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NEET 2022 | Class 12

NEET Important Chapter - Electrochemistry

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Introduction to Electrochemistry

Introduction to Electrochemistry


Last updated date: 23rd May 2024
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Electrochemistry is a discipline of physical chemistry concerned with the relationship between electrical potential as a quantifiable and quantitative phenomena and recognisable chemical change, with electrical potential as an outcome of a specific chemical change or vice versa.


Electrons move between electrodes via an electronically conducting phase (usually, but not always, an external electrical circuit like as in electroless plating), which is separated by an ionically conducting and electrically insulating electrolyte (or ionic species in a solution). Electrochemistry consists of different topics and different solved examples such as redox reaction NEET questions, and different solved examples and previous year questions.


Topics of Electrochemistry

  1. Redox Reactions:

Redox stands for "reduction-oxidation." These reactions involve the transfer of electrons between reactants. In a redox reaction, one substance gets oxidized, losing electrons, while another gets reduced, gaining those electrons. This exchange of electrons powers various chemical processes, from batteries and corrosion to photosynthesis and respiration. Understanding redox reactions is vital for comprehending the flow of energy in chemical systems. It plays a crucial role in electrochemistry, where it's employed in applications like electrolysis, galvanic cells, and fuel cells, driving technological advancements and chemical transformations in the real world.


  1. Conductance of Electrolytic Solutions:

It refers to the ability of an electrolytic solution to conduct electricity. The conductance is influenced by the concentration of ions in the solution and their mobility. Strong electrolytes, which dissociate completely into ions, exhibit high conductance, while weak electrolytes have limited ion dissociation and lower conductance.


Understanding conductance is essential in various applications, from designing electrochemical cells and batteries to predicting the behavior of ionic compounds in solution. It's a crucial aspect of electrochemistry, helping us harness the electrical conductivity of solutions for practical purposes and explaining the intricate world of ionic behavior in chemical systems.


  1. Specific and molar conductivity variation of conductivity in concentration:

Specific conductivity is the measure of a solution's ability to conduct electricity, while molar conductivity accounts for the concentration of ions in a given volume. With increasing concentration in an electrolytic solution, the Specific Conductivity typically increases. This is due to the greater number of ions available to carry the electric current. However, Molar Conductivity tends to decrease as concentration rises, indicating that each ion contributes less to conductivity due to increased ion-ion interactions. This variation is critical for designing electrochemical cells and predicting conductive properties, serving as a fundamental concept in electrochemistry.


  1. Kohlrausch’s Law:

Kohlrausch's Law is a fundamental principle that describes the molar conductivity of an electrolyte at infinite dilution. It states that the molar conductivity of an electrolyte at infinite dilution is the sum of the individual molar conductivities of its constituent ions, each multiplied by the number of ions produced when the electrolyte dissociates.


Kohlrausch's Law is a cornerstone in understanding how ions behave in solution and is crucial for accurately calculating the molar conductivity of strong electrolytes, aiding in the study of conductivity in electrolytic solutions, and predicting the behavior of ions in various chemical processes.


  1. Electrolysis and Laws of Electrolysis:

Electrolysis is a process where electrical energy is used to induce chemical reactions in an electrolyte. When an electric current passes through the electrolyte, it causes the ions to migrate and undergo redox reactions at the electrodes. This process is governed by Faraday's Laws of Electrolysis.


Faraday's First Law states that the amount of substance liberated at an electrode is directly proportional to the quantity of electricity passed through the electrolyte. Faraday's Second Law relates the amount of substance deposited during electrolysis to the equivalent weights of the substances involved. Understanding these laws is crucial for various applications, such as electroplating, purification of metals, and industrial processes, making them integral to the field of electrochemistry.


  1. Dry cell-electrolytic cells and Galvanic cells; lead accumulator:

Dry Cells, Electrolytic Cells, and Galvanic Cells are key concepts. Dry cells, like the common household batteries, are portable power sources that use a chemical reaction to generate electricity. Electrolytic Cells are where non-spontaneous redox reactions occur with the input of electrical energy. They are crucial for processes like electroplating. 


