Unit: Electrochemistry ⚡

Electrochemistry is the study of the relationship between electrical energy and chemical reactions.
It involves redox reactions, where electrons are transferred, leading to the conversion of chemical energy into electrical energy or vice versa.
The unit covers electrochemical cells, electrolysis, and the Nernst equation.

Electrochemical cells are devices that convert chemical energy into electrical energy. They consist of two half-cells connected by a salt bridge.
Galvanic cells or Voltaic cells produce electricity from spontaneous chemical reactions.
Example: Daniel cell, which consists of a zinc and copper electrode.

Electrolytic cells use external electric current to drive non-spontaneous reactions.
- Example: Electrolysis of water to produce hydrogen and oxygen.
Electrolysis involves the breaking of chemical bonds by the application of electric current.
- At the anode (positive electrode): Oxidation occurs.
- At the cathode (negative electrode): Reduction occurs.
Faraday’s Laws of Electrolysis:
- First law: The amount of substance deposited at the electrode is directly proportional to the quantity of charge passed through the electrolyte.
- Second law: The amounts of different substances deposited are proportional to their equivalent weights.

Conductance is the ability of a substance to conduct electricity. It depends on the number of ions in a solution.
Factors affecting conductivity:
1. Concentration of ions: Higher ion concentration increases conductivity.
2. Nature of the solvent: Some solvents, like water, are better conductors.
3. Temperature: Higher temperatures generally increase ion mobility, thus increasing conductivity.
Specific conductance (κ) refers to the conductivity of a solution with a unit length and unit area.
Molar conductance is the conductivity of a solution containing 1 mole of electrolyte dissolved in a given volume.

The electrochemical series is a list of elements (mainly metals) arranged based on their standard electrode potential.
Standard electrode potential is the potential difference when the electrode is in contact with a solution of its ions.
- More positive electrode potential means a stronger tendency to gain electrons (reduction).
- More negative electrode potential means a stronger tendency to lose electrons (oxidation).
Applications of the electrochemical series:
- It helps predict the direction of redox reactions.
- Used in understanding the strength of reducing or oxidizing agents.

The Nernst equation relates the electrode potential of a half-cell to the concentration of ions involved in the reaction.


  E = E^o - \frac{0.0591}{n} \log \left( \frac{[C]^c[D]^d}{[A]^a[B]^b} \right)

E = Electrode potential under non-standard conditions
E₀ = Standard electrode potential
n = Number of electrons transferred
[C], [D], [A], [B] = Concentrations of the reactants and products
The Nernst equation is essential for calculating the potential of electrochemical cells under different conditions.

Gibbs Free Energy (ΔG) determines whether a process is spontaneous or non-spontaneous.
- ΔG < 0: The reaction is spontaneous.
- ΔG > 0: The reaction is non-spontaneous.
Relationship between Gibbs Free Energy and Electrode Potential:


  \Delta G = -nFE

Where:

n = Number of moles of electrons transferred
F = Faraday’s constant (96,485 C/mol)
E = Electrode potential
Spontaneous reactions have a positive electrode potential and a negative Gibbs free energy.

Batteries:
- Primary batteries (e.g., dry cells) provide energy through a galvanic cell.
- Secondary batteries (e.g., lead-acid battery) can be recharged through electrolysis.
Electroplating:
- Using electrolysis to deposit a layer of metal onto another surface (e.g., gold plating on jewelry).
Corrosion:
- Electrochemical processes are involved in the corrosion of metals, such as rusting of iron.
Electrosynthesis:
- Electrochemical reactions are used to synthesize chemicals (e.g., chlorine gas from electrolysis of sodium chloride solution).

Electrochemistry plays a crucial role in various fields, including energy storage (batteries), metallurgy (extraction of metals), corrosion, and electroplating.
Understanding the concepts of electrochemical cells, Nernst equation, and Gibbs free energy allows us to predict and control chemical reactions involving electron transfer.