CUET UG Chemistry Booster Test – 2 Electrode Potential and Thermodynamics

At the electrode-electrolyte interface, two opposing processes occur simultaneously. Metal ions present in the solution tend to gain electrons and deposit on the metal surface, while metal atoms from the electrode may lose electrons and enter the solution as ions. When positive metal ions deposit on the electrode, they reduce the excess negative charge present on the metal surface and contribute toward making the electrode relatively positive. NCERT explains that these competing processes continue until equilibrium is established. The resulting separation of charges gives rise to the electrode potential. The tendency of metal ions to deposit on the electrode is therefore associated with making the electrode positively charged. This process is one of the key factors responsible for the development of the potential difference between the electrode and the electrolyte.

NCERT explains that at the electrode-electrolyte interface two competing processes occur continuously. Metal ions from solution may deposit on the electrode and tend to make it positively charged. Conversely, metal atoms from the electrode may enter the solution as positive ions, leaving excess electrons behind and tending to make the electrode negatively charged. Eventually, an equilibrium is established where the rates of deposition and dissolution become equal. At this point, there is no net current flow across the interface. The resulting charge separation between the electrode and electrolyte generates the electrode potential. This balance between opposing tendencies is fundamental to understanding the origin of electrode potentials and electrochemical cells. Therefore, the correct matching is 1-c, 2-d, 3-a and 4-b.

Standard reduction potential measures the tendency of a species to gain electrons and undergo reduction. A higher positive reduction potential indicates a stronger oxidising agent. According to the electrochemical series given in NCERT, fluorine has the highest standard reduction potential of approximately +2.87 V, making it the strongest oxidising agent. Copper has a standard reduction potential of +0.34 V, hydrogen is assigned 0.00 V, and lithium has a highly negative value of about −3.04 V. Therefore, the decreasing order of standard reduction potential is: F₂/F⁻ > Cu²⁺/Cu > H⁺/H₂ > Li⁺/Li Using the numbering given in the question: 3 > 4 > 1 > 2 This sequence also corresponds to decreasing oxidising strength. Such comparisons are essential for predicting spontaneous redox reactions and determining electron flow in electrochemical cells.

Modern IUPAC recommendations standardize the terminology used in electrochemistry to avoid ambiguity. NCERT follows these conventions and refers to standard reduction potentials as standard electrode potentials. These values are measured relative to the Standard Hydrogen Electrode and are used to compare the tendencies of different species to undergo reduction. Using a single terminology ensures consistency across scientific literature, textbooks and research publications. Standard electrode potentials are tabulated extensively and form the basis of electrochemical calculations, prediction of reaction feasibility and determination of cell emf. Therefore, the accepted IUPAC term is standard electrode potential.

According to NCERT, the Standard Hydrogen Electrode (SHE) is used as the universal reference electrode for measuring electrode potentials. It consists of a platinum electrode coated with platinum black, immersed in a solution containing hydrogen ions of unit activity, while hydrogen gas is bubbled at a pressure of 1 bar. Since it is impossible to measure the absolute potential of a single electrode, a reference electrode is required. By international convention, the standard electrode potential of the SHE is assigned the value 0.00 V at 298 K. The electrode potentials of all other electrodes are measured relative to this reference. If an electrode has a positive standard reduction potential relative to SHE, it has a greater tendency to undergo reduction than hydrogen ions. Thus, the value assigned to SHE is 0.00 V, making option A correct.

The reaction 2H⁺(aq) + 2e⁻ → H₂(g) represents the gain of electrons by hydrogen ions. According to NCERT, reduction is defined as the gain of electrons. Since hydrogen ions accept two electrons and are converted into hydrogen gas, the process is a reduction half-reaction. This reaction is extremely important in electrochemistry because it forms the basis of the Standard Hydrogen Electrode (SHE), whose standard electrode potential is assigned a value of 0.00 V. All other standard electrode potentials are measured relative to this electrode. The equation represents only one half of a redox process because oxidation is not shown. Therefore, it is classified as a reduction half-reaction and not as an overall redox reaction.

