CUET UG Chemistry Booster Test – 2 Basics of Electrochemistry

The Standard Hydrogen Electrode (SHE) serves as the universal reference electrode in electrochemistry. According to NCERT, it consists of a platinum electrode coated with finely divided platinum black, immersed in a solution where the hydrogen ion concentration corresponds to standard conditions. Pure hydrogen gas is continuously bubbled around the electrode at a pressure of 1 bar. The electrode potential of SHE is assigned a value of zero volt by convention at all temperatures, providing a reference point for measuring standard electrode potentials of other half-cells. The SHE is not restricted to functioning as an electrolytic cell. Depending on the half-cell connected to it, it may act either as an anode or as a cathode. Therefore statements 1, 2 and 3 are correct, whereas statement 4 is incorrect.

NCERT explains that the absolute electrode potential of a single half-cell cannot be measured directly because only potential differences are experimentally observable. Therefore, the standard electrode potential of a half-cell is determined by connecting it to the Standard Hydrogen Electrode under standard conditions and measuring the cell potential. Since the electrode potential of SHE is defined as zero volt, the measured cell potential directly gives the standard electrode potential of the other half-cell. This method provides a common reference scale for comparing oxidation and reduction tendencies of various electrodes. Therefore option B is correct.

Conductivity is defined as the reciprocal of resistivity. It indicates the ability of an electrolytic solution to conduct electric current. According to NCERT, the SI unit of conductivity is Siemens per metre . Since conductivity is the reciprocal of resistance per unit length and area, its unit is derived accordingly. Higher conductivity indicates greater ionic mobility and concentration within the solution. Conductivity measurements are widely used to study electrolyte solutions and determine ionic properties. Therefore option B is correct.

The given reaction is the overall cell reaction of the Daniell cell. Zinc loses electrons and undergoes oxidation: Copper ions gain electrons and undergo reduction: Since the cell potential is positive , the reaction is spontaneous. The reaction also involves displacement because zinc replaces copper from its ionic form. Thus it is a spontaneous redox displacement reaction. The positive cell potential confirms that electrical energy can be produced from this reaction without external assistance. Therefore option C is correct.

According to NCERT, the maximum useful electrical work obtainable from an electrochemical cell is achieved when the cell operates reversibly. Under reversible conditions, the system remains extremely close to equilibrium throughout the process. The relationship connects Gibbs free energy change with the maximum reversible electrical work. Any irreversible operation introduces energy losses due to resistance and other dissipative effects, reducing the useful work obtained. Therefore, to achieve the maximum possible electrical work from a galvanic cell, charge must be transferred reversibly. Hence option A is correct.

During the electrolysis of aqueous sodium chloride (brine), chloride ions migrate towards the anode. Since oxidation takes place at the anode, chloride ions lose electrons according to the reaction: As a result, chlorine gas is liberated at the anode. At the cathode, water is reduced to produce hydrogen gas and hydroxide ions. The overall process is industrially important because it simultaneously produces chlorine gas, hydrogen gas and sodium hydroxide. This electrolysis process is widely used in the chlor-alkali industry. Therefore, the gas liberated at the anode is chlorine gas and option C is correct.

The oxidising power of a substance is directly related to its tendency to gain electrons. A species with a higher standard reduction potential is a stronger oxidising agent because it more readily undergoes reduction. The given standard reduction potentials are: F₂ = +2.87 V Cl₂ = +1.36 V Br₂ = +1.09 V I₂ = +0.54 V Since oxidising power increases with increasing reduction potential, fluorine is the strongest oxidising agent, followed by chlorine, bromine and iodine. Therefore the decreasing order of oxidising power is: F₂ > Cl₂ > Br₂ > I₂ Using the numbering given in the question: 3 > 1 > 4 > 2 Hence option D is correct.

In the Daniell cell, zinc acts as the anode where oxidation occurs: Zn(s) → Zn²⁺ + 2e⁻ Copper ions are reduced at the cathode: Cu²⁺ + 2e⁻ → Cu(s) According to standard electrochemical notation, the anode is represented on the left side and the cathode on the right side. Therefore the correct matching is: 1-b, 2-c, 3-d, 4-a This notation helps identify the direction of electron flow and the nature of each half-cell in an electrochemical system.

The Daniell cell operates spontaneously because its cell potential is approximately 1.1 V. When an external opposing potential equal to the cell emf is applied, electron flow stops because the net driving force becomes zero. If the opposing potential exceeds 1.1 V, the spontaneous reaction is forced to run in the reverse direction. Under these conditions, the system behaves as an electrolytic cell. Copper metal begins to oxidize, while zinc ions are reduced, reversing the normal Daniell cell reaction. Thus electrical energy is used to drive a non-spontaneous chemical transformation. Therefore option C is correct.

