Electrochemistry Class 12 Notes

Electrochemistry

  MCQs  

Glvanic Cell or Voltaic Cell

An electrochemical cell that converts the chemical energy of spontaneous redox reactions into electrical energy is known as a galvanic cell or a voltaic cell.


Important Points of the Galvanic Cell
Oxidation occurs at Anode (−ve Electrode)
Reduction occurs at Cathode (+ve Electrode)
Current flows from Cathode to Anode
Electrone flows from Anode to Cathode
Salt bridge(Contains electrolytes) complete the circuit
Electromotive Force (emf) of the Cell is 1.1 Volt
Oxidation Half Cell Reaction- (at Anode)
Zn(s) → Zn⁺² + 2e
Reduction Half Cell Reaction- (at Cathode)
Cu⁺² + 2e → Cu(s)
Complete Cell Reaction-
Zn(s) + Cu⁺² → Zn⁺² + Cu(s)
Cell Representation-
Zn(s)|ZnSO₄(aq) || CuSO₄|Cu(s)

Electrmotive Force (emf)

emf stands for electromotive force. A galvanic cell consists of two electrodes each having its own potential. The difference of potential between these two electrodes of a cell causes a current to flow from electrode of higher potential to the electrode of lower potential is called emf of the cell.
emf is expressed in volts. Greater the emf (i.e. potential difference between two electrode) greater the electricity flow and greater is the tendency of the cell redox to occur.
Potential of a cell assembled of two electrodes can be determined from the two individual electrode potentials using this formula-
ΔV_cell = E_red,cathode − E_red,anode
or, ΔV_cell = E_red,cathode + E_oxy,anode

Electrode Potential or Half Cell Potential

Electrode potential is the electromotive force of a galvanic cell built from a standard reference electrode and another electrode to be characterized. The electrode potential has its origin in the potential difference developed at the interface between the electrode and the electrolyte.
When a metal(M) is dipped in a solution containing its own ion(M⁺), a potential is developed between them.
This is called Single Electrode Potential(E). It is not possible to measure accurately the absolute value of single electrode potential directly. Only the difference in potential between two electrodes can be measured experimentally with the help of reference electrode whose potential is already known. Cell potential is measured by the following formula-
E_Cell = E_Cathode + E_Anode
emf of single electrode potential is-
E_M/M⁺ⁿ = E^o_M/M⁺ⁿ + (RT/nF)ln a
When a = 1 then E = E^o. So, Standard electrode potential is the single electrode potential at unit activity.
M ⇌ M⁺ⁿ + ne single elecrode potential for the above equilibrium-
E_M/M⁺ⁿ = E^o_M/M⁺ⁿ + (RT/nF)ln [M]/[M⁺ⁿ]
or, E_M/M⁺ⁿ = E^o_M/M⁺ⁿ + (RT/nF)ln 1/[M⁺ⁿ] (as [M] = 1)
or, E_M/M⁺ⁿ = E^o_M/M⁺ⁿ − (RT/nF)ln [M⁺ⁿ]
or, E_M/M⁺ⁿ = E^o_M/M⁺ⁿ − (0.0591/n)log [M⁺ⁿ]
This is the general expression for electrode potential.

Oxidation Potential and Reduction Potential
The tendency of an electrode to lose electrons is called oxidation potential while tendency of electrode to gain electrons is called reduction potential.

Standard Electrode Potential

The potential of a half-reaction measured against the Standard Hydrogen Electrode under standard conditions is called the standard electrode potential for that half-reaction.
Standard conditions are-
Temperature = 298K
Pessure = 1atm
Concentration of the electrolyte = 1M.
Example-
Consider Zn and Hydrogen electrode
Zn ⇌ Zn²⁺ + 2e
H₂ ⇌ 2H⁺ + 2e
When these two electrodes (two equilibria) are brought into electrical contact using an external wire and a salt bridge, the electrons will be pushed from the zinc equilibrium (electrode) to the hydrogen equilbrium (electrode) with a force of - 0.76V (the negative sign simply indicates the direction of flow - from zinc to hydrogen ions).
So, the standard electrode potential of Zn is −0.76 volts and the overall reaction is-
Zn + 2H⁺ → Zn²⁺ + H₂

Uses of Standard Electrode Potentials

Uses of standard electrode potentials are given below –
1. It is used to measure relative strengths of various oxidants and reductants.
2. It is used to calculate standard cell potential.
3. It is used to predict possible reactions.
4. Prediction of equilibrium in the reaction.

