Electrochemistry Full Chapter Revision Guide

There may be only 4 seasons in Kerala, but higher secondary students get more than that! There is exam season, revision season, and much more! So, to welcome each season, let’s be prepared by taking a close look at the plus two chemistry electrochemistry chapter summary.

Electrochemistry is one of the highest-scoring and most concept-oriented chapters in Kerala Plus Two Chemistry. It connects chemical reactions with the flow of electricity, helping students understand how batteries, electrolysis, and redox reactions work.

In our world of scientific wonders, electrochemistry is the fascinating branch of chemistry that deals with the relationship between electrical energy and chemical reactions. To put it simply, it’s the study of 2 main processes:

1. Producing electrical energy from a chemical reaction: It’s the process of converting the energy released during a spontaneous chemical reaction into electricity. Typically, this is a process that happens on its own.
2. Using electrical energy to cause chemical reactions: It’s the process of using an external electricity source to drive a non-spontaneous chemical reaction.

To clear up the confusion, the core concept linking the 2 forms of energy is the movement of electrons (which is electricity) during a redox reaction.

In this full electrochemistry class 12 chapter revision blog, let’s walk through the key concepts, formulas, equations, and important board exam questions. Make sure to read and analyze till the last to face the exam seasons with confidence! .

2. Electrode potentials

Electrode potential is the measure of the tendency of an electrode to either gain or lose electrons when it is in contact with a solution of its own ions. Electrode potential depends on the nature of the metal, the ion concentration, and the temperature. The tendency of a metal to do this is measured as its electrode potential, Ecell (E°). It tells us how easily a metal gets oxidized or reduced. The Standard Electrode Potential (E°) is measured under standard conditions (1M, 25°C, 1 atm).

• If it easily loses electrons, → has a negative potential.

• If it easily gains electrons, → has positive potential.

3. Nernst Equation

While studying about electrochemical cells and redox reactions in class 12, it is important to understand the Nernst equation. The Nernst equation allows the calculation of the electrode potential (E) of a cell or half-cell under non-standard conditions. Which happens when the concentration or the partial pressure is unity, i.e, concentration ≠ 1 M

Ecell = E0 cell – RT/ nF ln Q

At standard temperature (298 K), the equation can be simplified to:

Ecell = E0cell – (0.0592/n) log Q

Components of the Nernst equation

Ecell : The cell potential

E0cell  : The standard cell potential

R : The gas constant

T : The temperature in Kelvin

n : The number of moles of electrons transferred

F : Faraday’s constant

Q : The reaction quotient 

4. Electrolysis and Faraday’s Law

Electrolysis can be simply termed as the decomposition of an electrolyte, and this can be accomplished by using electrical energy to cause a non-spontaneous chemical reaction!

Faraday’s Law of electrolysis is the relation between the amount of electricity passed and the amount of substance produced or consumed during electrolysis. There are basically 2 laws associated with Faraday’s law of electrolysis.

• 1st law: The mass of a substance deposited or liberated at any electrode is directly proportional to the quantity of electric charge passed through the electrolyte.

• 2nd law: When the same quantity of electricity passes through different electrolytes, the mass of substances deposited is proportional to their equivalent weights.

While the whole process sounds complex, it’s used for electroplating, metal extraction, and the purification of copper.

5. Batteries and Fuel Cells

Batteries are basically a galvanic cell in which the chemical energy of the redox reaction is converted to electrical energy. There are 2 types of batteries, primary and secondary.

• Primary batteries: These are only meant to be used once. The reaction is irreversible and stops when reactants are consumed. (Eg: Dry cell and mercury cell)

• Secondary batteries: These are rechargeable batteries, as they act as a galvanic cell while discharging and an electrolytic cell while charging (Eg: Lead storage battery and nickel-cadmium cell)

While primary and secondary batteries seem like the main characters, they are not! It’s the fuel cells that steal the spotlight when it comes to batteries.

• Fuel cells: These are galvanic cells that convert the energy of combustion of fuels like hydrogen, methane, methanol, etc., directly into electrical energy. (e.g., H₂-O₂ fuel cell).

The fuel cells are pollution-free and work continuously as long as reactants are supplied. The main and interesting thing to know about fuel cells is that they have been used in the Apollo space program.

6. Corrosion—A Real-Life Electrochemical Process

Many might confuse corrosion with the rusting of iron, but few understand that the rusting of iron is only an example of corrosion! Make sure you are among the few who know what corrosion really is.  

Slow destruction of metal due to the attack of atmospheric gases and moisture on its surface, resulting in the formation of compounds such as oxides, sulfides, carbonates, etc., is called corrosion. Some examples might include the rusting of iron, tarnishing of silver, and the formation of a green coating on copper and bronze.

As you might understand now, corrosion is a natural process. There are multiple ways to avoid or prevent corrosion. The one to keep in mind is the sacrificial method! An electrochemical method is to provide a sacrificial electrode of another metal, like Mg, Zn, etc., when it corrodes itself but saves the object, and this method is called the sacrificial method.