CUET UG Chemistry Booster Test – 2 Chemical Properties

The maximum oxidation states of the given transition elements are Mn (+7), Cr (+6), V (+5), and Ti (+4). As we move from titanium toward manganese in the first transition series, the number of electrons available for participation in bonding increases. Consequently, higher oxidation states become possible. Manganese attains the highest oxidation state of +7 in compounds such as permanganate (MnO₄⁻). Chromium reaches +6 in chromates and dichromates, vanadium reaches +5, and titanium commonly exhibits a maximum oxidation state of +4. Therefore, the decreasing order of maximum oxidation state is Mn > Cr > V > Ti, which corresponds to 3, 1, 4, 2.

In the p-block, the inert pair effect causes heavier elements to favour lower oxidation states, making Statement 4 correct. In contrast, among transition elements, higher oxidation states become increasingly stable for heavier members, making Statement 2 correct. For example, Mo(VI) and W(VI) are significantly more stable than Cr(VI), so Statement 3 is correct. Statement 1 is incorrect because dichromate ion containing Cr(VI) in acidic medium is a strong oxidising agent and readily undergoes reduction to Cr(III).

Cr²⁺ possesses a d⁴ configuration and tends to lose an electron to form Cr³⁺ (d³). The d³ configuration is particularly stable in an octahedral field because the t₂g orbitals become half-filled. Therefore, Cr²⁺ acts as a reducing agent. Mn³⁺ also has a d⁴ configuration but tends to gain an electron and form Mn²⁺ (d⁵). The d⁵ configuration is exceptionally stable because it contains a half-filled d-subshell with maximum exchange energy. Consequently, Mn³⁺ behaves as a strong oxidising agent.

In the given reaction, Cu⁺ undergoes simultaneous oxidation and reduction. One Cu⁺ ion loses an electron and becomes Cu²⁺, while another Cu⁺ ion gains an electron and forms metallic copper, Cu. Since the same species acts as both the oxidising agent and reducing agent, the process is called disproportionation. Copper(I) compounds are often unstable in aqueous solution because Cu²⁺ and Cu are thermodynamically more stable than Cu⁺.

Enthalpy of atomisation is the energy required to convert one mole of a solid metal into its gaseous atoms under standard conditions. Since it represents an energy change per mole of substance, the standard unit is kilojoules per mole (kJ mol⁻¹). Transition metals generally possess high enthalpies of atomisation because of strong metallic bonding arising from participation of both ns and d-electrons in bonding.

In the lanthanoid series, electrons are progressively added to the 4f orbitals. These 4f electrons do not effectively shield one another from the increasing nuclear charge. As a result, the effective nuclear attraction experienced by the outer electrons gradually increases across the series. This causes a steady decrease in atomic and ionic radii from one lanthanoid to the next. This gradual decrease in size is known as the lanthanoid contraction. The phenomenon has important consequences for the similarities between the second and third transition series elements.

Across the first transition series, there is a general tendency for standard electrode potentials to become less negative. Therefore, Statement 2 is correct. Zinc exhibits a particularly stable d¹⁰ configuration, which contributes significantly to its characteristic E° value, making Statement 3 correct. The E° values of Mn, Ni and Zn are more negative than expected because of additional stability factors associated with their electronic configurations; hence Statement 4 is correct. Statement 1 is incorrect because the large energy required to convert Cu(s) into Cu²⁺(aq) is not fully compensated by hydration enthalpy, resulting in the positive E° value of copper.

The conversion of copper metal into Cu²⁺ ions involves atomisation and ionisation processes that require a large amount of energy. Although hydration of Cu²⁺ releases energy, this release is insufficient to compensate for the energy consumed during oxidation. Consequently, the overall process is thermodynamically unfavourable, leading to a positive standard electrode potential. Because of this positive E° value, copper does not readily liberate hydrogen from non-oxidising acids such as dilute HCl or dilute H₂SO₄.

Most first-transition-series metals dissolve in mineral acids because they readily undergo oxidation and liberate hydrogen gas. Copper is an exception because it has a positive standard electrode potential (+0.34 V), making oxidation thermodynamically unfavourable under ordinary conditions. Therefore, copper does not dissolve in dilute non-oxidising acids such as hydrochloric acid or dilute sulphuric acid. It reacts only with oxidising acids like concentrated nitric acid.

In acidic medium, permanganate ion (MnO₄⁻) exhibits the greatest oxidising power because manganese in the +7 oxidation state is readily reduced to the highly stable Mn²⁺ state. Dichromate ion (Cr₂O₇²⁻) is also a strong oxidising agent but is weaker than permanganate. Vanadium(V) species such as VO₂⁺ possess moderate oxidising ability. Titanium(IV) species, represented by TiO²⁺, are comparatively stable and show the least tendency to undergo reduction. Therefore, the decreasing order of oxidising power is MnO₄⁻ > Cr₂O₇²⁻ > VO₂⁺ > TiO²⁺, corresponding to 3, 4, 1, 2.

