D And F Block Elements – NEET UG ChemistryPractice Questions & PYQs

The structural geometry of transition metal oxoanions is influenced by multiple bonding that stabilizes the central atom alongside sigma framework requirements. In species like manganate and permanganate, this multiple bonding arises from a lateral overlap between the filled p-orbitals of oxygen and the vacant d-orbitals of the manganese center. This specific $d\pi-p\pi$ interaction reinforces the tetrahedral arrangement by minimizing electron pair repulsions and satisfying directional bonding constraints.

The spin-only magnetic moment of a transition metal ion is determined by the formula $\mu = \sqrt{n(n+2)}$ B.M., where $n$ denotes the number of unpaired $d$-electrons. Cobalt(III) with a $3d^6$ electronic configuration has four unpaired electrons, yielding a magnetic moment of $\sqrt{24}$ B.M. Chromium(III) exhibits a $3d^3$ configuration with three unpaired electrons, corresponding to $\sqrt{15}$ B.M. Iron(III) contains a half-filled $3d^5$ configuration with five unpaired electrons, resulting in $\sqrt{35}$ B.M., whereas nickel(II) has a $3d^8$ configuration with two unpaired electrons, matching $\sqrt{8}$ B.M.

The presence of unpaired electrons dictates both the magnetic behavior and the possibility of electronic transitions within the d-orbitals of transition metal species. Evaluating the electronic configurations shows that manganate possesses a $3d^1$ configuration with one unpaired electron, whereas the other listed ions contain fully emptied d-shells with zero unpaired electrons. This single unpaired electron enables both net paramagnetism and characteristic d-d electronic transitions under light absorption.

When mercuric chloride is treated with aqueous iodide ions, it initially precipitates as mercuric iodide, which subsequently dissolves in excess iodide to form a soluble tetraiodomercurate(II) complex anion. Simultaneously, molecular iodine combines with iodide ions in the solution to produce triiodide ions. Consequently, the coexistence of these species in the presence of excess iodide yields both tetraiodomercurate(II) and triiodide ions as the final dissolved products.

Sulfur dioxide acts as a strong reducing agent due to the presence of sulfur in a plus four oxidation state, allowing it to reduce intensely colored permanganate ions. When bubbled through an acidified potassium permanganate solution, it reduces the purple manganese seven ions into colorless manganese two ions while being oxidized to sulfate, thereby completely discharging the color.

Actinoids exhibit a broader spectrum of oxidation states compared to lanthanoids because their 5f, 6d, and 7s subshells possess exceptionally comparable energy levels. Due to this minimal energy gap, electrons residing in all three valence levels can be readily utilized for chemical bonding. This close proximity in energy allows multiple electrons to participate flexibly under varying chemical environments, leading to numerous stable oxidation states.

Acidified potassium dichromate acts as a strong oxidizing agent that converts $\mathrm{SO}_2$ into sulfate while itself being reduced to chromium(III) ions. During this redox reaction, the orange color of the dichromate solution changes as green $\mathrm{Cr}_2(\mathrm{SO}_4)_3$ is produced. The sulfur atom in sulfur dioxide increases its oxidation state from +4 to +6, while chromium is reduced from +6 to +3.

Lanthanoids feature deeply buried 4f orbitals that remain shielded from external ligands, rendering their chemical reactivity distinct from and generally lower than typical reactive metals rather than universally exceeding that of aluminum. Conversely, the other options accurately reflect established lanthanoid chemistry, such as cerium in the plus four oxidation state acting as a strong oxidant, europium readily exhibiting a stable plus two oxidation state, and basicity systematically dropping across the series as ionic radii decrease from praseodymium to lutetium.

Lanthanide electronic configurations are governed by the extra stability associated with half-filled $f$-subshells. Europium achieves a stable half-filled configuration with seven electrons in the $4f$ orbital and a completely filled $6s$ orbital. Gadolinium continues by placing the next electron into the $5d$ orbital rather than disrupting the half-filled $4f^7$ shell, yielding a $4f^7 5d^1 6s^2$ configuration. Terbium then adds electrons directly back into the $4f$ orbital since the half-filled stability has been passed, resulting in $4f^9$ alongside the filled $6s$ subshell.

The electron filling of lanthanide elements prioritizes exchange energy stabilization, particularly when achieving half-filled subshell configurations. For gadolinium with an atomic number of 64, filling the seven degenerate $4f$ orbitals creates a symmetrical, half-filled shell that provides exceptional quantum mechanical stability. Rather than forcing an eighth electron into this already half-filled orbital set, the next available electron enters the $5d$ orbital. Combined with the fully occupied $6s$ orbitals, this produces the characteristic configuration of xenon followed by $4f^7 5d^1 6s^2$.

