion effect, electrolysis, electrical conductivity, electrophoresis, etc.
• The ionic theory can explain the abnormal and unpredictable colligative properties. When electolyte gets dissolved in water, number of particle in the solution always increase than total number of molecules are dissolved due to ionization.
Limitations of Arrhenius theory:
• Arrhenius theory is applicable to aqueous solutions and not to non-aqueous solutions and gaseous solutions, as it defines electrolyte in terms of aqueous solution and not as a substance.
• The role of solvent is not responsible for deciding the nature of strength of an electrolyte. Example: HCl is a strong acid in the presence of water but it is a weak acid in the presence of benzene.
• Organic solvents have not been explored as much as non-organic solvents.
• Theories based (Ostwald’s dilution law) on Arrhenius theory of dissociation have proved to be effective only for weak electrolytes.
• It is proved and observed that in the absence of water also, strong electrolyte conducts electricity. This is found to be contradictory to the Arrhenius theory.
• Factors affecting the degree of dissociation are not very well-explained.