CUET UG Chemistry Booster Test – 2 Electronic configuration

The d-orbitals of transition elements extend significantly towards the outer region of the atom compared with many other orbitals. Because these orbitals project towards the periphery, they are more exposed to surrounding atoms, ions and molecules. As a result, d-electrons are strongly influenced by the chemical environment and participate actively in bonding interactions. This behaviour is responsible for many characteristic properties of transition elements such as complex formation, coloured compounds, catalytic activity and magnetic behaviour. The ability of d-orbitals to interact with surrounding species makes transition metals chemically versatile and distinguishes them from many representative elements.

In transition metals, metallic bonding is strengthened because not only the ns electrons but also several (n−1)d electrons participate in interatomic bonding. Therefore, Statement 1 is correct. The melting points, hardness and metallic character of transition elements depend on the combined contribution of both s and d electrons rather than ns electrons alone, making Statement 2 incorrect. Zinc, Cadmium and Mercury possess ns electrons and therefore Statement 3 is incorrect. Up to manganese, the maximum oxidation state often corresponds to the total number of available s and d valence electrons, making Statement 4 correct. Hence Statements 1 and 4 are correct.

Vanadium (V) has the electronic configuration [Ar]3d³4s² and therefore matches configuration 1. Iron (Fe) possesses [Ar]3d⁶4s² and corresponds to configuration 2. Nickel (Ni) has [Ar]3d⁸4s² and therefore matches configuration 3. Molybdenum (Mo) is an exception in the 4d series and exhibits the configuration [Kr]4d⁵5s¹, corresponding to configuration 4. Correct matching of these configurations requires understanding of electron filling patterns and exceptions among transition elements. Therefore, the correct matching is 1-b, 2-d, 3-a and 4-c.

Palladium exhibits one of the most unusual electronic configurations among transition elements. Instead of following the expected pattern of having one or two electrons in the outermost ns orbital, Palladium possesses the configuration [Kr]4d¹⁰5s⁰. This means the entire valence electron population is accommodated within the 4d subshell while the 5s orbital remains empty. The completely filled d-subshell provides extra stability and explains this exceptional configuration. Because no electrons occupy the 5s orbital, Palladium represents the most extreme deviation from the general electronic configuration of transition elements.

Chromium(II) ions are readily oxidized to Chromium(III) ions in aqueous solution. The conversion can be represented as: Cr²⁺ → Cr³⁺ + e⁻ Since an electron is lost during this process, the reaction is classified as oxidation. Chromium(III) is thermodynamically more stable than Chromium(II) in aqueous medium because of its favourable electronic arrangement and greater hydration stability. The tendency of Cr²⁺ to undergo oxidation is an important characteristic of chromium chemistry and helps explain the predominance of the +3 oxidation state in many compounds.

According to the modern periodic table, Copper belongs to Group 11, Nickel belongs to Group 10, Cobalt belongs to Group 9 and Iron belongs to Group 8. When arranging these elements in decreasing order of group number, the element with the highest group number must be placed first. Therefore, the correct order is Copper (11), Nickel (10), Cobalt (9) and Iron (8). Using the shuffled numbering given in the question, the correct sequence becomes 2, 4, 3 and 1. These elements occur consecutively in the first transition series and illustrate the gradual filling of d-orbitals.

Manganese has the electronic configuration [Ar]3d⁵4s². After losing two 4s electrons, Mn²⁺ attains the configuration [Ar]3d⁵, which is a highly stable half-filled d-subshell arrangement. To remove another electron and form Mn³⁺, one electron must be removed from this stable d⁵ configuration. Because half-filled subshells possess extra stability due to exchange energy and symmetrical electron distribution, significantly more energy is required. As a result, the third ionization enthalpy of manganese is unusually high compared with neighbouring elements.

Zinc possesses the electronic configuration [Ar]3d¹⁰4s². During ion formation, the two 4s electrons are removed first, producing Zn²⁺ with the configuration [Ar]3d¹⁰. Because both the atom and its common ion possess completely filled d-subshells, Zinc does not satisfy the IUPAC definition of a transition element. According to this definition, a transition element must possess an incompletely filled d-subshell in either its atom or one of its common oxidation states. Therefore, Zinc is excluded from the category of true transition metals despite belonging to the d-block.

Palladium possesses the unique electronic configuration [Kr]4d¹⁰5s⁰. Therefore, Statement 1 is correct. Palladium belongs to the second transition series, also known as the 4d series, making Statement 2 correct. Statement 3 is incorrect because Palladium contains no electrons in the 5s orbital. Statement 4 is also incorrect because Palladium does not follow the usual (n−1)d¹⁻¹⁰ns¹⁻² pattern; it is one of the most notable exceptions among transition elements. Thus, only Statements 1 and 2 are correct.

According to the IUPAC definition, a transition element must possess an incompletely filled d-subshell either in its atom or in one of its common oxidation states. Zinc, Cadmium, Mercury and Copernicium have the general electronic configuration (n−1)d¹⁰ns². When they form their most common +2 ions, the ns electrons are removed while the d¹⁰ subshell remains completely filled. Because neither the neutral atoms nor their common ions contain partially filled d-orbitals, these elements are excluded from the category of true transition metals. Their filled d-subshells distinguish them from typical transition elements.