The Nernst equation contains the reaction quotient, which includes the activities or concentrations of species participating in the reaction. According to NCERT, the activity of a pure solid is taken as unity. Since the metal M is present as a pure solid in the electrode reaction, its activity remains constant and does not affect the electrode potential. As a result, the term corresponding to the solid metal is omitted from the final form of the Nernst equation. This simplification is commonly used in electrochemistry because the concentration of a pure solid does not change significantly during the reaction. Therefore, the concentration of solid M is absent from the final expression because its activity is taken as unity.

The universal gas constant R is one of the fundamental constants used in chemistry and physics. In the Nernst equation, it relates temperature to the energy changes associated with electrochemical processes. NCERT gives the value of R as 8.314 J K⁻¹ mol⁻¹. The unit indicates that R relates energy (joules) to temperature (kelvin) per mole of substance. This constant appears in thermodynamics, gas laws and electrochemistry. Correct identification of its unit is important when deriving the Nernst equation and performing electrochemical calculations. Therefore, the correct unit of the gas constant is J K⁻¹ mol⁻¹.

For a galvanic cell, NCERT gives the expression: E°cell = E°cathode − E°anode In the Daniell cell, copper acts as the cathode and zinc acts as the anode. Therefore: E°cell = (+0.34 V) − (−0.76 V) E°cell = 0.34 + 0.76 E°cell = 1.10 V A positive value of cell emf indicates that the reaction is spontaneous under standard conditions. The Daniell cell is one of the most common examples of a galvanic cell and demonstrates the conversion of chemical energy into electrical energy. Therefore, the standard emf of the Daniell cell is 1.10 V.

The electromotive force (emf) of an electrochemical cell is the maximum potential difference between two electrodes when no current is drawn from the cell. Since emf is a form of electrical potential difference, its SI unit is the volt (V). One volt is defined as one joule of energy per coulomb of charge. NCERT consistently expresses electrode potentials, cell potentials and standard reduction potentials in volts. The unit allows comparison of the driving force of different electrochemical reactions and is fundamental to all electrochemical calculations. Therefore, the correct unit of cell potential difference is volt.

The Nernst equation in its general form is: where R is the gas constant, T is the temperature in Kelvin, F is the Faraday constant and Q is the reaction quotient. NCERT converts the natural logarithm (ln) into common logarithm (log₁₀) using the relation: Substituting T = 298 K, R = 8.314 J mol⁻¹ K⁻¹ and F = 96487 C mol⁻¹ gives: Therefore, the Nernst equation becomes: This simplified form is widely used in electrochemistry problems at room temperature. The constant 0.059 is not an independent physical constant but a derived value obtained by combining R, T, F and the conversion factor 2.303. Hence option A is correct according to NCERT.

The Nernst equation is one of the most important equations in electrochemistry. According to NCERT, it establishes a relationship between cell potential and the concentrations of reactants and products through the reaction quotient Q. Therefore statement 1 is correct. The equation is originally expressed using natural logarithms, and conversion to common logarithms is achieved using the relation ln x = 2.303 log x; hence statement 2 is correct. The quantity n represents the number of electrons transferred in the balanced redox reaction, making statement 4 correct. Statement 3 is incorrect because the Nernst equation is specifically used to determine electrode and cell potentials under non-standard conditions. Standard-state potentials are represented by E° values. Thus, statements 1, 2 and 4 are correct and statement 3 is incorrect. Therefore option C is the correct answer.

For the cell reaction: the reaction quotient is: According to the Nernst equation: When the concentration of Ag⁺ decreases, the denominator of Q becomes smaller. As a result, the value of Q increases. Since log Q increases, the subtraction term becomes larger and the cell potential decreases. This agrees with NCERT's discussion of concentration effects on cell potential. The cathode reaction depends directly on Ag⁺ ions, and reducing their concentration reduces the tendency of silver ions to gain electrons. Consequently, the driving force of the electrochemical reaction becomes weaker and Ecell decreases. Therefore option B is correct.

The Nernst equation contains several important quantities. The reaction quotient Q expresses the ratio of product concentrations to reactant concentrations raised to their stoichiometric powers. Therefore Q corresponds to the reaction quotient. The expression [C]ᶜ[D]ᵈ represents the product concentration term because products appear in the numerator of Q. Similarly, [A]ᵃ[B]ᵇ represents the reactant concentration term because reactants appear in the denominator of Q. The symbol n represents the number of electrons transferred in the balanced redox reaction. These definitions are directly used in NCERT while deriving and applying the Nernst equation. Matching these concepts correctly gives 1-b, 2-c, 3-d and 4-a. Hence option D is correct.