Galvanic Cells, on the other hand, are spontaneous redox reactions that produce electricity. They are the foundation of batteries, including lead-acid accumulators used in vehicles, where chemical reactions between lead and lead dioxide generate electrical energy. Understanding these cell types is essential for applications like energy storage, electroplating, and various electrochemical processes.


  1. EMF of a cell:

Electromotive Force (EMF) of a cell is a crucial concept. It represents the cell's ability to drive electrons through an external circuit. The EMF is responsible for the flow of electric current in both galvanic and electrolytic cells. For a galvanic cell (like a battery), the EMF arises from spontaneous chemical reactions. In an electrolytic cell, it's applied externally to initiate non-spontaneous reactions. Understanding the EMF of a cell is vital in designing and optimizing electrochemical devices, such as batteries and fuel cells, as it dictates their performance and efficiency in providing electrical energy for various applications.


  1. Standard Electrode Potential:

Standard Electrode Potential is a critical concept. It refers to the electrode potential of a half-cell under standard conditions, where all ions are at a concentration of 1 Molar, gases at 1 bar pressure, and a temperature of 25°C. The Standard Hydrogen Electrode (SHE) serves as the reference point, with an assigned potential of 0 volts. Standard electrode potential values help us compare the tendency of different electrodes to gain or lose electrons in redox reactions. A positive potential indicates a tendency to undergo reduction, while a negative value suggests oxidation. Understanding standard electrode potentials is essential in predicting the direction and feasibility of various electrochemical reactions, aiding in the study of electrochemistry and its practical applications.


  1. Relation between Gibbs energy change and EMF of a cell:

There's a profound connection between Gibbs Energy and the Electromotive Force (EMF) of a cell. Gibbs energy, a measure of a system's capacity to perform work, is closely linked to EMF, which drives the flow of electrons in electrochemical cells. The relationship can be expressed through the Nernst equation, which relates EMF to the Gibbs energy change, temperature, and reaction quotient. This equation helps us understand the thermodynamic feasibility of redox reactions in cells. By studying this connection, we gain insights into the spontaneity and energy transformations within electrochemical processes, facilitating our comprehension of cell potentials and their practical applications in energy generation and storage.


  1.  Fuel cells; corrosion:

Fuel Cells and Corrosion are significant topics. Fuel cells are electrochemical devices that convert the energy from chemical reactions, typically involving hydrogen and oxygen, into electrical power. They offer efficient and clean energy solutions, with applications in vehicles and power generation. 


Corrosion, on the other hand, is the destructive process where metals deteriorate due to chemical reactions with the environment. Understanding corrosion is crucial for preventing and mitigating the degradation of structures and materials. Both fuel cells and corrosion illustrate the practical implications of electrochemistry, from sustainable energy production to preserving the integrity of metals and infrastructure.


Important Topics of Electrochemistry

  • Electrochemical Cell

  • Electrode Potential

  • Cell Potential or EMF of Cell

  • Nernst Equation

  • Kohlrausch’s Law

  • Electrolytic Conductance


Important Definitions of Electrochemistry

Parameter

Explanation

Electrochemistry

The study of generating electricity from the energy released during a spontaneous chemical reaction, as well as the application of electrical energy to non-spontaneous chemical changes, is known as electrochemistry.

Electrochemical Cells

A spontaneous chemical reaction is one that can happen on its own, and the system's Gibbs energy decreases as a result. After that, the energy is converted into electrical energy. These interconversions are carried out using electrochemical cells.

Electrolytic Cell

This cell's electrodes are immersed in an electrolytic solution containing both cations and anions. When current is applied, the ions migrate to electrodes with opposite polarity, where they are reduced and oxidised simultaneously.

Electrode Potential

When an element comes into contact with its own ions, it tends to lose or gain electrons, leading it to become positively or negatively charged.

Electrochemical Series

The half-cell potential values are standard, and they are represented as standard reduction potential values in the Electrochemical Series table at the end.

Batteries

The term "battery" refers to a system in which Galvanic cells are wired together in series to create a higher voltage.