Ionization enthalpy is defined as the energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms. Since it represents an energy change per mole of atoms, its unit is kilojoule per mole (kJ mol⁻¹). In transition elements, ionization enthalpy reflects the ease or difficulty of removing electrons from ns and d orbitals. The value depends on factors such as effective nuclear charge, atomic size, electron-electron repulsion and exchange energy. Understanding ionization enthalpy is essential for explaining oxidation states and reactivity trends among transition metals.

The small energy difference between 3d and 4s orbitals explains the anomalous electronic configurations of Chromium and Copper. Therefore, Condition 1 matches Consequence c. When transition metals form doubly charged ions, the outermost ns electrons are removed before the (n−1)d electrons, so Condition 2 matches Consequence d. The shielding effect among 4f electrons is relatively poor compared to that among d-electrons. This ineffective shielding results in a gradual decrease in atomic and ionic radii across the lanthanoid series, known as lanthanoid contraction. Therefore, Condition 3 matches Consequence a. In the actinoids, the 5f, 6d and 7s orbitals possess comparable energies. As a result, electrons from these orbitals can participate in bonding, producing a wider range of oxidation states. Hence, Condition 4 matches Consequence b. Thus, the correct matching is 1-c, 2-d, 3-a and 4-b.

Chromium possesses the anomalous electronic configuration [Ar]3d⁵4s¹. The five d-electrons occupy separate d-orbitals with parallel spins according to Hund's rule, while the single 4s electron also remains unpaired. Therefore, Chromium contains a total of six unpaired electrons in its ground state. Manganese has the configuration [Ar]3d⁵4s² and therefore contains only five unpaired electrons because the two 4s electrons are paired. Iron and Vanadium possess fewer unpaired electrons. Consequently, Chromium exhibits the maximum number of unpaired electrons among the given elements.

The second transition series extends from Yttrium (Y) to Cadmium (Cd), making Statement 1 correct. Silver possesses the configuration [Kr]4d¹⁰5s¹, so its d-subshell is completely filled, making Statement 2 correct. Molybdenum belongs to the 4d series and exhibits the configuration [Kr]4d⁵5s¹, making Statement 4 correct. Statement 3 is incorrect because the 4d transition metals generally have higher enthalpies of atomization than many corresponding 3d-series elements owing to stronger metallic bonding. Therefore, Statements 1, 2 and 4 are correct.

Enthalpy of atomization reflects the strength of metallic bonding. In the 5d series, Tungsten exhibits one of the highest enthalpies of atomization because of its strong metallic bonding and large number of bonding electrons. Tantalum also possesses a very high value, followed by Platinum. Mercury has a much lower enthalpy of atomization because its filled d¹⁰ configuration results in weak metallic bonding. Therefore, the decreasing order is Tungsten > Tantalum > Platinum > Mercury. Using the shuffled numbering given in the question, the correct sequence becomes 3 > 4 > 1 > 2. This trend illustrates the relationship between electron participation in bonding and metallic bond strength.

Rutherfordium (Rf) is the element with atomic number 104 and is the first element of the 6d transition series after Actinium. It belongs to Group 4 of the periodic table and is classified as a transactinide element. Rutherfordium is produced artificially in nuclear reactions and is highly radioactive. The element derives its name from the physicist Ernest Rutherford, who made significant contributions to atomic structure. Its position in the periodic table corresponds to the beginning of the 6d series, where electrons start occupying the 6d orbitals. Therefore, the element represented by Z = 104 is Rutherfordium.

The first ionization enthalpy depends on effective nuclear charge, atomic radius and electronic configuration. Zinc possesses a stable 3d¹⁰4s² configuration and therefore has the highest first ionization enthalpy among the given elements. Copper follows because of its stable 3d¹⁰4s¹ arrangement. Iron has a lower value than Copper, while Scandium, being near the beginning of the series with a larger atomic size and lower effective nuclear charge, exhibits the lowest ionization enthalpy. Therefore, the decreasing order is Zn > Cu > Fe > Sc.

The conversion of Fe³⁺ to Fe²⁺ involves the gain of one electron: Fe³⁺ + e⁻ → Fe²⁺ Any process involving the gain of electrons is classified as reduction. In this reaction, iodide ions donate electrons and are oxidized to iodine, while Fe³⁺ accepts electrons and is reduced to Fe²⁺. This is a common redox process used in analytical chemistry. Since Fe³⁺ gains an electron during the reaction, the correct classification is reduction.

Mn²⁺ possesses the electronic configuration 3d⁵, which is a half-filled d-subshell. This arrangement provides maximum exchange energy because the five d-electrons occupy separate orbitals with parallel spins. Such a configuration is exceptionally stable. Fe²⁺, on the other hand, has a 3d⁶ configuration. The addition of the sixth electron reduces the exchange-energy advantage present in the d⁵ arrangement. Consequently, Mn²⁺ is more stable than Fe²⁺, making electron removal from Mn²⁺ more difficult and leading to a higher ionization enthalpy. Therefore, the difference arises from the exchange-energy stabilization of the d⁵ configuration.

A d⁵ configuration represents a perfectly half-filled d-subshell. According to Hund's rule, each electron occupies a separate orbital with parallel spin before pairing occurs. This arrangement produces the maximum possible exchange energy among d-electron configurations. Exchange energy lowers the overall energy of the atom or ion and thereby increases stability. As a result, species such as Cr⁺ and Mn²⁺ that possess d⁵ configurations are especially stable. This additional stability explains many irregular trends in ionization enthalpy, oxidation states and redox behaviour observed among transition elements.