CUET UG Chemistry Booster Test – 3 Electrode Potential and Thermodynamics

According to NCERT, when a metal rod is dipped into a solution containing its ions, two opposing processes occur simultaneously. Metal atoms may lose electrons and pass into solution as positive ions, while metal ions present in the solution may gain electrons and get deposited on the metal surface. These competing tendencies establish an equilibrium and create a potential difference between the metal and the solution, known as electrode potential. Depending on which process predominates, the electrode may acquire a positive or negative charge relative to the electrolyte. Therefore, statements 1, 2 and 4 are correct. Statement 3 is incorrect because standard electrode potential is defined under standard conditions, where ionic concentration is 1 mol L⁻¹, gases are at 1 bar pressure and temperature is generally 298 K. The concentration is never taken as zero because electrochemical equilibrium cannot be established without ions. Thus, NCERT clearly supports statements 1, 2 and 4, making option B the correct answer.

Standard electrode potentials are measured relative to the standard hydrogen electrode (SHE), whose standard reduction potential is assigned a value of 0.00 V. A metal that readily loses electrons and undergoes oxidation has a low tendency to gain electrons. Consequently, its reduction potential becomes negative relative to SHE. According to NCERT, metals such as lithium, sodium and potassium exhibit highly negative standard reduction potentials because they strongly prefer oxidation. Such metals are excellent reducing agents. The more readily a metal loses electrons, the more negative its reduction potential becomes. Therefore, if an electrode has a higher intrinsic tendency for oxidation than reduction under standard conditions, its standard electrode potential will generally be negative relative to SHE. This relationship forms the basis of the electrochemical series and helps predict the direction of electron flow in electrochemical cells. Hence, option B is correct.

The electrochemical series arranges electrodes according to their standard reduction potentials. Species with highly positive reduction potentials have a strong tendency to gain electrons and therefore act as powerful oxidising agents. Fluorine possesses the highest standard reduction potential (+2.87 V), making F₂ the strongest oxidising agent in the series. Conversely, lithium has a highly negative standard reduction potential (−3.05 V), indicating a strong tendency to lose electrons. Therefore, lithium metal is the strongest reducing agent. NCERT emphasizes that oxidising strength increases with increasing reduction potential, while reducing strength increases with decreasing reduction potential. This understanding is essential for predicting redox reactions and determining which species will undergo oxidation or reduction spontaneously. Since fluorine is at the top and lithium at the bottom of the electrochemical series, option D correctly describes their chemical behavior.

Certain redox systems contain only ions, gases or liquids and therefore require an inert electrode to facilitate electron transfer. In the bromine/bromide half-cell, neither bromine nor bromide ions can serve as a metallic conducting electrode. Therefore, an inert platinum electrode is used. According to NCERT conventions, the inert electrode is written first, followed by the oxidised and reduced forms of the redox couple. Hence, the correct representation is Pt(s) | Br₂(aq) | Br⁻(aq). Platinum participates only as a conductor and does not undergo any chemical change during cell operation. This notation clearly indicates the presence of the bromine/bromide redox system and the supporting inert electrode. Proper understanding of cell notation is important for writing half-cell reactions and complete electrochemical cell representations. Thus, option C is the correct answer.

The Standard Hydrogen Electrode (SHE), represented as Pt(s) | H₂(g) | H⁺(aq), serves as the reference electrode and is assigned a standard electrode potential of 0.00 V. Therefore, item 1 matches with characteristic c. Fluorine possesses the highest positive standard reduction potential (+2.87 V), making item 2 correspond to characteristic a. Lithium exhibits a highly negative standard reduction potential (−3.05 V), so item 3 matches with characteristic d. Platinum and gold are inert electrodes that provide a conducting surface without participating in the reaction, making item 4 correspond to characteristic b. These facts are directly taken from NCERT's discussion of standard electrode potentials and the electrochemical series. Understanding these characteristic values helps compare oxidising and reducing strengths of chemical species and predict spontaneity in electrochemical cells. Therefore, the correct matching is 1-c, 2-a, 3-d and 4-b.

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.

According to NCERT, the reaction quotient Q is written using only the concentrations of species present in aqueous or gaseous form. Pure solids and pure liquids are assigned unit activity and therefore do not appear in the expression. For the reaction: nickel metal and silver metal are solids and are excluded from Q. The concentration of the product ion Ni²⁺ appears in the numerator, while the concentration of the reactant Ag⁺ appears in the denominator. Since the stoichiometric coefficient of Ag⁺ is 2, its concentration is raised to the power 2. Thus, The denominator therefore contains . This form is directly used in the Nernst equation to calculate cell potential under non-standard conditions. Hence option B is correct.