Standard Hydrogen Electrode (SHE)

Standard Hydrogen Electrode is a reference electrode consists of a container, containing solution kept at 298K.
A wire containing Platinum electrode coated with platinum black is immersed in the solution.
Pure hydrogen gas is bubbled in the solution at 1bar pressure.
The potential of standard hydrogen electrode is taken as zero volt at all temperatures.
Standard hydrogen electrode may act as anode or cathode depending upon the nature of the other electrode.
If its acts as anode, the oxidation reaction taking place is
H₂(gas) → 2H⁺(aq) + 2e
If it acts as cathode then the reduction half reaction occurring is
2H⁺(aq) + 2e → H₂(gas)

Nernst's Equation or Relation between emf and Activity of a Cell

A galvanic cell is the cell which generates electricity from energy change accompanying a chemical reaction. The emf of such cell arises due to the reaction occuring in the cell. If the following reaction occurs in a reversible chemical cell-

Relation Between Free Energy Change and Equilibrium Constant

Relation Between emf and Free Energy Change of a Cell

A reactant possesses a specific internal energy which is called free energy.
Change in free energy-
−ΔG = G_final − G_initial
As the result of cell reaction occuring in a galvanic cell, the change in free energy is equal to the product of emf(E) of the cell and the quantity of electricity (nF) that has passed through the cell.
i.e. −ΔG = nEF
where- 'n' is the valency of the element
'F' is 1 Faraday = 96500 Coulombs.
Thus, if the passage of the 'n' Faraday of electricity liberates one gram atom of the element.
i.e. −ΔG = nEF
Thus,
1. If ΔG is negative, FE is positive or vice versa and cell reaction will occur spontaneously.
2. If ΔG is positive, FE is negative and cell reaction will not possible.
3. If ΔG is zero, and E = 0, the cell reaction will be in equilibrium and hence, the reaction will be reversible.

Electrochemical Series

Electrochemical series describes the arrangement of elements in order of their increasing electrode potential values. By measuring the potentials of various electrodes versus standard hydrogen electrode (SHE), a series of standard electrode potentials has been established.
Electrodes with positive E° values for reduction half reaction act as cathodes versus SHE, while those with negative E° values of reduction half reactions behave as anodes versus SHE.
The negative sign of standard reduction potential indicates that an electrode when joined with SHE acts as anode and oxidation occurs on this electrode. For example, standard reduction potential of zinc is -0.76 volt. When zinc electrode is joined with SHE, it acts as anode (-ve electrode) i.e., oxidation occurs on this electrode.
Similarly, the +ve sign of standard reduction potential indicates that the electrode when joined with SHE acts as cathode and reduction occurs on this electrode.
On moving down the series-
Reduction Potential Decreases
Oxidation Poitential Increases
Strength of Oxidizing Agent Decreases
Strength of Reducing Agent Increases
Reactivity of Metals Increases
Reactivity of Nonmetals Decreases

Conductance (π)

Any substance that allows an electric current to pass through is called conductor and the capacity to conduct electricity is called the conductivity or conductance of a conductor. Conductor offers to resistance to flow the electric current.
So, the conductance is resiprocal to resistance (R)
π = 1/R
Unit of π = 1/ohm = ohm⁻¹ = mhos
SI Unit is Siemen (S)

Specific or Electrolytic Conductance (κ)

The resiprocal of specific resistance (ρ) is called specific conductance.
We know that-
R ∝ l/a
where, l is length of the wire and a is cross sectional area of the wire
or, R = ρ. l/a
or, 1/ρ = 1/R . l/a
or, κ = π . l/a
Unit of κ = ohm⁻¹ . cm/cm²
ohm⁻¹ . cm⁻¹
Its S.I. unit is Sm⁻¹
when, l = 1 and a = 1
then, κ = π
So, specific conductance is the conductance of a conductor of unit length and unit cross sectional area. For electrolytic solutions, the conductance of one cc of the solution is called its specific conductance.

