Arrhenius Theory of Ionization and Dissociation

Arrhenius Theory of Electrolytic Dissociation: 

Electrolyte:  

  • According to Arrhenius theory, a substance (acid, base or salt ) which when dissolved in water splits up spontaneously into positively and negatively charged ions and aqueous solution has electrical conductivity is called as an electrolyte.
  • Examples:

NaCl(aq)     →    Na+ (aq)   +    Cl(aq) 

H2SO4 (aq)     →    2 H+ (aq)   +    SO42-(aq) 

  • In modern theory, it is assumed that the solid electrolytes consist of two types of charged particles, one carrying a positive charge and other carrying a negative charge. They are held together by the electrostatic force of attraction. When such solid electrolytes are dissolved in a solvent, these forces weakened and electrolyte undergoes dissociation into ions. The process is also called ion solvation.
  • non -electrolyte is a substance which in its aqueous solution or in fused state does not conduct electricity (due to no formation of ions). Examples: sugar, urea, ethanol, starch, acetone. etc.

Strong electrolytes: 

  • Substances which dissociate almost completely in their aqueous solutions even at moderate dilutions are called as strong electrolytes.
  • Their dissociation reaction is irreversible.
  • Examples:
    • Strong acids like HCl, HNO3 H2SO4  etc.
    • Strong bases like NaOH, KOH etc.
    • Salts like NaCl, KCl etc.
  • Characteristics of Strong Electrolytes:

    • Substances which dissociate almost completely in their aqueous solutions even at moderate dilutions are called as strong electrolytes.
    • The degree of dissociation is high.
    • Law of mass action is not applicable since dissociation is irreversible.
    • Their solution has high conductivity.
    • For strong electrolyte dissociation constant has a higher value.

Weak electrolytes :

  • Substances which dissociate to a little (limited) extent in their aqueous solutions are called weak electrolytes.
  • Examples:
    • weak acids like CH3COOH, HCOOH etc.
    • Weak bases like NH4OH etc.
    • salts like CH3COONH4, CH3COOAg etc.
    • Characteristics of Weak Electrolytes:

    • Substances which dissociate to a little (limited) extent in their aqueous solutions are called weak electrolytes.
    • The degree of dissociation is low.
    • Law of mass action is applicable since dissociation is reversible.
    • Their solution has low conductivity.
    • For weak electrolyte dissociation constant has the lower value.

Ionisation:

  • It is the formation of the ions from molecules which are not initially in the ionic state.
  • Example: In HCl molecule, H and Cl atoms are covalently bonded. But when dissolved in water forms H+  and Cl ions.

HCl(aq)     →    H+ (aq)   +    Cl(aq) 

Dissociation:

  • The spontaneous splitting of a substance into positively and negatively charged ions in an aqueous solution is called as dissociation.
  • Example: In NaCl molecule, Na and Cl atoms are bonded with the ionic bond. They exist in the ionic state even after formation of the compound.

NaCl(aq)     →    Na+ (aq)   +    Cl(aq) 



Degree of dissociation (α):

  • The fraction of the total number of moles of weak electrolyte that ionises into ions in an aqueous solution at equilibrium is called as the degree of dissociation. It is denoted by ‘α’
  • Degree of dissociation = no. of molecules dissociated as ions / total number of molecules present
  • Percentage dissociation or ionisation = Degree of dissociation × 100
  • The degree of dissociation or ionisation depends on following factors.
    • The nature of solute:
      • When the ionizable parts of a molecule of a substance are held more by covalent bond than by electrovalent bond, fewer ions are furnished in the solution. e.g. H2S, HCN, CH3COOH, NH4OH etc.
      • When the ionizable parts of a molecule of a substance are held mainly by electrovalent bond, more ions are furnished in the solution. e.g. NaCl, KOH etc.
    • The nature of solvent:
      • The main function of solvent is to weaken the electrostatic force of attraction between the ions.  By Coulomb’s law, the magnitude of the force between two charged particles is inversely proportional to the dielectric constant of the medium between the charged particles.
      • The solvent having more dielectric constant has the higher capacity of separating the ions. Water (85) > Methyl alcohol (35) > Ethyl alcohol (27) > Acetone (21). Thus water is a good solvent.
    • 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 if the dilution increases (concentration decreases) the degree of ionisation increases.
    • Temperature:
      • Due to increase in temperature, the kinetic energy of the molecules increases and thus attractive force between the ions in the molecule decreases, resulting in easier ionisation (dissociation). Thus if temperature increases the degree of ionisation increases.


