Introduction of Arrhenius Theory of Electrolyte Dissociation
The Arrhenius theory of electrolyte dissociation was based on the observation that the ionisation potential of water was greater than that of sodium chloride and KCl. This was the reason that, when two solutions of NaCl and KCl were mixed, ions of Cl-, and OH- ions were generated. In other words, it was observed that the Cl- ion concentration of solution decreased as NaCl and KCl were mixed. This is called the Arrhenius dissociation theory.
In 1885, Henry demonstrated that electrolyte ions present in dilute solutions can exist in stable species. The dissociation constant, K, the value of all electrolytes increases with their concentration. Henry showed that when dilute solutions of HCl and NH3 are mixed, the ions of H- and OH- can exist in stable species. electrolytes are Lewis acids. Henry's theory was not considered as a theory of electrolyte dissociation but as an explanation of the selective conductivity of solutions of chloride ions in comparison with those of water. electrolytes dissociation theory is based on the selective conductivity of solutions of chloride ions in comparison with those of water.
Henry's theory was not considered as a theory of electrolyte dissociation but as an explanation of the selective conductivity of solutions of chloride ions in comparison with those of water. 20, 21 Henry's theory also allowed chemists to calculate the conductivity of dilute aqueous solutions of chloride ions in comparison with those of water, which is often called Henry's theory. Henry's theory also allowed chemists to calculate the conductivity of dilute aqueous solutions of chloride ions in comparison with those of water, which is often called Henry's law. electrolytes dissociation theory is based on the selective conductivity of solutions of chloride ions in comparison with those of water.
Henry's theory also allowed chemists to calculate the conductivity of dilute aqueous solutions of chloride ions in comparison with those of water, which is often called Henry's law. Henry's theory also allowed chemists to calculate the conductivity of dilute aqueous solutions of chloride ions in comparison with those of water, which is often called Henry's law. 2.2.1 Henry's Law was established by Henry in 1807 and was first expressed as Equation 1 The conductivity of dilute solutions of a solute (solute ion) at a constant temperature, depends on the concentration of solute ions, or, Equation 2
Henry's law
Henry's law is generally expressed as 1 T = R·n1·e−[K+]·[Cl-]·(s)/(RT) Substituting Equation 2 into Equation 1, one gets: T=R·n1·e −[K+]·[Cl-]·(s)/(RT) Equation 3 Thus, the conductivity of solutions can be calculated by measuring the concentration of solute ions, and temperature, (in Kelvin) Solubility of Cl− in Water at Different Temperatures and Concentrations Solubility of Cl- in Water at Different Temperatures and Concentrations Solubility of Cl- in Water at Different Temperatures and Concentrations Since cm represents the concentration of mobile ions (chloride ions) in water, it must be known to calculate conductivity.
Svante Arrhenius
In the 1880s, Svante Arrhenius set the foundation for the theory of electrolytic dissociation. He was awarded the Nobel Prize in 1903 for this theory, after which the theory gained importance. Based on the original theory, if a fraction α mole of an electrolyte dissociates in water, it forms into 2α moles of ions, and rests 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 Van't Hoff factor.
The measure of the degree of dissociation (α) is the 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 – Van't Hoff and Oswald and were later on known as ‘Ionists’.
The modifications and applicability of theory were tried to extend to concentrated solutions and allowance and the 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 in 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 were evaluated from vapour pressure data, instead of conductivity ratio. This explained and proved the non-ideal properties of electrolytes over a large concentration range. Further, it was found that the modalities and hydration number of free waters were different at the surface and in the bulk of the 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.” The best example is NaCl, being dissociated into Na+ and Cl-
NaCl(aq) → Na+ (aq) + Cl-(aq)
H2SO4 (aq) → 2 H+ (aq) + SO42-(aq)
Definition of Electrolytic Dissociation
Dissociation simply means breaking up compounds into simpler constituents, which can recombine again under other conditions. In ionic or electrolytic dissociation, the addition of an 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 the electrical conductivity of the electrolyte and the compound.
Overview
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, the electrostatic force between charged particles is weakened, leading to the separation of dissociation into a single charged entity. This is known as electrolytic dissociation of ion solvation. Based on the capability of electrolytes, their types are as follows:
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 electrolyte has high conductivity. The Law of mass action is inapplicable as the dissociation is irreversible.
Example: strong acids – H2SO4, HCl, HNO3
Strong bases – NaOH, KOH
Salts – NaCI, KCl
Weak Electrolytes – Those electrolytes which dissociate to a limited extent are called weak electrolytes. These electrolytes have a low degree of ionisation and lower dissociation constant value. They have low electrical conductivity. The dissociation is reversible; hence the law of mass action is applicable. Example: acetic acid, formic acid, weak bases like ammonium hydroxide and salts like ammonium acetate and silver acetate.
Characteristics of Electrolytic Dissociation
Dissociation is the process of separation of charged particles that already exist in a compound.
Dissociation involves ionic compounds.
Dissociation will either produce charged particles or electrically neutral particles.
Dissociation is reversible.
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 that are not ionic in nature. It involves covalent compounds. It is irreversible in nature. An appropriate solvent is required to start the process of ionisation and it is also called ion solvation.
