CUET UG Chemistry Booster Test – 3 Electrochemical Cells

In electrochemistry, the relationship between Gibbs free energy and cell potential is one of the most important NCERT concepts. For a non-standard electrochemical cell, the Gibbs free energy change is given by ΔrG = -nFEcell, where n is the number of electrons transferred and F is the Faraday constant. Under standard conditions, the relation becomes ΔrG° = -nFE°cell. The standard cell potential is also related to the equilibrium constant by the expression E°cell = (2.303RT/nF) log Kc. At equilibrium, the reaction has no tendency to proceed in either direction and therefore the cell potential becomes zero. This is represented by Ecell = 0. These relationships connect thermodynamics and electrochemistry and are extensively used in NCERT for predicting spontaneity, equilibrium conditions and the electrical energy produced by galvanic cells. Since each parameter matches uniquely with its corresponding relation, option A represents the correct matching sequence.

According to the NCERT discussion on the Nernst equation, the potential of an electrochemical cell depends on the concentrations of reactants and products involved in the cell reaction. Therefore statement 1 is correct. Statement 3 is also correct because the reaction quotient Q is written as the concentration of products raised to their stoichiometric powers divided by the concentration of reactants raised to their stoichiometric powers. For a Daniell cell, the expression Ecell = E°cell − (0.059/2) log([Zn²⁺]/[Cu²⁺]) shows that decreasing the concentration of Zn²⁺ decreases the value of the logarithmic term and therefore increases the cell potential. Hence statement 4 is correct. Statement 2 is incorrect because Gibbs free energy change depends on the amount of substance participating in the reaction and is therefore an extensive property, not an intensive property. Thus statements 1, 3 and 4 are correct, making option B the correct answer.

The Standard Hydrogen Electrode (SHE) is the primary reference electrode used in electrochemistry. It consists of a platinum electrode coated with platinum black immersed in a solution containing hydrogen ions of unit activity while hydrogen gas at one atmosphere pressure is bubbled over the electrode. By international convention, the standard electrode potential of SHE is assigned a value of zero volt. NCERT uses SHE as the reference for measuring and tabulating standard reduction potentials of all other electrodes. When an unknown electrode is connected to SHE, the potential difference obtained allows determination of the electrode potential of the unknown half-cell. Because it serves as the fundamental reference point for electrochemical measurements, Pt(s)|H₂(g)|H⁺(aq) is called the Standard Hydrogen Electrode. The other options refer to entirely different electrochemical systems and therefore cannot serve as the universal reference defined by NCERT.

A galvanic cell operates on the basis of a spontaneous redox reaction. According to thermodynamics, a process is spontaneous when the Gibbs free energy change (ΔrG) is negative. NCERT establishes the relationship: ΔrG = −nFEcell where n is the number of electrons transferred, F is the Faraday constant, and Ecell is the cell potential. For a galvanic cell, Ecell is positive because the reaction proceeds spontaneously. Consequently, ΔrG becomes negative. This negative value indicates that the chemical reaction can occur without any external energy supply. During this process, oxidation occurs at the anode and reduction occurs at the cathode. The energy released by the spontaneous redox reaction is converted directly into electrical energy, which produces current in the external circuit. Therefore, a negative ΔrG signifies spontaneous redox energy conversion, making option B correct. This principle forms the foundation of batteries and all galvanic cell operations discussed in NCERT Electrochemistry.

Faraday's First Law of Electrolysis states that the mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte. Mathematically, m ∝ Q, where m is the mass deposited and Q is the electric charge. Since Q = It, the mass deposited is also proportional to current and time. This law forms the quantitative basis of electrochemical calculations in NCERT. The law demonstrates that chemical change occurring at an electrode is directly controlled by the number of electrons transferred through the external circuit. When more charge passes through the electrolyte, more ions are discharged and a larger amount of substance is produced. This principle is used in electroplating, metal purification and industrial electrolysis processes. Therefore the quantity of electricity passed through the electrolyte is the determining factor, making option C the correct answer.

Standard reduction potential indicates the tendency of a species to gain electrons and undergo reduction. According to NCERT electrochemical series, Ag⁺/Ag has a standard reduction potential of +0.80 V, Cu²⁺/Cu has +0.34 V, H⁺/H₂ has 0.00 V, and Na⁺/Na has −2.71 V. A higher positive value corresponds to a greater tendency for reduction. Therefore, silver ions are reduced most readily, followed by copper ions, hydrogen ions, and finally sodium ions. Since sodium possesses a highly negative reduction potential, its reduction is least favourable under standard conditions. This order is important in predicting electrode reactions during electrolysis and galvanic cell operation. NCERT frequently uses standard reduction potentials to determine the direction of electron flow and feasibility of redox reactions. Hence the decreasing order of E° values is Ag⁺/Ag > Cu²⁺/Cu > H⁺/H₂ > Na⁺/Na, corresponding to 3, 4, 1 and 2 respectively.

