CUET UG Chemistry Booster Test – 2 Electrochemical Cells

Electrochemical cells play an important role in both energy conversion and analytical chemistry. Galvanic cells convert the Gibbs free energy of spontaneous redox reactions into useful electrical work. Electrolytic cells consume electrical energy and store it as chemical energy through non-spontaneous reactions. Electrochemical methods are also used for determining pH, electrode potentials and ionic concentrations. Modern electrochemical technologies, including batteries, fuel cells and electrochemical sensors, are considered relatively energy efficient and environmentally friendly compared with many conventional industrial processes. Their applications extend from portable electronics to pollution control and green energy systems. Therefore all four statements are correct and option D is the correct answer.

The Daniell cell possesses a standard emf of approximately 1.1 V. When no opposing voltage is applied, the full cell potential drives electron flow and the current is maximum. As an external opposing voltage is gradually increased, the effective potential difference available for electron flow decreases. Consequently, the current decreases continuously. At an opposing voltage of 1.1 V, the external potential exactly balances the cell emf, causing the net driving force to become zero and the current to stop. The given values are: 1. 1.0 V 2. 0.0 V 3. 1.1 V 4. 0.5 V Therefore the decreasing order of current magnitude is: 0.0 V > 0.5 V > 1.0 V > 1.1 V Using the numbering given: 2 > 4 > 1 > 3 Hence option A is correct.

The maximum electrical work obtainable from a galvanic cell is directly related to the Gibbs free energy change of the cell reaction. According to NCERT, where is the number of electrons transferred and is the Faraday constant. A spontaneous redox reaction has a negative Gibbs energy change, and this decrease in Gibbs energy becomes available as useful electrical work. The greater the magnitude of negative Gibbs energy, the larger the electrical work that can be obtained. Therefore electrical energy produced by a galvanic cell corresponds directly to the decrease in Gibbs energy of the spontaneous reaction.

A galvanic cell operates through a spontaneous redox reaction in which oxidation and reduction occur in separate half-cells. The separation of these reactions forces electrons to travel through an external circuit, thereby generating electrical current. This coupling of spontaneous oxidation and reduction processes is known as spontaneous redox coupling. The Daniell cell provides a classic example where zinc undergoes oxidation and copper ions undergo reduction. Because the reaction is spontaneous, chemical energy is converted into electrical energy without external assistance. Therefore option C is correct.

The fundamental distinction between an electrolytic cell and a galvanic cell is the direction of energy conversion. A galvanic cell generates electrical energy from a spontaneous chemical reaction. In contrast, an electrolytic cell uses an external source of electrical energy to force a non-spontaneous reaction to occur. For example, when an opposing voltage greater than 1.1 V is applied to a Daniell cell, the spontaneous reaction reverses and the system behaves as an electrolytic cell. Electrical energy is then converted into chemical energy. Therefore option B correctly describes the distinguishing feature of an electrolytic cell.

In a normal Daniell cell, zinc acts as the anode and copper acts as the cathode. However, when an external potential greater than 1.1 V is applied, the spontaneous reaction is reversed. Copper metal is oxidized: and zinc ions are reduced: According to IUPAC convention, the anode is written on the left and the cathode on the right. Since copper now acts as the anode and zinc acts as the cathode, the correct cell notation becomes: Therefore option A is correct.

At the electrode-electrolyte interface, two opposing processes occur. Metal ions from the solution may deposit on the electrode, making it relatively positive. Simultaneously, metal atoms of the electrode may enter the solution as positive ions, leaving their electrons behind on the metal surface. When the tendency of metal atoms to ionize is greater, excess electrons accumulate on the electrode and it becomes negatively charged relative to the solution. This charge separation leads to the development of an electrode potential. In the zinc electrode of the Daniell cell, zinc atoms readily enter the solution as ions, leaving electrons behind and producing a negative electrode potential. Therefore option D is correct.

When a metal electrode is immersed in a solution containing its ions, two competing processes occur: metal atoms may enter the solution as ions, and metal ions from the solution may deposit on the electrode. Eventually, equilibrium is established and a separation of charges develops between the electrode and the electrolyte. This charge separation produces a potential difference known as the electrode potential. The magnitude and sign of this potential depend on the relative tendencies of oxidation and reduction at the interface. Electrode potentials form the basis for measuring standard reduction potentials and calculating cell emf. Therefore option C is correct.

The standard reduction potentials given in NCERT are: Therefore the correct matching is: 1-c, 2-a, 3-d, 4-b which corresponds to option A.

Some electrochemical half-cells involve only ions, gases or dissolved species and do not contain a solid conducting material. In such cases, an inert electrode is required to provide a conducting surface for electron transfer. Platinum is commonly used because it is chemically stable, resistant to corrosion and an excellent conductor of electricity. In the bromine-bromide half-cell, platinum does not participate in the chemical reaction. Instead, it merely allows electrons to enter or leave the system and provides a surface on which the oxidation or reduction process can occur. Therefore platinum functions as an inert electrode and option C is correct.