For most compounds of the first transition series, the orbital contribution to the magnetic moment is effectively quenched due to interactions with surrounding ligands. As a result, the magnetic behavior depends primarily on the spin angular momentum of the unpaired electrons. Therefore, the magnetic moment can be calculated using the spin-only formula: μ = √n(n+2) BM where n is the number of unpaired electrons. The calculated values obtained from this formula closely match the experimentally observed magnetic moments for most first-row transition-metal complexes. Hence, the orbital angular momentum contribution is generally neglected.

The spin-only magnetic moment is calculated using: μ = √n(n+2) BM V²⁺ has configuration 3d³ and contains 3 unpaired electrons, giving 3.87 BM. Cr²⁺ has configuration 3d⁴ with 4 unpaired electrons, giving 4.90 BM. Mn²⁺ possesses a half-filled 3d⁵ configuration and has 5 unpaired electrons, resulting in 5.92 BM. Ni²⁺ has configuration 3d⁸ with 2 unpaired electrons and therefore has a magnetic moment of 2.84 BM. Thus, the correct matching is 1-d, 2-c, 3-a and 4-b.

Using the spin-only magnetic moment formula: μ = √n(n+2) For one unpaired electron: μ = √1(1+2) = √3 = 1.73 BM This value represents the magnetic moment arising solely from spin angular momentum. It is commonly observed in d¹ or d⁹ systems having one unpaired electron. Therefore, the correct magnetic moment is 1.73 BM.

Cobalt has atomic number 27 and electronic configuration [Ar] 3d⁷4s². Upon forming Co²⁺, the two 4s electrons are removed, giving [Ar] 3d⁷. In the hydrated ion, three electrons remain unpaired. Applying the spin-only formula: μ = √n(n+2) μ = √3(3+2) μ = √15 μ = 3.87 BM Therefore, the spin-only magnetic moment of hydrated Co²⁺ is 3.87 BM.

The colour of transition-metal complexes arises from d-d electronic transitions. The energy required for these transitions depends on the splitting of d orbitals in the ligand field. Different ligands produce different magnitudes of crystal-field splitting. Consequently, the frequency of visible light absorbed changes according to the ligand attached to the metal ion. Therefore, the nature of the ligand is the primary factor determining the frequency of absorbed light and the resulting colour of the complex.

When a transition-metal complex absorbs visible light, an electron is promoted from a lower-energy d orbital to a higher-energy d orbital. The colour corresponding to the absorbed wavelength is removed from the transmitted or reflected light. As a result, the colour perceived by the observer is the complementary colour of the absorbed light. For example, if red light is absorbed, the complex may appear green. The exact wavelength absorbed depends upon the ligand field splitting energy, which is influenced by the nature of the ligand. This principle explains the characteristic colours of many transition-metal ions and complexes.

Ligands such as carbon monoxide possess both σ-donor and π-acceptor properties. They donate electron density to the metal through σ-bonding and simultaneously accept electron density from filled metal d orbitals through π-back bonding. This synergic bonding stabilizes metals in low oxidation states, including oxidation state zero. Nickel tetracarbonyl, Ni(CO)₄, is a classic example where the metal remains in the zero oxidation state because of strong metal-to-ligand back donation. Such stabilization is an important feature of transition-metal carbonyl chemistry.

In strongly alkaline media, iron can attain the oxidation state +6 and form the oxoanion [FeO₄]²⁻. The oxidation number of iron is calculated as follows: let the oxidation state of iron be x. Since oxygen contributes −8 and the overall charge is −2, x − 8 = −2, giving x = +6. Therefore, the ion is named ferrate(VI). This species demonstrates the ability of transition metals to exhibit high oxidation states and form oxoanions similar to chromate and permanganate ions.

Transition metals are widely used as catalysts because of their variable oxidation states and ability to form intermediate complexes. Vanadium(V) oxide is used as a catalyst in the Contact Process for manufacturing sulphuric acid, making Statement 3 correct. Finely divided iron serves as the catalyst in Haber's Process for ammonia synthesis, making Statement 4 correct. Nickel is widely used in catalytic hydrogenation of vegetable oils, making Statement 1 correct. Statement 2 is incorrect because PdCl₂ is associated with the Wacker process rather than the Ziegler process. Therefore, Statements 1, 3 and 4 are correct.

Palladium(II) chloride is used in the Wacker-type oxidation process, where unsaturated hydrocarbons are oxidized in the presence of palladium catalysts. A well-known example mentioned in NCERT is the oxidation of ethyne to ethanal. The catalyst cycles between different oxidation states during the reaction and facilitates electron transfer. This catalytic application demonstrates the importance of transition metals in industrial oxidation processes.