For the Daniell cell, the Nernst equation is Ecell = E°cell − (0.059/2) log([Zn²⁺]/[Cu²⁺]) When the concentration of Cu²⁺ is increased tenfold, the denominator of the concentration ratio becomes ten times larger. Consequently, log([Zn²⁺]/10[Cu²⁺]) = log([Zn²⁺]/[Cu²⁺]) − 1 Substituting into the Nernst equation gives Ecell = E°cell − (0.059/2)[log([Zn²⁺]/[Cu²⁺]) − 1] = Original Ecell + (0.059/2) Thus the cell potential increases by 0.0295 V, which is equal to (0.059/2) V. This result agrees with NCERT's explanation that increasing the concentration of Cu²⁺ ions favours reduction at the cathode and increases the cell potential. The Nernst equation quantitatively predicts this change. Therefore option C is the correct answer.

The Nernst equation at 298 K is Ecell = E°cell − (0.059/n) log Q where n represents the number of electrons transferred in the balanced overall redox reaction. In the Daniell cell, the oxidation half-reaction is Zn → Zn²⁺ + 2e⁻ and the reduction half-reaction is Cu²⁺ + 2e⁻ → Cu Combining these half-reactions gives the overall reaction Zn + Cu²⁺ → Zn²⁺ + Cu A total of two electrons are transferred during the reaction. Therefore n = 2. This value appears in the denominator of the Nernst equation and determines how strongly concentration changes influence the cell potential. NCERT emphasizes that the value of n must always be obtained from the balanced redox equation. Thus the factor 2 represents the number of electrons involved in the balanced overall reaction, making option B correct.

In a bromine-bromide half-cell, platinum serves merely as an inert conductor that allows electron transfer between the external circuit and the solution. The electrode itself does not participate chemically in the reaction. According to the Nernst equation, the electrode potential depends on the activities or concentrations of the oxidised and reduced forms of the redox couple. For the Br₂/Br⁻ system, the potential depends on the relative concentrations of bromine and bromide ions. Changes in these concentrations alter the reaction quotient and therefore change the electrode potential. Factors such as electrode area, solution volume and distance between electrodes do not directly determine the equilibrium electrode potential. NCERT explains that inert electrodes are used when no solid metallic conductor is naturally present in the redox system. Hence the ratio of concentrations of Br₂ and Br⁻ determines the half-cell potential, making option C correct.

Reduction is defined as the gain of electrons or a decrease in oxidation number. In the given half-reaction, Fe³⁺ gains one electron and is converted into Fe²⁺. The oxidation state decreases from +3 to +2, which clearly indicates reduction. Standard electrode potentials are conventionally written for reduction half-reactions in NCERT and electrochemical tables. The positive value of +0.77 V indicates that Fe³⁺ has a significant tendency to accept electrons under standard conditions. This reaction commonly appears in electrochemical calculations involving iron redox systems. Since electron gain is the defining characteristic of reduction, the given equation represents a reduction half-reaction. Understanding such half-reactions is essential for balancing redox equations and calculating cell potentials. Therefore option D is the correct answer.

According to NCERT, a species having a positive standard reduction potential is reduced more readily than hydrogen ions and therefore acts as a stronger oxidizing agent. Conversely, species having negative reduction potentials tend to lose electrons more readily and can reduce hydrogen ions or be oxidized themselves. The anode is always the electrode at which oxidation takes place, whereas the cathode is the electrode at which reduction occurs. These definitions remain valid for both galvanic and electrolytic cells, although electrode signs may differ. Standard reduction potentials provide a quantitative measure of the tendency of a species to undergo reduction. Therefore, negative E° values correspond to species capable of reducing hydrogen ions or being oxidized more easily, while positive E° values indicate species more readily reduced than hydrogen. Matching these concepts correctly gives 1-b, 2-c, 3-d and 4-a.

NCERT derives the important relationship between equilibrium constant and standard cell potential: E°cell = (2.303RT/nF) log Kc This equation shows that if the standard cell potential is known, the equilibrium constant can be calculated directly. At equilibrium, there is no net tendency for the reaction to proceed in either direction and therefore Ecell becomes zero. Statements 1, 2 and 3 are therefore correct. Statement 4 is incorrect because the value of n is not chosen arbitrarily; it is fixed by the balanced chemical equation. Since Kc depends on the actual reaction stoichiometry, it does not change due to an arbitrary selection of electron number. The relationship between E°cell and Kc remains valid only when n corresponds to the correctly balanced redox equation. Hence option A is correct.

NCERT follows the convention that the anode is written on the left-hand side and the cathode on the right-hand side in cell notation. A single vertical line represents a phase boundary while a double vertical line represents the salt bridge. Since the Standard Hydrogen Electrode acts as the anode, it must appear on the left side as Pt(s)|H₂(g)|H⁺(aq). Copper ions undergo reduction at the cathode and therefore Cu²⁺(aq)|Cu(s) is written on the right side. Platinum is included because it serves as an inert conductor in the hydrogen electrode. Thus the correct notation becomes Pt(s)|H₂(g,1 bar)|H⁺(aq,1 M)||Cu²⁺(aq,1 M)|Cu(s). This notation clearly represents oxidation occurring at the left electrode and reduction occurring at the right electrode.