Conductance (G)

It is the reciprocal of resistance and is defined as the ease with which electric current travels through a conductor.

Conductivity

It's the inverse of resistivity (⍴).

Electrolyte

An electrolyte is a material that dissociates in solution to produce ions and so conducts electricity when dissolved or molten.

Kohlrausch’s Law

The limiting molar conductivity of an electrolyte is defined as the sum of the individual contributions of the electrolyte's anion and cation.


Types of Electrochemical Cell

  • A spontaneous chemical reaction is one that can occur on its own and reduces the system's Gibbs energy. 

  • The energy is then transformed into electrical energy. Electrochemical cells are used to carry out these interconversions.

  • External energy in the form of electrical energy can also be used to cause non-spontaneous processes to occur.

  • There are two types of electrochemical cells: galvanic and electrolytic cells. Galvanic cells convert chemical energy into electrical energy, whereas electrolytic cells turn electrical energy into chemical energy.


Galvanic Cell

  • To extract cell energy, a spontaneous chemical process or reaction is employed, which is subsequently converted to electric current.

  • A Daniell Cell, for example, is a Galvanic Cell that performs the redox process utilising Zinc and Copper.

  • The reaction of Daniell Cell is given below:
    Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

  • The two halves of the reaction are,
    Oxidation Half: Zn(s) → Zn2+(aq) + 2e-
    Reduction Half: Cu2+(aq) + 2e- → Cu(s)

  • These reactions take place separately.

  • Zn is the reducing agent, whereas Cu2+ is the oxidising agent.

  • Half cells are also known as electrodes. The oxidation half is the anode, while the reduction half is the cathode.

  • Electrons travel from the anode to the cathode in the external circuit. The anode is given negative polarity. The cathode is given positive polarity.

  • Daniell Cell is a fictional character that Daniell Cell has developed. The cathode is Cu, and the anode is Zn.


Electrode Potential

  • The discharge process becomes competitive when more than one cation or anion is present. 

  • Any ion that needs to be discharged requires energy, and if numerous ions are present, the one that requires the most energy is discharged first.

  • When an element comes into contact with its own ions, it tends to lose or gain electrons, leading it to become positively or negatively charged.

  • The electrode potential is referred to as oxidation or reduction potential depending on whether oxidation or reduction has happened.

  • Characteristics:

    • The oxidation and reduction potentials have the same magnitude and sign.

    • The values of E do not add up since it is not a thermodynamic characteristic.


Standard Electrode Potential (Eo)

  • It's defined as the electrode potential of a given electrode when compared to a standard hydrogen electrode under standard conditions.

  • The standard conditions are as follows:

    • Each ion in the solution has a concentration of 1M.

    • A temperature of 298 K.

    • The pressure of each gas is one bar.


Cell Potential or EMF of a Cell

  • The difference between the electrode potentials of two half cells is known as cell potential. 

  • Electromotive force occurs when no current is drawn from the cell (EMF).

  • Ecell = Ecathode + Eanode. We use the anode's oxidation potential and the cathode's reduction potential in this equation.

  • Because the anode is on the left and the cathode is on the right, the following is the result:
    = ER + EL

  • As a result, for a Daniel Cell:
    Eocell = EoCu2+/Cu - EoZn/Zn2+ = 0.34 + (0.76) = 1.10 V.


Nernst Equation

  • It connects the electrode voltage with the ion concentration. As a result, the reduction potential grows in tandem with the ion concentration. 

  • For a sort of electrochemical reaction in general,
    aA + bB →cC + dD.

  • The Nernst equation is as follows:
    Ecell = Eocell - $\frac{2.303}{nF}.RT.log\frac{\left [ C \right ]^{c}\left [ D \right ]^{d}}{\left [ A \right ]^{a}\left [ B \right ]^{b}}$


Factors Affecting Electrolytic Conductance

  • An electrolyte is a material that dissociates in solution to produce ions and so conducts electricity when dissolved or molten.

  • Strong electrolytes, such as HCl, NaOH, KCl, and weak electrolytes, such as CH3COOH, NH4OH, are examples.