Equivalent Conductance (Λ)

It is denoted by capital lambda(Λ) and is conductance of V cc of the solution containing one gm-equivalent of an electrolyte. So, it is the product of Specific or Electrolytic Conductance (κ) and the volume of the solution (V) in cc containing one-gm equivalent of the electrolyte.
or, Λ = κ . V
If the concentration of a solution is C g-equivalent/liter, then-
C g-equivalent is present in 1000cc of the solution.
so, 1 g-equivalent is present in 1000cc/C of the solution.
or, Volume in cc containing one gm-equivalent = 1000/C
so, Λ = 1000.κ/C
Unit of Λ = κ/C = ohm⁻¹ . cm⁻¹/eq-cm⁻³
or, ohm⁻¹ . cm² eq⁻¹

Molecular Conductance (μ)

It is denoted by capital mu(μ) and is conductance of V cc of the solution containing one mole of an electrolyte. So, it is the product of Specific or Electrolytic Conductance (κ) and the volume of the solution (V) in cc containing one mole of the electrolyte.
or, μ = κ . V
If the concentration of a solution is C mole/liter, then-
C mole is present in 1000cc of the solution.
so, 1 mole is present in 1000cc/C of the solution.
or, Volume in cc containing one mole = 1000/C
so, μ = 1000.κ/C
When concentration approaches zero, the molar conductivity is known as limiting molar conductivity and is represented by E^o_m.
Unit of μ = κ/C = ohm⁻¹ . cm⁻²mol⁻¹
or, Sm² . cm²mol⁻¹

Relationship between molar conductivity and Concentration

For strong electrolyte, molar conductivity increases with dilution and can be represented by the equation-
μ = E^o_m − AC^1/2
where A is a constant depending upon the type of the electrolyte, the nature of the solvent and the temperature.
The equation is called Debye Huckel-Onsager equation and is found to hold good at low concentrations.
If we plot μ against c^1/2, we obtain a straight line with intercept equal to E^o_m and slope equal to '–A'.

Cell Constant

We know that the resistance (R) is given by-
R ∝ l/a
or, R = ρ. l/a
or, 1/ρ = 1/R . l/a
or, κ = π . l/a
where, ρ is specific conductance, l is length of wire, a is cross sectional area of the wire κ is specific conductance and π is conductance of conductor.
The ratio of l to a is called cell constant.
Hence, cell constant = κ / π = κ . R
Unit of cell constant:
l/a = cm/cm² = cm⁻¹

Effect of dilution on conductivity

The specific conductance depends on the number of ions present per cc of the solution. Though degree of dissociation increases on dilution but the number of ions per cc decreases. So, the specific conductance decreases on dilution.
The equibvalent conductance is the product of specific conductance and volume of the solution containing one gm-equivalent of the electrolyte.
Λ = κ . V
As the decreasing κ value is more than compensated by the increasing V value, hence, Λ increases. on dilution.

Kohlrausch's Law

The equivalent conductivity of an electrolyte at infinite dilution (Λ_o) is the sum of the ionic conductivities of their cations and anions.
Λ_o = λ₊ + λ₋
where- λ₊ and λ₋ are cationic and anionic conductivities at infinite dilution respectively.
Applications:
Useful in
calculating equivalent conductivity at infinite dilution.
Calculation of degree of dissociation of an electrolyte.
Calculation of solubility of sparingly soluble salt,
Calculation of ionic product of water.

Q. Calculate the limiting molar conductivity of CH₃COOH. The molar conductivities of CH₃COONa, HCl and NaCl at infinite dilution are 90.1 S.cm².mol^-1, 426.16 S.cm².mol^-1 and 126.45 S.cm².mol^-1 respectively.

Given-
λ_CH3COONa = 90.1 S.cm².mol^-1
λ_HCl=426.16 S.cm².mol^-1
λ_NaCl=126.45 S.cm².mol^-1
According to Kohlrausch law-
λ_CH3COOH = λ_CH3COONa + λ_HCl – λ_NaCl
or, λ_CH3COOH = 90.1 + 426.16 – 126.45
or, λ_CH3COOH = 390.71 S.cm².mol^-1
So, the limiting molar conductivity of CH₃COOH is, 390.71S.cm².mol^-1.

Q. From the given molar conductivities at infinite dilution, calculate λ_m for NH₄OH.
λ_m for Ba(OH)₂ = 457.6 ohm^-1 cm²mol^-1
λ_m for Ba(Cl)₂ = 240.6 ohm^-1 cm²mol^-1
λ_m for NH₄Cl = 129.8 ohm^-1 cm²mol^-1