The Nernst equation contains the factor: where R is the gas constant, T is temperature, n is the number of electrons transferred and F is the Faraday constant. Considering the units: Multiplying R and T gives units of J mol⁻¹. Dividing by F gives: A joule per coulomb is defined as one volt. Therefore, the term has units of volts. This is consistent with the fact that the Nernst equation calculates electrode potential or cell potential, which is measured in volts. Thus, option C is correct.

The Nernst equation shows that cell potential depends on the reaction quotient Q. For the reaction involving magnesium and silver ions: the reaction quotient is: Since the concentration of Ag⁺ is only 0.0001 M, the denominator becomes extremely small. As a result, the value of Q becomes very large. According to the Nernst equation: a larger Q increases the subtraction term, causing a significant decrease in cell potential. The reduction from 3.17 V to 2.96 V is therefore a direct consequence of the very low silver ion concentration. NCERT emphasizes that concentration changes strongly influence electrode and cell potentials. Hence option B is correct.

For each electrochemical cell, the reaction quotient is obtained from the balanced cell reaction. Pure solids are omitted. For Zn|Zn²⁺||Cu²⁺|Cu: For Ni|Ni²⁺||Ag⁺|Ag: For Mg|Mg²⁺||Ag⁺|Ag: For Cu|Cu²⁺||Ag⁺|Ag: Matching these expressions gives: 1-b, 2-d, 3-a and 4-c. Thus option D is correct. This procedure is frequently used in NCERT numerical problems involving the Nernst equation.

NCERT explains that an electrochemical cell reaches equilibrium when there is no net tendency for the reaction to proceed in either direction. At this stage, the reaction quotient becomes equal to the equilibrium constant: Substituting this condition into the Nernst equation results in a cell potential of zero. This equilibrium condition does not require reactants to be completely consumed. In fact, both reactants and products coexist at equilibrium. Likewise, does not necessarily equal because standard conditions are generally different from equilibrium conditions. The evaporation of cell compartments has no thermodynamic significance in defining equilibrium. Therefore, only statement 2 correctly identifies the condition of thermodynamic equilibrium. Hence, option A is correct.

According to NCERT, at equilibrium the cell potential becomes zero because there is no net tendency for the reaction to proceed in either direction. Starting from the Nernst equation: At equilibrium: and Substituting these conditions gives: Rearranging: This important NCERT equation connects the standard cell potential with the equilibrium constant. It shows that a larger positive standard emf corresponds to a larger equilibrium constant and a more product-favoured reaction. Therefore, option B is correct.

NCERT explains the relationship: When the stoichiometric equation is multiplied by a factor, the number of electrons transferred (n) also increases by the same factor. Consequently, changes proportionally because it depends directly on n. Therefore, Gibbs free energy is an extensive property. However, the cell potential remains unchanged because it represents the driving force per unit charge and is independent of the amount of reactants involved. For example, doubling a Daniell cell reaction doubles but leaves unchanged. This distinction between extensive and intensive properties is fundamental in thermodynamics and electrochemistry. Therefore, option C correctly explains the difference.

The NCERT relation connecting Gibbs free energy and cell potential is: For the Daniell cell reaction: the number of electrons transferred is 2. However, the given reaction is multiplied by 2: Therefore, the total number of electrons transferred becomes: Substituting into the equation: The negative sign indicates that a spontaneous galvanic cell releases free energy. Since the reaction has been doubled, the Gibbs energy change doubles, while the cell potential remains unchanged. Hence option B is correct.

NCERT derives the relationship: at 298 K. This equation shows that is directly proportional to the standard cell potential. Therefore, larger positive values of correspond to larger equilibrium constants. To arrange equilibrium constants in increasing order, we simply arrange the standard cell potentials in increasing order: Using the numbering provided: Therefore, the equilibrium constants follow exactly the same order. The reaction with will have the largest equilibrium constant, while the reaction with will have the smallest. Hence option A is correct.

Two important NCERT equations are: and Since both expressions represent the same standard Gibbs free energy change, they can be equated: This equation forms a bridge between electrochemistry and chemical thermodynamics. It shows how the standard cell potential determines spontaneity and equilibrium. A positive standard cell potential corresponds to a negative Gibbs free energy change and a large equilibrium constant. Conversely, a negative standard cell potential corresponds to a positive Gibbs free energy change and a small equilibrium constant. This relationship is one of the most important results in the NCERT Electrochemistry chapter. Therefore, option D is correct.