  • The conductance of electricity by ions in solutions is known as electrolytic or ionic conductance. 

  • The following parameters influence the flow of electricity via an electrolyte solution.

  • Electrolyte Nature or Interionic Attractions: The greater the freedom of ion mobility and the higher the conductance, the lower the solute-solute interactions.

  • Ion Solvation: The extent of solvation increases as the number of solute-solvent interactions grows, and the electrical conductance drops.

  • The Nature of the Solvent and its Viscosity: The higher the viscosity and the greater the solvent's resistance to ion flow, and hence the lower the electrical conductance, the larger the solvent-solvent interactions are.

  • Temperature: Solute-solute, solute-solvent, and solvent-solvent interactions decrease when the temperature of an electrolytic solution rises, leading electrolytic conductance to rise.


Solved Examples from Chapter

Question 1: Find the charge in coulomb on 1 g-ion of N3-.

Solution: 

  • An ion is formed by the gain or loss of an electron. 

  • Therefore, Charge of an ion = Charge of an electron
    ∴ Charge of an ion = 1.6 x 10-19 Coulomb (C) .

  • Following this, the charge on the given N3- ion is,
    Charge on N3- ion = 3 x 1.6 x 10-19.

  • The number of ions in 1gram of N3- is given by the Avogadro’s Number = 6.02 x 1023 ions. 

  • Therefore, the total charge in 1 g of N3- is given as
    Charge in 1g of N3- = 3 x 1.6 x 10-19 x 6.02 x 1023 = 2.89 x 105 Coulomb (C) .

Hence, the answer is 2.89 x 105 Coulomb (C).

Key Points to Remember: The charge of an electron and the amount of constituent particles in a given unit amount of sample i.e. the Avogadro’s number are the two important concepts in use in the solution.


Question 2: Estimate the Eo from the half-reaction M+(aq) + e- → M(s) based on the following observations:

(i) M interacts with H2SO4(aq) but not with HI(aq); M displaces Au+(aq) but not Fe3+(aq).

(ii) The metal M reacts with HI(aq) to produce H2(g), but neither Al3+(aq) nor Na+(aq) are displaced.

Solution: 

  • (i) When a metal dissolves in H2SO4, it has a lower reduction potential than ESO42-(aq)/SO2(g) = 0.17 V. 

  • It has a reduction potential greater than EH+(aq)/H2(g) = 0 V if it does not dissolve in HI. 

  • It has a reduction potential smaller than EAu+(aq)/Au(s) = 1.68 V if it displaces Au+(aq) from solution.

  • However, if it does not remove Fe3+(aq) from solution, its reduction potential is greater than,
    EoFe3+(aq)/Fe2+(s) = 0.769 V.

  • As a result,  0 V < Eo < 0.17 V .

  • (ii) When a metal dissolves in HI(aq), it has a lower reduction potential than EH+(aq)/H2(g) = 0 V. 

  • Its reduction potential is greater than EAl3+(aq)/Al(s) = 1.676 V if it does not displace Al3+(aq) from solution.

  • Its reduction potential is greater than ENa+(aq)/Na(s) = 2.7144 V if it does not displace Na+(aq) from solution. As a result, -1.7676 < E0 < 0 V.


Key Points to remember: Electrode potential is different for different sets of ions and whether they are undergoing an oxidation or a reduction reaction. 


Solved Examples from Previous Year Questions

Question 1: Which of the following relationships for the values of ΔGo and Keq is valid if Eocell for a particular reaction is negative?

(a) ΔGo < 0; Keq > 1

(b) ΔGo < 0; Keq < 1

(c) ΔGo > 0; Keq < 1

(d) ΔGo > 0; Keq > 1

Solution: 

  • The general relationship between free energy, ΔGo and Eocell is given as,
    ΔGo = -nFEocell.

  • Also, the relationship between ΔGo and Keq at equilibrium is given as;

    • ΔGo > 0 then Keq < 1 and at equilibrium, reactants are preferred above products.

    • ΔGo = 0 then Keq = 1 and products and reactants are equally preferred in an equilibrium state.