The salt bridge plays a crucial role in maintaining electrical neutrality in both half-cells of a galvanic cell. During operation of the Daniell cell, zinc metal dissolves into Zn²⁺ ions, causing positive charge accumulation in the anode compartment. Simultaneously, Cu²⁺ ions are reduced at the cathode, decreasing positive charge in that compartment. Without a salt bridge, these charge imbalances rapidly build up. The resulting electrical opposition prevents further movement of electrons through the external circuit. Consequently, the cell reaction stops and the measured cell potential effectively falls to zero. NCERT emphasizes that the salt bridge completes the internal circuit and allows ions to migrate to maintain neutrality. Therefore, in the absence of a salt bridge, sustained current cannot flow and the galvanic cell ceases to function.

During charging of a lead storage battery, an external voltage greater than the battery emf must first be applied. This forces current to flow in the reverse direction of discharge. Electrons are supplied to the negative plate containing lead sulfate deposits. As charging continues, lead sulfate is reduced back to lead at one electrode and oxidized to lead dioxide at the other electrode. These reverse reactions regenerate the original active materials of the battery. Simultaneously, sulfate ions return to the electrolyte, increasing the concentration of sulfuric acid. NCERT describes charging as the reverse of the discharge process, restoring the battery's chemical energy. Therefore the correct chronological sequence is application of external voltage, electron supply to lead sulfate, conversion of lead sulfate into lead and lead dioxide, followed by increase in sulfuric acid concentration.

In a Daniell cell, electrons normally flow spontaneously from zinc to copper because of the cell emf. When an external voltage is applied in the opposite direction, it opposes this spontaneous electron flow. If the external voltage becomes exactly equal to the cell potential, the net potential difference across the circuit becomes zero. Since current is driven by a potential difference, no macroscopic current can flow under these conditions. As a result, oxidation at the zinc electrode and reduction at the copper electrode effectively stop. The system reaches a dynamic state in which there is no net chemical change. NCERT explains that when the opposing potential exactly balances the cell emf, the cell neither performs electrical work nor undergoes net chemical reaction. Therefore no observable current flows and the cell reaction ceases, making option B correct.

The relation ΔrG° = −nFE°cell connects thermodynamics with electrochemistry. Here n is the number of moles of electrons transferred, F is the Faraday constant with unit C mol⁻¹, and E°cell has unit volt (V). Since 1 volt equals 1 joule per coulomb (J C⁻¹), multiplying C mol⁻¹ by J C⁻¹ gives J mol⁻¹. Therefore the unit of ΔrG° becomes joules per mole. NCERT uses this relation to determine whether a reaction is spontaneous and to calculate energy changes associated with electrochemical reactions. Because Gibbs free energy represents the maximum useful work obtainable from a process, its standard unit must be an energy unit. Thus the correct unit is J mol⁻¹.

The Standard Hydrogen Electrode has a standard reduction potential of zero volt. Any redox couple having a positive standard reduction potential possesses a greater tendency to gain electrons than the H⁺/H₂ system. As a result, it acts as a stronger oxidizing agent. Since oxidizing and reducing strengths are inversely related, a species that is easily reduced is a weaker reducing agent. NCERT explains that metals with highly positive reduction potentials do not readily lose electrons and therefore are poor reducing agents. Conversely, species with highly negative reduction potentials readily lose electrons and behave as strong reducing agents. Thus a positive standard reduction potential indicates that the species is more easily reduced than hydrogen ions and therefore functions as a weaker reducing agent than the H⁺/H₂ couple.

In the Nernst equation, E = E° − (RT/nF) ln Q the logarithmic term is dimensionless. Therefore the factor RT/nF must possess the same unit as electrode potential. The gas constant R has unit J mol⁻¹ K⁻¹ and temperature T has unit K. Their product RT therefore has unit J mol⁻¹. The Faraday constant F has unit C mol⁻¹. Dividing RT by F gives (J mol⁻¹)/(C mol⁻¹) = J/C A joule per coulomb is defined as one volt. Hence the factor RT/nF possesses the unit volt. This ensures dimensional consistency throughout the Nernst equation. NCERT frequently uses this unit analysis while deriving the simplified Nernst equation at 298 K.

A hydrogen–oxygen fuel cell generates electricity through the controlled oxidation of hydrogen and reduction of oxygen. In the alkaline fuel cell discussed in NCERT, oxidation occurs at the anode. Hydrogen reacts with hydroxide ions to produce water and electrons according to: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻ The released electrons travel through the external circuit and perform electrical work. At the cathode, oxygen accepts electrons and reacts with water to regenerate hydroxide ions. Since oxidation always occurs at the anode, the hydrogen oxidation reaction represents the anode half-reaction. Fuel cells are highly efficient because they directly convert chemical energy into electrical energy without combustion. NCERT highlights hydrogen–oxygen fuel cells as environmentally friendly energy sources because water is the principal product. Therefore option A correctly represents the anode reaction.