Electrochemistry is the branch of chemistry which deals with the chemical changes caused in the matter by passage of electric current and conversion of chemical energy into electrical energy and vice versa.
Conductors and Non Conductors
Substances around us can be divided into two classes based on their ability of conduct electricity:
Comparison of Electrolytic and Metallic Conduction
Electric current flows by movement of electrons.
Electric current flows by movement of ions.
No chemical change occurs.
Ions are oxidized or reduced at the electrodes.
It does not involve the transfer of any matter.
It involves transfer of matter in the form of ions.
Ohm's law is followed.
Ohm's low is followed.
Resistance increases with increase of temperature.
Resistance decreases with increase of temperature.
Faraday law is not followed.
6 Faraday law is followed.
Specific Resistance or resistivity (ρ):
R = ρ× l /A
R = Resistance
A=Area of cross sections of electrodes
l = Distance between the electrodes.
Specific Conductance or Conductivity (κ):
κ = 1/ ρ
Units: W–1 cm–1 or Sm–1
Equivalent Conductance (Λ):
Conducting power of all the ions produced by one g-equivalent i.e. one equivalent of an electrolyte in a given solution
Λ = 1000 κ /C
Where, C be the normality of solution i.e. concentration of electrolytic solution in equivalent/L.
Molar Conductance (Λm):
Conductance of solution due to all the ions produced by one mole of the dissolved electrolyte in a given solution.
Λm = k/c
Where, c = concentration of solution in mol m-3.
Λm = k×1000/M
Where, M is molarity of solution.
Units: S cm2 mol-1
Relation between Λ and Λm
Λm = n Λ
Where n = n-factor of the electrolyte = total charge carried by either ion = M/E
Variation of Conductance with Dilution:
Specific conductance: Decrease with dilution due to decrease in number of ions per c.c. of the solution.
Molar and Equivalent Conductance: Increases with increase in dilution.
Λm = Λ0m – (A+B Λ0m) √C
A & B = Debye-Hückel constants.
C = Molar concentration
Λ0m = Limiting molar conductivity i.e. molar conductivity at infinite dilution.
At 298 K,
Λm = Λ0m – (60.2 – 0.299 Λ0m) √C
Kohlransch’s Law of Independent Migration of Ions:
"At infinite dilution, when dissociation is complete, each ion makes a definite contribution towards equivalent conductance of the electrolyte irrespective of the nature of the ion with which it Is associated and the value of equivalent conductance at infinite dilution for any electrolyte is the sum of contribution of its constituent ions", i.e., anions and cations. Thus,
According to Kohlrausch’slaw. “conductivity of ions is constant at infinite dilution and it does not depend on nature of co-ions.
”For AxBy type electrolyte,
Here Z+and Z- are the charges present on cation and anion.
Here m and n are the number of moles of cations and anions.
Faraday’s Laws of Electrolysis:
First Law: “The mass of a substance deposited or liberated at any electrode is directly proportional to the amount of charge passed.”
w = zQ
Where z = electrochemical equivalent i.e. the mass of the substance in grams deposited or liberated by passing one coulomb of charge,
Second Law: “Mass of a substance deposited or liberated at any electrode on passing a certain amount of charge is directly proportional to its chemical equivalent weight”.
Charge on one mole electrons = 1 F = 96487 C
That is w a E where w is the mass of the substance in grams while E is its chemical equivalent weight in gms per equivalent=
The charge possessed by 1 mole of electrons
This charge is called as 1 Faraday.
If we pass one Faraday of charge, it means that we are passing one mole of electron and by passing 1 Faraday of charge 1gm equivalent weight of the substance will be deposited or liberated.
By combining the first and second law, we get
Difference in Electrolytic Cell and Galvanic Cell:
Electrical energy is converted into chemical energy.
Chemical energy is converted into electrical energy.
Anode positive electrode. Cathode negative electrode
Anode negative electrode. Cathode positive electrode.
Ions are discharged on both the electrodes.
Ions are discharged only on the cathode
If the electrodes are inert, concentration of the electrolyte decreases when the electric current is circulated
Concentration of the anodic half-cell increases while that of cathodic half-cell decreases when the two electrodes are joined by a wire
Both the electrodes can be fitted in the same compartment
The electrodes are fitted in different compartment
Standard electrode potential: The potential difference developed between metal electrode and the solution of its ions of unit molarity (1M) at 25°C (298 K)
IUPAC Cell Representation: Anode (Molarity of electrolyte at anode) || Cathode (Molarity of electrolyte at cathode)
E0cell = E0cathode - E0anode
The Nernst Equation:
For a general reaction such as
M1A + m2B ..... n1X + n2Y + .... .......(i)
Occurring in the cell, the Gibbs free energy change is given by the equation
'a' represents the activities of reactants and products under a given set of conditions and
?Go refers to free energy change for the reaction when the various reactants and products are present at standard conditions.
The free energy change of a cell reaction is related to the electrical work that can be obtained from the cell, i.e.,
?Go = -nFEcell and ?Go = -nFEo.
On substituting these values in Eq. (ii) we get
This equation is known as Nearnst equation.
Some other important relations:
Some Important Half Cells:
Ferrous – Ferric Electrode:
Quin – Hydrone Electrode:
Oxidation and Reduction
Oxidation Number / State Method For Balancing Redox Reactions:
This method is based on the principle that the number of electrons lost in oxidation must be equal to the number of electrons gained in reduction. The steps to be followed are :
Half-Reaction or Ion-Electron Method For Balancing Redox Reactions
This method involves the following steps :
Common Oxidising and Reducing Agents
Decrease in Oxidation Number
KMnO4 in acid solution
MnO4 - → Mn2+
KMnO4 in alkaline solution
MnO4 - → MnO2
K2Cr2O7 in acid solution
Cr2O72- → Cr3+
NO3- → NO
NO3- → NO2
SO42- → SO2
manganese (IV) oxide
MnO2 → Mn2+
Cl → Cl-
chloric (I) acid
ClO- → Cl-
KlO3 in dilute acid
IO3- → I
KlO3 in concentrated acid
IO3- → I-
Increase in Oxidation Number
iron (II) salts (acid)
Fe2+ → Fe3+
tin (II) salts (acid)`
Sn2+ → Sn4+
C2O42- → CO2
S2- → S
iodides (dilute acid)
I- → I
iodides (concentrated acid)
I- → I+
metals, e.g. Zn
Zn → Zn2+