Revision: Electrochemistry CUET (UG) Electrochemistry

Electrochemistry is the study of the production of electricity from energy which is released during spontaneous chemical reactions, as well as the use of electrical energy to bring about non-spontaneous chemical transformations.

The electrode at which the reduction occur is called cathode.

Fuel cells are the galvanic cells in which the energy of combustion of fuels like hydrogen, methanol, etc., is directly converted into electrical energy.

The electrode at which the oxidation occur is called anode.

Cell constant is the ratio of the distance between the electrodes divided by the area of cross-section of the electrode. It is denoted by b.
Thus, Cell constant = b =`l/a`. It is expressed in unit m⁻¹.

Electrochemistry is the branch of chemistry that deals with the production of electricity from the energy released during spontaneous reactions and the use of electrical energy to drive non-spontaneous reactions.

It is the inverse of resistance R and may be simply defined as the speed through which current flows in a conductor.

c = `1/R = A/(pl)`

k = `A/l`

Here k is the specific conductance. The SI unit of conductance is Siemens, which is denoted by the symbol ‘S’ and is equal to ohm⁻¹ or Ω⁻¹.

The limiting molar conductivity of an electrolyte is defined as its molar conductivity when the concentration of the electrolyte in the solution approaches zero.

When the concentration of an electrolytic solution placed between electrodes of a conductivity cell placed at a unit distance having an area of cross-section sufficient to accommodate enough volume of solution containing one mole of electrolyte approaches zero, then the conductance of the solution is known as limiting molar conductivity.

Molar conductivity is the conductance of a volume of solution containing 1 mole of dissolved electrolyte when placed between two parallel electrodes 1 cm apart and large enough to contain between them all the solution.

The conductivity, which is shown by all the ions when 1 mol of electrolyte is dissolved in the solution, is called molar conductivity; it is expressed by ∧_m (lambda). If 1 mol of electrolyte is present in V_m cm³ of electrolyte solution, then ∧_m = κ × V

= `(kappa xx 1000)/"Molarity" = (kappa xx 1000)/M`

Its unit is ohm⁻¹ cm² mol⁻¹ or S cm² mol⁻¹.

A fuel cell is a galvanic cell in which the reactants are not placed within the cell, but are continuously supplied from outside, where one reactant acts as a fuel (such as hydrogen or methanol) and the other as an oxidant (such as oxygen).

Corrosion is the gradual damage of metals caused by their reaction with components of the atmosphere, such as oxygen and moisture.

Kohlrausch law states that at infinite dilution of the solution, each ion of electrolyte migrates independently of its co-ions and contribute independently to the total molar conductivity irrespective of the nature of other ion.

Kohlrausch’s law states that the molar conductivity of an electrolyte at infinite dilution is the same as the sum of the anions' and cations' limited molar conductivities.

`∧_m^° = v_+  λ_+^° + v_-  λ_-^°`

Here `λ_+^°` and `λ_-^°` are limiting molar conductivities of cations and anions.

Electrolysis of NaCl

1. Molten NaCl:

• Oxidation: Cl⁻ → Cl₂ (gas)

• Reduction: Na⁺ → Na (metal)

• Products: Na (cathode), Cl₂ (anode)

2. Aqueous NaCl:

• Oxidation: Cl⁻ → Cl₂

• Reduction: H₂O → H₂ + OH⁻

• Products: H₂ (cathode), Cl₂ (anode), NaOH formed

Components of a Galvanic Cell

6. Cell Notation 

• Anode written on the left, cathode on the right

• Example:

  Cu(s) | Cu²⁺(aq) || Ag⁺(aq) | Ag(s)

• Single line (|) → phase boundary

• Double line (||) → salt bridge

The Nernst equation is used to calculate the electrode or cell potential under non-standard conditions.

\[E_{cell}=E_{cell}^\circ-\frac{RT}{nF}\ln Q\]

At 25°C, it becomes:

\[E_{cell}=E_{cell}^\circ-\frac{0.0591}{n}\log Q\]

Where E°cell is the standard cell potential, n is the number of electrons transferred, and Q is the reaction quotient.

The equation helps determine the direction and spontaneity of a reaction:

• Ecell > 0 → spontaneous
• Ecell = 0 → equilibrium (Q = K)

It also relates to Gibbs energy:

ΔG = −nFE_cell

Thus, the Nernst equation is important for electrochemical calculations and equilibrium analysis.

The Nernst equation is used to calculate the electrode or cell potential under non-standard conditions.

\[E_{cell}=E_{cell}^\circ-\frac{RT}{nF}\ln Q\]

At 25°C, it becomes:

\[E_{cell}=E_{cell}^\circ-\frac{0.0591}{n}\log Q\]

Where E°cell is the standard cell potential, n is the number of electrons transferred, and Q is the reaction quotient.

The equation helps determine the direction and spontaneity of a reaction:

• Ecell > 0 → spontaneous
• Ecell = 0 → equilibrium (Q = K)

It also relates to Gibbs energy:

ΔG = −nFE_cell

Thus, the Nernst equation is important for electrochemical calculations and equilibrium analysis.

Electrical conductance and resistance:

\[\mathrm{K}=\mathrm{G}\frac{l}{A}\]

K = Conductivity

G = Conductance

\[\mathrm{G}=\mathrm{}\frac{1}{R}\]

R = Resistance

\[\mathrm{K}=\mathrm{}\frac{l}{RA}\]

Reactions

• Anode:
  2H₂ + 4OH⁻ → 4H₂O + 4e⁻
• Cathode:
  O₂ + 4H₂O + 4e⁻ → 4OH⁻
• Overall reaction:
  2H₂ + O₂ → 2H₂O

Applications

• Spacecraft (electric power)
• Power generators (homes, hospitals)
• Automobiles (experimental)
• Clean energy for industries

Drawbacks

• Hydrogen gas is hazardous
• High cost of hydrogen preparation