THEORY OF ELECTROLYTIC CONDUCTANCE
Arrhenius theory of electrolytic conductance is also known as Arrhenius theory of ionisation since electrolytic dissociation into ions is considered here.
Postulates of Arrhenius Theory :
1. When dissolved in water, neutral electrolyte molecules are split up into two types of charged particles.
These particles were called ions and the process was termed ionisation. The positively charged particles were called cations and those having negative charge were called anions.
The theory assumes that the ions are already present in the solid electrolyte and these are held together by electrostatic force. When placed in water, these neutral molecules dissociate to form separate anions and cations.
A+ B- -- -- > A+ + B-
For that reason, this theory may be referred to as the theory of electrolytic dissociations.
2. The ions present in solution constantly reunite to form neutral molecules. Thus there is a state of equilibrium between the undissociated molecules and the ions.
AB < -- --- > A+ + B-
Applying the Law of Mass Action to the ionic equilibrium we have,
[ A + ][ B - ] / [AB] = K
where K is called the Dissociation constant.
3.The charged ions are free to move through the solution to the oppositely charged electrode. This is called as migration of ions. This movement of the ions constitutes the electric current through electrolytes. This explains the conductivity of electrolytes as well as the phenomenon of electrolysis.
4.The electrical conductivity of an electrolyte solution depends on the number of ions present in solution. Thus the degree of dissociation of an electrolyte determines whether it is a strong electrolyte or a weak electrolyte.
We know that electrolytes dissociate in solution to form positive ions (cations) and negative ions (anions).
AgNO3 -- -- - > Ag+ + NO3-
CuSO -- -- - > Cu2+ + SO 2-
H2SO-- -- - > 2H+ + SO 2-
As the current is passed between the electrode of the electrolytic cell, the ions migrate to the opposite electrodes. Thus in the electrolytic solution of AgNO3, the cations (Ag+) will move to the cathode and anions (NO3- ) will move to the anode. Usually different ions move with different rates. The migration of ions through the electrolytic solution can be demonstrated by the following experiments.
5.The properties of solution of electrolytes are the properties of ions. The solution of electrolyte as a whole is electrically neutral unless an electric
field is applied to the electrodes dipped into it. Presence of hydrogen ions (H+) renders the solution acidic while presence of hydroxide ions (OH- ) renders the solution basic.
6.There are two types of electrolytes. Strong electrolytes are those when dissolved in water are completely dissociated (ionised) into ions of positive and negative charges. The total number of cations and anions produced are equal to those in the formula of the electrolyte.
Al2(SO4)3 -- > 2Al3+ + 3SO 2-
NaCl, KCl, AgNO3 etc., are few examples of strong electrolytes.
In the case of weak electrolytes, there is partial dissociation into ions in water and an equilibrium exists between the dissociated ions and the undissociated electrolyte.
(e.g.,) CH3COOH < -- --- > CH3COO- + H+. Acetic acid is a weak
electrolyte in water and unionised acetic acid molecules are in equilibrium with the acetate anions and H+ ions in solution.
Evidences of Arrhenius theory of electrolytic dissociation
1. The enthalpy of neutralisation of strong acid by strong base is a constant value and is equal to -57.32 kJ. gm.equiv -1. This aspect is well explained by adopting Arrhenius theory of electrolytic dissociation. Strong acids and strong bases are completely ionised in water and produce H+ and OH- ions respectively along with the counter ions. The net reaction in the acid-base neutralisation is the formation of water from H+ and OH- ions.
H+ + OH- -- -- - > H2O, DHro = -57.32 kJ.mol -1
2.The colour of certain salts or their solution is due to the ions present. For example, copper sulphate is blue due to Cu2+ ions. Nickel salts are green due to Ni2+ ions. Metallic chromates are yellow due to CrO42- ions.
3.Ostwalds dilution law, common ion effect and solubility product and other such concepts are based on Arrhenius theory.
4.Chemical reactions between electrolytes are almost ionic reactions. This is because these are essentially the reaction between oppositely charged ions. For example,
Ag+ + Cl- --- -- > AgCl ¯
5.Electrolytic solutions conduct current due to the presence of ions which migrate in the presence of electric field.
6.Colligative properties depend on the number of particles present in the solution. Electrolytic solution has abnormal colligative properties. For example, 0.1 molal solution of NaCl has elevation of boiling point about twice that of 0.1 molal solution of non-electrolyte. The abnormal colligative properties of electrolytic solutions can be explained with theory of electrolytic dissociation.
Ostwald's dilution law for weak electrolytes
According to Arrhenius theory, weak electrolytes partially dissociate into ions in water which are in equilibrium with the undissociated electrolyte molecules. Ostwald's dilution law relates the dissociation constant of the weak electrolyte with the degree of dissociation and the concentration of the weak electrolyte. Consider the dissociation equilibrium of CH3COOH which is a weak electrolyte in water.
CH3COOH <-- -- > CH3COO- + H+
Ka = [ H + ][CH 3COO - ] / [CH3COOH]
a is the degree of dissociation which represents the fraction of total concentration of CH3COOH that exists in the completely ionised state. Hence (1 - a) is the fraction of the total concentration of CH3COOH, that exists in the unionised state. If 'C' is the total concentration of CH 3COOH initially, then at equilibrium Ca, Ca and C (1 - a) represent the concentration of H+, CH3COO- and CH3COOH respectively.
Then Ka = (Ca .C a) / C (1-a ) / a2 C / (1-a)
If a is too small, then Ka = a2C
a = root(Ka/C) also [H+] = [CH3COO-] = Ca
[H+] = root(Ka.C)
Ka= a2C / (1-a) is known as the Ostwalds dilution law. For weak bases,
Kb= a2C / (1-a) and a = rt(Kb/C) at a = small values.
Kb = dissociation constant for weak base.
This law fails for strong electrolytes. For strong electrolytes, a tends to 1.0 and therefore Ka increases tremendously.
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