The volume or moles of an oxidizing agent required depends directly on the total number of electrons lost per mole of the reducing agent. In $\mathrm{FeSO}_4$, only the iron(II) ion undergoes oxidation to iron(III), releasing a single electron per formula unit. Conversely, compounds like $\mathrm{FeSO}_3$, $\mathrm{FeC}_2\mathrm{O}_4$, and $\mathrm{Fe}(\mathrm{NO}_2)_2$ contain anions that are also oxidizable, substantially increasing the total reducing capacity per mole. Since $\mathrm{FeSO}_4$ possesses the lowest oxidizable electron yield among the choices, it consumes the minimum quantity of acidified $\mathrm{KMnO}_4$.

Lanthanoid contraction is the steady decrease in atomic and ionic radii across the lanthanoid series, which progressively increases the effective nuclear charge. Because of this cumulative size reduction prior to the third transition series, elements in the second and third transition rows exhibit nearly identical atomic dimensions. Zirconium and hafnium are prime examples of this phenomenon, residing in group four where the expected increase in size down the group is completely offset by the intervening lanthanoids.

Spin-only magnetic moment is calculated using the formula $\mu = \sqrt{n(n+2)}$ Bohr Magnetons, where $n$ represents the number of unpaired electrons. A magnetic moment value of 2.84 B.M. corresponds to two unpaired electrons, since $\sqrt{2(2+2)} = \sqrt{8}$ evaluates to approximately 2.83. Nickel with atomic number 28 has an outer electronic configuration of $3d^8 4s^2$, meaning the $\mathrm{Ni}^{2+}$ ion retains a $3d^8$ configuration. Distributing eight electrons among the five degenerate $3d$ orbitals leaves precisely two unpaired electrons.

Oxidation involves an increase in the oxidation state of an element due to the loss of electrons. In the formation of $\mathrm{Fe}(\mathrm{CO})_5$, elemental iron reacts with carbon monoxide to form a neutral coordination complex where iron retains its zero oxidation state. Conversely, displacement reactions involving $\mathrm{CuSO}_4$, steam reactions yielding hydrogen, and sheet rusting all force iron to lose electrons and increase its positive oxidation state.

Permanganate ion acts as a powerful oxidizing agent in acidic medium, oxidizing hydrogen peroxide into molecular oxygen while itself undergoing reduction. During this redox process, the manganese in $\mathrm{KMnO}_4$ changes from a high oxidation state of plus seven down to plus two as manganese sulfate. Simultaneously, the peroxide species loses electrons to form oxygen gas, releasing protons and water as byproducts in the sulfuric acid environment.

The spin-only magnetic moment is calculated using the formula $\mu = \sqrt{n(n+2)}$ BM, where $n$ represents the number of unpaired electrons. Matching a magnetic moment of 2.83 BM to this relationship indicates that $n = 2$. Evaluating the electronic configuration of nickel as $\mathrm{Ni}^{2+}$ results in a $3d^8$ valence state, which distributes eight electrons across the five $d$ orbitals to leave precisely two unpaired electrons.

Lanthanoid contraction arises from the diffuse and complex spatial shape of inner orbitals that fail to effectively shield outer valence shells from the nucleus. As atomic number increases across the lanthanide series, electrons are successively added to this buried sublevel with very poor shielding characteristics. Consequently, the outer electrons experience a steadily rising net attractive pull toward the center, causing a progressive shrinkage in atomic and ionic radii.

A transition element is formally defined as an element that has an incompletely filled $d$ subshell either in its ground state or in its common oxidation states. Scandium features a partially filled $3d^1$ configuration in its ground state, qualifying it as a true transition metal. Conversely, zinc possesses a completely filled $3d^{10}$ configuration in both its ground state and its stable $+2$ oxidation state, which excludes it from being classified as a transition element.

Interstitial compounds are formed when small atoms like hydrogen, carbon, or nitrogen become trapped inside the crystal lattices of transition metals, imparting exceptional hardness and high melting points. Because of their tightly bound structures, these compounds retain metallic conductivity while remaining chemically inert. Asserting that they are chemically reactive contradicts their inherent stability and resistance to chemical attack.

Magnetic properties of lanthanoid ions are determined by the presence or absence of unpaired electrons within their $4f$ orbitals. Evaluating the electronic configurations shows that $\mathrm{Yb}^{2+}$ possesses a completely filled $4f^{14}$ subshell with zero unpaired electrons, making it diamagnetic. In contrast, the other listed ions retain unpaired electrons in their respective subshells, rendering them paramagnetic.