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 concentration 1 M, while hydrogen gas is bubbled at a pressure of 1 bar. By international convention, the electrode potential of SHE is assigned a value of exactly 0.00 V. This allows the standard electrode potentials of all other half-cells to be measured relative to it. Because absolute electrode potentials cannot be measured directly, the SHE serves as the universal standard reference electrode in electrochemical measurements.

According to IUPAC convention, the anode is written on the left side and the cathode on the right side in cell notation. In a galvanic cell, oxidation occurs spontaneously at the anode, releasing electrons. These electrons travel through the external circuit and are consumed at the cathode. Therefore reduction occurs spontaneously at the cathode. For example, in the Daniell cell: Cu²⁺(aq) + 2e⁻ → Cu(s) This reduction half-reaction takes place at the right-hand side half-cell. Hence the correct answer is spontaneous reduction.

In standard electrochemical cell notation, a single vertical line (|) represents a phase boundary, such as between a metal electrode and its ionic solution. A double vertical line (||) represents the salt bridge separating the two half-cells. According to IUPAC convention, the anode is written on the left and the cathode on the right. This arrangement is chosen so that the cell emf is positive for a spontaneous galvanic cell. Therefore statements 1, 2 and 4 are correct, while statement 3 is incorrect because it reverses the accepted convention.

A species with a higher standard reduction potential has a greater tendency to gain electrons and undergo reduction. Therefore it preferentially acts as the cathode when paired with another half-cell. The approximate reduction potentials are: Au³⁺/Au = +1.50 V Ag⁺/Ag = +0.80 V Zn²⁺/Zn = –0.76 V Li⁺/Li = –3.05 V Arranging these in decreasing order gives: Au³⁺/Au > Ag⁺/Ag > Zn²⁺/Zn > Li⁺/Li Using the numbering given: 3 > 1 > 2 > 4 Hence option B is correct.

The Standard Hydrogen Electrode has a potential of 0.00 V. If a half-cell possesses a negative standard reduction potential, it means that its tendency to undergo reduction is lower than that of the H⁺/H₂ couple. Consequently, when connected to SHE, the half-cell tends to undergo oxidation while hydrogen ions undergo reduction. In thermodynamic terms, hydrogen gas is relatively more stable than the reduced form of the species. Metals such as zinc and lithium have negative standard reduction potentials and therefore act as strong reducing agents.

According to standard galvanic cell convention, the left electrode represents the anode where oxidation occurs, while the right electrode represents the cathode where reduction occurs. Electrons are released at the anode and travel through the external circuit toward the cathode. Conventional current is defined opposite to electron flow. Therefore electrons move from the negative electrode to the positive electrode, while current flows from the positive electrode to the negative electrode. Thus the correct matching is: 1-c, 2-b, 3-a, 4-d

The Faraday constant (F) is defined as the total electric charge carried by one mole of electrons. It is numerically equal to the product of Avogadro's number and the electronic charge: F = Nₐ × e ≈ 6.022 × 10²³ × 1.602 × 10⁻¹⁹ ≈ 96485 C mol⁻¹ Since it represents electric charge associated with one mole of electrons, its SI unit is coulomb per mole (C mol⁻¹). The Faraday constant plays an important role in Faraday's laws of electrolysis and electrochemical calculations.

An individual electrode potential cannot be measured directly because electrical measurements always involve a difference in potential between two points. A single half-cell by itself cannot sustain a measurable potential difference without being connected to another half-cell. Therefore electrode potentials are always determined relative to a reference electrode, most commonly the Standard Hydrogen Electrode (SHE). By constructing a complete electrochemical cell and measuring the cell emf, the potential of the unknown half-cell can be calculated. Hence a complete circuit linking two half-cells is essential.

The Nernst equation is: E = E° − (RT/nF) ln Q where R is the universal gas constant. It relates cell potential to concentration and temperature. The SI unit of R is joule per kelvin per mole (J K⁻¹ mol⁻¹), and its value is approximately 8.314 J K⁻¹ mol⁻¹. This constant appears throughout thermodynamics and electrochemistry because it connects energy, temperature and amount of substance.

Fluorine possesses the highest standard reduction potential (+2.87 V), making it the strongest oxidising agent and the species with the greatest tendency to undergo reduction. Lithium metal has a highly negative standard reduction potential (−3.05 V), making it an exceptionally strong reducing agent. Standard electrode potentials are also related to Gibbs free energy and equilibrium constants through thermodynamic relationships. Statement 3 is incorrect because a negative E° indicates that the species is a weaker oxidising agent than the H⁺/H₂ couple, not a weaker reducing agent. Therefore statements 1, 2 and 4 are correct.