The Nernst Equation is one of the most important relationships in electrochemistry. It describes how electrode potential and cell potential vary with concentration and temperature. The general form is: where is the gas constant, is the temperature, is the number of electrons transferred and is the Faraday constant. At 298 K, the equation simplifies to: This equation clearly shows that cell potential depends on concentration and is not constant under all conditions. Therefore statements 1, 3 and 4 are correct, while statement 2 is incorrect.

Aluminium is extracted industrially by the Hall-Héroult process. In this process, purified aluminium oxide (Al₂O₃) is dissolved in molten cryolite (Na₃AlF₆). Cryolite serves two important functions. First, it lowers the melting point of aluminium oxide from about 2323 K to around 1140 K, making the process economically feasible. Second, it increases the electrical conductivity of the molten electrolyte. During electrolysis, aluminium ions are reduced at the cathode to produce molten aluminium metal. Therefore aluminium oxide is electrolysed in the presence of cryolite.

During electrolysis of aqueous sodium chloride (brine), water is preferentially reduced at the cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻ Thus hydrogen gas is liberated. At the anode, chloride ions are oxidized: 2Cl⁻ → Cl₂ + 2e⁻ producing chlorine gas. The sodium ions remain in solution and combine with hydroxide ions to form sodium hydroxide. Sodium metal is not deposited because water is reduced more readily than sodium ions. Therefore statements 1, 2 and 3 are correct, while statement 4 is incorrect.

Biological communication occurs through electrochemical processes involving ions such as Na⁺, K⁺, Ca²⁺ and Cl⁻. Cell membranes maintain differences in ion concentration, producing charge separation across the membrane. The movement of these ions generates electrical signals that travel through neurons and between cells. These processes form the basis of sensory perception, muscle contraction and nervous coordination. Therefore electrochemical communication fundamentally depends on charge separation and ion transport.

Electrochemistry is not limited to chemical industries or laboratory studies. It contributes to understanding biological communication, batteries, corrosion, fuel cells, metallurgy and environmental technologies. Since it connects concepts from chemistry, physics, biology, engineering and materials science, electrochemistry is considered a highly interdisciplinary subject. The electrochemical transmission of sensory signals demonstrates its importance beyond traditional chemistry and highlights its role in understanding living systems.

Electrochemical technologies such as fuel cells, advanced batteries and electrolytic production methods are actively researched because of their potential to improve energy efficiency and reduce environmental impact. Many electrochemical processes convert energy directly without multiple intermediate stages, thereby minimizing energy losses. They can also reduce pollutant emissions compared with conventional fossil-fuel-based technologies. These advantages make electrochemical systems attractive for sustainable energy generation, energy storage and environmentally friendly industrial production.

The hydrogen-oxygen fuel cell operates through the electrochemical reaction between hydrogen and oxygen to produce electrical energy. The overall reaction is: 2H₂(g) + O₂(g) → 2H₂O(l) Unlike thermal power plants that burn fossil fuels and release pollutants such as carbon dioxide, sulfur oxides and nitrogen oxides, the hydrogen-oxygen fuel cell produces water as its primary byproduct. This makes the technology environmentally friendly and highly attractive for sustainable energy applications. The clean production of water, combined with high energy efficiency, is one of the major reasons fuel cells are considered an important future energy technology.

The passage explains that thermal power plants first convert the chemical energy of fuel into heat, then into steam, mechanical energy and finally electrical energy. Each conversion stage results in energy loss. Fuel cells differ fundamentally because they convert the chemical energy of fuel directly into electrical energy through electrochemical reactions. Since intermediate stages are avoided, fuel cells operate with higher efficiency and lower pollution. This direct conversion process is the defining feature that distinguishes fuel cells from thermal power plants.

The hydrogen-oxygen fuel cell used in the Apollo space programme generated electricity through the reaction of hydrogen with oxygen. The reaction produced water as the only significant byproduct. This water condensed into liquid form and was utilized by astronauts as drinking water. Thus the fuel cell not only supplied electrical power for spacecraft operations but also provided a valuable source of water during the mission. This dual benefit made hydrogen-oxygen fuel cells particularly suitable for space applications.

One of the major theoretical contributions of electrochemistry is its connection with chemical thermodynamics. The standard cell potential (E°cell) can be related to the Gibbs free energy change and subsequently to the equilibrium constant of a reaction. Since equilibrium constants are often difficult to determine directly, electrochemical measurements provide an alternative and highly accurate method. This demonstrates that electrochemistry is not merely an industrial technology but also a powerful theoretical tool for understanding chemical behavior. Therefore the correct answer is broad theoretical importance.

Corrosion protection by sacrificial electrodes is based on electrochemical principles. Metals such as magnesium and zinc have more negative electrode potentials than iron and therefore oxidize more readily. When connected to iron, the sacrificial metal acts as the anode and undergoes oxidation preferentially: Zn → Zn²⁺ + 2e⁻ The electrons released prevent oxidation of the iron surface. As a result, iron remains protected while the sacrificial metal gradually corrodes. Since the oxidation of the sacrificial metal occurs naturally due to its higher reactivity, the protecting reaction is a spontaneous preferential oxidation process.