Answer: 238.3 ohm^-1 cm²mol^-1

Electrolytes

Compounds which either in solution or in molten state conduct electricity are called electrolytes.
Acids, Bases and Salts are the examples of electrolytes.
HCl ⇌ H⁺ + Cl⁻
CH₃COOH ⇌ CH₃COO⁻ + H⁺
NaOH ⇌ Na⁺ + OH⁻
CuSO₄ ⇌ Cu⁺² + SO₄⁻²
Electricity is carried out by the mobile ions present in them.
Compounds which don't ionize in solution consequently can not conduct electricity are called nonelectrolytes.
Sugar, Coal, Wood etc are non electrolytes.
Electrolytes are of two types-
1. Strong electrolyte
2. Weak electrolyte
1. Strong electrolyte
The electrolytes which are completely ionised are called strong electrolytes. They have high electrical conductivity and does not applicable to Ostwald’s dilution law.
Example-
HCl ⇌ H⁺ + Cl⁻
2. Weak electrolyte
The electrolytes which are partially ionised called strong electrolytes. hey have low electrical conductivity and applicable to Ostwald’s dilution law.
Example-
CH₃COOH ⇌ CH₃COO⁻ + H⁺

Electrolytic Cell

Electrolytic cell is the device by which electrolysis is carried out. Electrolyte in molten state or in dissolved state is taken in a glass vessel and two metallic strips(electrodes) are dipped in the electrolyte. These strips are connected to a battery. One metal strip acts as cathode (-ve electrode) while the other metal strip acts as anode (+ve electrode). On passing current, reduction occurs at cathode and oxidation at anode.

Applications of Electrolytic Cells

Followings are some imortant applications of Electrolytic Cells-
1. The primary application of electrolytic cells is for the production of oxygen gas and hydrogen gas from water.
2. They are also used for the extraction of aluminium from bauxite.
3. Electrolytic cells is also use in electroplating, which is the process of forming a thin protective layer of a specific metal on the surface of another metal.
4. The electrorefining of many non-ferrous metals is done with the help of electrolytic cells. 5. To remove the impurities from the metal, the electrorefining method is followed.

Electrolysis

If an electric current is passed through an electrolyte taken either in solution or in fused state chemical change occurs. This phenomenon is called electrolysis.
Thus in electrolysis, electrical energy converted into chemical energy.
Electrolysis takes place in electrolytic cell in which two electrodes are dipped in a electrolytic solution. Chemical reaction occurs by passing current through electrodes. Oxidation and reduction takes place at the anode and cathode respectively as takes palce on galvanic cell.

Electrolysis of Water

Cathode-
H₂O + e → 1/2H₂ + OH⁻
Anode-
H₂O → 1/2O₂ + 2H⁺ + 2e

Electrolysis of Molten NaCl

NaCl ---Heat---> Na⁺ + Cl⁻
Cathode-
Na⁺ + e → Na
Anode-
Cl⁻ → 1/2Cl₂ + e

Electrolysis of Brine Solution of NaCl

Cathode-
Na⁺ + e → Na
H₂O + e → 1/2H₂ + OH⁻
At Cathode Na is diposited as its E^o is more negative.
Anode-
Cl⁻ → 1/2Cl₂ + e
H₂O → 1/2O₂ + 2H⁺ + 2e
At Anode chlorine is diposited as the amount of oxygen is less.

Electrolysis of aquous solution of AgNO₃ with Silver electrode

Reaction in solution-
AgNO₃ ⇌ Ag⁺ + NO⁻
H₂O ⇌ H⁺ + OH⁻
Cathode-
Ag⁺ + e → Ag
Anode-
Ag + NO₃⁻ → AgNO₃ + e

Electrolysis of aquous solution of AgNO₃ with Pt- electrode

Reaction in solution-
AgNO₃ ⇌ Ag⁺ + NO⁻
H₂O ⇌ H⁺ + OH⁻
Cathode-
Ag⁺ + e → Ag
Anode-
Due to Pt-electrode, self ionization of water will take place.
H₂O → 2H⁺ + 1/2O₂ + 2e
Ag + NO₃⁻ → AgNO₃ + e

Electrolysis of dilute solution of H₂SO₄ with Pt- electrode

Reaction in solution-
H₂SO₄ ⇌ 2H⁺ + SO₄⁻²
H₂O ⇌ H⁺ + OH⁻
Cathode-
H⁺ + e → 1/2H₂
Anode-
Due to Pt-electrode, self ionization of water will take place.
H₂O → 2H⁺ + 1/2O₂ + 2e

Electrolysis of aqueous solution of CuCl₂ with Pt- electrode

Reaction in solution-
CuCl₂ ⇌ 2H⁺ + 2Cl⁻
H₂O ⇌ H⁺ + OH⁻
Cathode-
Cu⁺² + 2e → Cu
Anode-
2Cl⁻ → Cl₂ + 2e

Electrolysis CuSO₄

Cathode-
Cu⁺² + 2e → Cu
Anode-
H₂O → 1/2O₂ + 2H⁺ + 2e

Faraday's Law of Electrolysis