Example: In the case of the 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)
Asymmetry Effect
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 a large number of either of the charges, the charge density increases at one end. This causes decreased conductance, however, symmetry is achieved after a short time. The effect is represented:
B=8.2x1050/(DT)3/2
D is dielectric constant, η is viscosity in poises, T is the absolute temperature
Electrophoretic Effect
A single ion is surrounded by solvent molecules and other ions, thus the ionic atmosphere of the central ion involves forces of both. The movement of the central ion in a direction opposite to that of the ionic atmosphere causes the withdrawal force by solvent molecules on the movement of the central ion. This new retarding force on the central ion due to friction between ion and solvent is known as an electrophoretic effect. This causes a 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 the 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 the degree of dissociation. It is denoted by ‘α’. It can be represented in the equation as:
The Degree of Dissociation and its Value is Found to be Dependent on the Following Factors
Nature of Solute: If the ionisable part of the molecule is bonded with covalent bonds, fewer ions are produced. And if the ionisable part of the molecule is held by an electro-covalent bond, more ions are produced.
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, the 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 the dissociation of charged particles.
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 a 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.
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 interparticle attraction force and resulting in more dissociation of ions.
Evidence in Support of Arrhenius Theory
X-ray diffraction studies show the presence of ions in electrolytes. It also shows that they conduct electricity infused states.
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:
Na+ (aq) + Ag+(aq) + NO3– (aq) + Cl–(aq) → AgCl(aq) + NaNO3(aq)
Based on Arrhenius theory, undissociated water is obtained in some systems, which leads to the change in the enthalpy of the system. The phenomenon is known as the enthalpy of neutralisation.
The colour of the electrolyte is due to the presence of an ion.
This theory forms the basis of solubility products, hydrolysis, common-ion effect, electrolysis, electrical conductivity, electrophoresis, etc.
The ionic theory can explain the abnormal and unpredictable colligative properties. When an electrolyte gets dissolved in water, the number of particles in the solution always increases, then a total number of molecules are dissolved due to ionisation.
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.
FAQs on Arrhenius Theory of Electrolyte Dissociation
1.Why is there a difference in the acidity of sodium chloride, hydrochloric acid and sodium hydroxide?
Sodium chloride is NaCl, which is an ionic compound of sodium and chlorine. When the crystals are dissolved in water it forms a solution of NaCl – NaCl solution that has no electrical conductivity. When it is put in a solution of hydrochloric acid or any other strong acid, the substance turns into ions and molecules lose their electrical properties. Similarly, when we dissolve a salt like sodium chloride in strong acid the substance turns into ions and becomes hydrated and loses its electrical properties. Hydrogen ionisation is more effective in hydrochloric acid than in any other acid (except sulfuric acid) because of its stronger effect on hydrophilic molecules.
2. Why is the atomic radius of fluorine (delta-F, δ) slightly larger than that of chlorine (delta-Cl, δ)?
Because of its orbitals and thus its valence electrons, fluorine is only slightly more reactive than chlorine. However, since it is significantly more electronegative than chlorine, fluorine has a smaller radius (F > Cl) because its valence electron repulsion is less effective in cancelling its more negative charges. Thus, fluorine atoms and molecules have a small number of electrons to be distributed amongst, resulting in fewer covalent bonds than in the case of chlorine. Fluorine's proton affinity (a measure of how easily a proton can be donated to a molecule) is only slightly less than chlorine's.
3. Why does the atomic weight of caesium differ from that of sodium?
Caesium and sodium have similar chemical properties; they are both in Group 1 (or alkali metals). Their bond energies are similar, and since the two elements have similar electronegativity, their molecular orbital energies should be also similar. These three factors indicate that their atoms should have similar atomic radii. However, when comparing the atomic radii of caesium and sodium, the differences are apparent. One of the ways in which covalent bonding is observed among the members of a group, such as those in Group 1, is through the nature of their ionic radii. The larger atoms (oxygen, nitrogen, fluorine, etc.) are usually the ones that form covalent bonds. Because of the electron affinities of the two elements, and their similar bond energies, caesium should, in theory, have a smaller ionic radius than sodium, but it has a greater atomic radius. This apparent contradiction, where larger atoms have smaller ionic radii than their lighter counterparts, is often resolved by noting that the charge of the atom (the electrostatic energy) increases as the covalent bond angle increases.
4. Does the nature of salt cause an electric charge to be built up in the molecule? Explain.
In order to answer this question, we must first look at the nature of the covalent bond between the two atoms, as well as what elements are bonded to them. There are two types of covalent bonds between elements. They are ionic and covalent. An ionic bond occurs when the electron from one atom tries to attach itself to another atom that has higher energy than the first atom. This can occur when the atom with high energy has a slightly smaller electron density than the atom with low energy.
5. Why does the behaviour of strong electrolytes differ from that of weak electrolytes?
When there is a large difference in charge on the atoms of two different elements, electrostatic forces will be drawn between the atoms. In order to explain the behaviour of strong electrolytes, we must first look at ionic and covalent bonds. Ionic bonds are an attractive force that attracts atoms with a difference in charge. Covalent bonds are relatively short-lived, and there is no tendency for the two atoms to hold together. the weak electrolytes will attract one another in order to maintain neutrality, while the strong electrolytes can withstand a large amount of charge separation due to the forces of ionic bonds.