    • ΔGo < 0 then Keq > 1 and at equilibrium, products are preferred over reactants.

  • Since, the given value of Eocell is negative, the resulting value of ΔGo is positive.
    ∴ ΔGo > 0

  • Since, ΔGo > 0 then Keq < 1.

  • As a result, option (c) is the correct answer.


Question 2: For the cell reaction

2Fe3+(aq) + 2I-(aq) → 2Fe2+(aq) + I2(aq).

E-cell = 0.24 V at 298 K. The standard Gibbs Energy (ΔrG-) of the cell reaction is:

(Given that Faraday Constant is F = 96,500 C mol-1)

(a) - 46.32 kJ/mol

(b) - 23.6 kJ/mol

(c) 46.32 kJ/mol

(d) 23.16 kJ/mol.

Solution: 

  • Using the formula for free energy,
    ΔrG- = -nFEocell

  • ΔrG- = - 2 x 96,500 x 0.24 J mol-1.

  • ΔrG- = - 46,320 J mol-1

  • ΔrG- = - 46.32 Jk mol-1 

Hence, the final answer is (d) ΔrG- = - 46.32 kJ mol-1 


Question 3: Without losing its concentration ZnCl2 solution cannot be kept in contact with: 

(a) Au

(b) Al

(c) Pb

(d) Ag

Solution: 

  • Al is situated above Zn in the electrochemical series, while all other elements are found below Zn. 

  • As a result, zinc is displaced from the ZnCl2 solution by aluminium. As a result, it is unable to communicate with Al.

  • And the reaction is 2Al + 3ZnCl2 → 2AlCl3 + 3Zn.

  • As a result, option (b) is the correct answer.


Practice Questions

Question 1: When a copper wire is submerged in an AgNO3 solution, the solution turns blue because copper:

(a) With AgNO3, creates a soluble compound.

(b) Is converted to Cu2+ by oxidation

(c) Is converted to Cu2+ by reduction

(d) Splits and dissolves into atomic form

Answer: (b) Is converted to Cu2+ by oxidation


Question 2: The half-cell reactions are listed below:

Mn2+ + 2e- → Mn; Eo = -1.18 V

2Mn3+ + 2e- → 2Mn2+; Eo = +1.51 V.

Therefore, 3Mn2+ → 2Mn3+ + Mn will be,

(a) -2.69 V; The reaction will not take place.

(b) -2.69 V; The reaction will take place.

(c) -0.33 V; The reaction will not take place.

(d) -0.33 V; The reaction will take place.

Answer: (a) -2.69 V; The reaction will not take place.


Conclusion of Electrochemistry NEET

The study of chemical reactions that cause electrons to move is known as electrochemistry. This flow of electrons is known as electricity, and it can be generated by electrons moving from one element to another in an oxidation-reduction reaction. The electrochemistry NEET 2022 notes thus are valuable to your preparation. 


Delving into the complexities of electrochemistry proves indispensable for NEET aspirants. This chapter not only illuminates the fundamental principles governing chemical processes but also lays the groundwork for understanding various biological and industrial applications. Aspiring medical students should grasp the intricacies of electrochemical concepts, ensuring a solid foundation for success in the NEET examination.

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FAQs on NEET Important Chapter - Electrochemistry

FAQ

1. What is electrochemistry, and what is an example of it?

The study of chemical reactions that cause electrons to move is known as electrochemistry. It's all about how electrical energy interacts with chemical transformation. Electrochemistry, for example, is concerned with the study of electrochemical cells. It is concerned with cells that convert chemical to electrical energy.

2. What purpose does electrochemistry serve?

In everyday life, electrochemistry is used in a variety of ways. Chemical reactions are utilised to generate electricity in all types of batteries, from flashlights to calculators to automobiles. Decorative metals like gold and chromium are applied to things using electricity.

3. What does electrochemistry have to do with anything?

Electrochemistry plays a key role in a variety of important technological applications. Batteries, for example, are critical not just for storing energy for mobile devices and cars, but also for load levelling, which allows renewable energy conversion technologies to be used.