Evidences in Favour of Arrhenius Theory :

  • X-ray diffraction studies have shown that electrolytes are composed of ions. For example, NaCl is present as Na+Cl. Each Na+ ion is surrounded by six Cl ions. In turn each Cl ion is surrounded by six Na+ ions. A total number of Na+  ions is equal to total number of Cl ions. It conducts electricity in the fused state.
  • Electrolytic solutions obey Ohm’s law. This is only possible if ions are already present in the solution.
  • Following ionisation, reaction is possible due to existence of ions

Ag+(aq) + NO3 (aq)  + Na+ (aq) + Cl(aq) →  AgCl(aq)  + NaNO3(aq)

  • A similar reaction of AgNO3 with CCl4, CH3Cl, CH2Cl2, CHCl3 is not possible as these substances are not ionic compounds.
  • By Arrhenius, theory neutralisation is the reaction in which H+ ion from acid and OH ion from base react together to give practically un-dissociated water. Due to which there is a change in enthalpy of the system. This change in enthalpy is known as enthalpy of neutralisation.
  • Abnormal behaviour of electrolytes towards colligative properties can be explained on the basis of ionic theory only. When an electrolyte is dissolved in water, the number of particles in the solution is always more than the number of molecules actually dissolved due ionisation. The van’t Hoff factor is defined as

i = Observed colligative property / Calculated colligative property

Value of i is always more than unity. i.e  i = 1 + (n – 1)α

Where n is the number of ions produced from one molecule of electrolyte and α is



and α is the degree of dissociation.

  • Colour of electrolytic solutions is due to the presence of ions.
  • Ionic theory successfully explains the concept of common ion effect, solubility product, hydrolysis, electrolysis, the conductivity of electrolytic solutions etc.

Limitations of Arrhenius  Ionic Theory :

  • Arrhenius theory defines electrolyte in terms of their aqueous solution and not in terms of the substances themselves. Hence this theory is applicable to aqueous solutions only and not applicable to non-aqueous and gaseous reactions.
  • The theory does not consider the role of solvent in the deciding the nature of strength of an electrolyte. For e.g. HCl is strong acid when dissolved in water but it is weak acid when dissolved in benzene.
  • Ostwald’s dilution law which is based on Arrhenius theory is applicable to weak electrolytes only.
  • Strong electrolyte conduct electricity in a fused state (in absence of water). This is a contradiction to Arrhenius theory.
  • Arrhenius theory fails to explain the factors those affect degree of dissociation.

Preferential Discharge Theory:

  • If an electrolytic solution contains more than two ions and electrolysis is done, it is observed that all the ions are not discharged at the electrode simultaneously but certain ions are liberated at electrodes in preference to other. This phenomenon can be explained on the basis of preferential discharge theory.
  • It states that if more than one type of ions are attracted towards particular electrode, then the one discharged is the ion which requires the least energy.
  • The potential at which the ion is discharged or deposited on the appropriate electrode is called discharge or deposition potential. Discharge potential is different for different ions.

Example:

  • In case of NaCl in water, there are two equilibria Thus there are four ions involved.

NaCl(aq)    →  Na+ (aq) + Cl(aq)

H2O  →  H+ (aq) + OH(aq)

  • Now discharge potential of H+ is lower than that of Na+. Hence at cathode H+ ions will get discharged preferably. Similarly, discharge potential of Cl  ion is lower than OH ions. Hence at anode, Cl  ions will get discharged preferably.
  • Thus Na+ and OH ions remain in solution and when the solution is evaporated crystals of sodium hydroxide (NaOH) are obtained.
  • The decreasing order of discharge potential for cations is K+ > Na+ > Ca+2  > Mg+2 >  Zn+2 > H+ > Cu+2  > Hg +2 > Ag+ . The decreasing order of discharge potential for anions is SO4-2 > NO3  > OH >  Cl > Br > I
  • Note: When Hg is used as cathode, Na+ ions have lower discharge potential than H+ ions.

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