A galvanic cell generates electrical energy through a spontaneous redox reaction. As the reaction proceeds, reactants are consumed and products are formed. Consequently, reactant ion concentrations decrease while product ion concentrations increase. According to the Nernst equation, the reaction quotient Q continuously changes during cell operation. As Q increases, the value of the term increases, causing the cell potential to decrease. NCERT explains that the cell potential falls gradually until equilibrium is reached. At equilibrium, there is no net tendency for the reaction to proceed and the cell potential becomes zero. Therefore the gradual change in concentration is the fundamental reason for the decrease in cell voltage. This behavior reflects the close relationship between thermodynamics and electrochemistry discussed in the NCERT chapter. Hence option D is correct.

According to NCERT, the driving force for a galvanic cell is the free energy change associated with the redox reaction. As the cell operates, reactants are consumed and products are formed, causing the reaction quotient Q to increase. Eventually, the system reaches chemical equilibrium where the forward and reverse reactions occur at equal rates. At this stage, there is no net tendency for the reaction to proceed in either direction. The Gibbs free energy change becomes zero and the cell can no longer perform electrical work. Since the cell potential is directly related to the Gibbs free energy change through the relation ΔG = –nFEcell, the cell potential also becomes zero. Therefore, a voltmeter connected across the cell terminals will show no potential difference. This condition is an important consequence of the Nernst equation and demonstrates the relationship between electrochemistry and chemical equilibrium. Hence, option C is correct.

The standard cell potential is an intensive property and does not depend on the amount of substance involved in the reaction. According to NCERT, when a balanced electrochemical reaction is multiplied by any factor, the number of electrons transferred and the Gibbs free energy change are multiplied by the same factor. Since the relation between them is: both ΔG° and n increase proportionally. As a result, the value of E°cell remains unchanged. This behavior is similar to temperature, density and pressure, which are also intensive properties. For example, if a reaction with E°cell = 1.10 V is doubled, the value of E°cell remains 1.10 V and does not become 2.20 V. Understanding this principle is essential when calculating thermodynamic quantities and equilibrium constants using electrochemical data. Therefore, option C is correct.

The Gibbs free energy change represents the maximum useful work obtainable from a chemical process under specified conditions. In electrochemistry, NCERT explains that the relation: is valid when electrical energy is obtained reversibly from the cell. Under reversible conditions, the system remains extremely close to equilibrium at every stage, ensuring that the maximum amount of useful electrical work is extracted. If the process occurs irreversibly, some energy is dissipated as heat or lost through other inefficiencies, resulting in less useful work than predicted by ΔG. Therefore, reversible transfer of charge is essential for establishing the exact thermodynamic relationship between Gibbs free energy and cell potential. This principle forms the basis for connecting electrochemical measurements with thermodynamic quantities. Hence, option A is correct.

NCERT establishes the relationship between standard cell potential and the equilibrium constant using thermodynamic principles. A positive value of E°cell indicates that the cell reaction is spontaneous under standard conditions, making statement 1 correct. Since spontaneous reactions favor product formation, the equilibrium constant K is generally greater than 1, making statement 2 correct. At 298 K, the relation is: which makes statement 4 correct. Statement 3 is incorrect because E°cell is a standard-state quantity and does not become zero at equilibrium. Instead, the actual cell potential Ecell becomes zero when equilibrium is reached. This distinction between E°cell and Ecell is emphasized in NCERT and is frequently tested in examinations. Therefore, statements 1, 2 and 4 are correct, making option D the correct answer.

According to NCERT, the relationship between standard cell potential and equilibrium constant at 298 K is: For the given reaction: the number of electrons transferred is: Substituting the given values: Thus, the equilibrium constant is extremely large, indicating that the reaction strongly favors product formation under standard conditions. This result is consistent with the positive standard cell potential and the spontaneous nature of the reaction. Therefore, the correct answer is option B.