Classification Of Elements – JEE Main ChemistryPractice Questions & PYQs

The sequence ${Al}^{3+}< {Mg}^{2+}< {Na}^{+}< {F}^{-}$ involves isoelectronic species where the ionic radius decreases as nuclear charge increases, effectively matching (A) to (IV). In (B), the ionization enthalpy order is dictated by the stable half-filled $2p$ subshell of nitrogen, which pushes it to a higher value than oxygen, linking (B) to (I). The trend in (C) reflects metallic character, which rises moving down a group and falls moving across a period, while the electronegativity increase observed in (D) directly aligns with the progression across the third period, identifying (D) with (III).

The enthalpy of atomization of transition metals is primarily determined by the strength of metallic bonding, which depends on the number of electrons available for interaction. Among the specified metals, $\mathrm{Cr}$ possesses the lowest enthalpy of atomization in comparison to $\mathrm{Fe}$, $\mathrm{Co}$, and $\mathrm{Ni}$. Given the electronic configuration of chromium is $3d^5 4s^1$, the number of valence electrons present in the metal is 6.

Electronegativity generally increases as you move from left to right across a period in the periodic table because the increasing effective nuclear charge pulls bonding electrons more tightly toward the nucleus. By observing the valence shell electron configurations, we can identify these as second-period elements: $1s^2 2s^2 2p^2$ is Carbon (D), $1s^2 2s^2 2p^3$ is Nitrogen (A), $1s^2 2s^2 2p^4$ is Oxygen (B), and $1s^2 2s^2 2p^5$ is Fluorine (C). Arranging these in increasing order of their ability to attract shared electrons results in Carbon having the lowest electronegativity and Fluorine the highest, establishing the sequence $D < A < B < C$.

For isoelectronic species, the ionic radius is inversely related to the nuclear charge, meaning that as the number of protons decreases for a constant number of electrons, the attractive pull on the electron cloud weakens and the size increases, which validates the order $\text{Mg}^{2+} < \text{Na}^{+} < \text{F}^{-} < \text{O}^{2-}$. In terms of electron gain enthalpy, chlorine possesses a higher magnitude than fluorine because the compact $2p$ orbitals of fluorine create significant inter-electronic repulsion when an additional electron is introduced, leading to the verified sequence $\text{Cl} > \text{F} > \text{Br} > \text{I}$.

Phosphorus serves as the essential building block of the apatite family of minerals. When considering the periodic trends of group 15 elements, the first ionization enthalpy follows the order $\mathrm{N} > \mathrm{P} > \mathrm{As} > \mathrm{Sb}$ because the energy required to remove an electron decreases as the atomic radius increases down the group. Based on this established periodic trend, the first ionization enthalpy specific to phosphorus is 1012 kJ mol⁻¹.

The atomic radius decreases as you move from left to right across a period in the periodic table, which identifies Boron as the element with the smallest atomic radius among those listed. Because Boron is a Group 13 element, it forms an oxide with the chemical formula $\mathrm{B}_2\mathrm{O}_3$, corresponding to the general type $\mathrm{A}_2\mathrm{O}_3$.

The magnitude of electron gain enthalpy is primarily determined by the interplay between nuclear charge and inter-electronic repulsion, which results in specific anomalies for the first member of a group. While one might expect the first element to always have the most negative value, chlorine actually exhibits a more negative electron gain enthalpy than fluorine due to the excessive inter-electronic repulsion within the small atomic volume of fluorine, rendering Statement I incorrect. Meanwhile, within the oxygen family, oxygen possesses the least negative electron gain enthalpy relative to its heavier group members, making Statement II factually accurate.

The size of an ion is determined by the balance between the nuclear charge and the number of electrons; removing an electron to form a cation increases the effective nuclear charge per electron, drawing the electron cloud inward, while adding an electron to form an anion creates greater inter-electronic repulsion, which causes the cloud to expand. Applying this to sodium, the metallic radius of Na is 1.86 Å, and forming the Na+ cation reduces this radius, which confirms that Statement I is correct. Statement II is incorrect because anions are consistently larger than their corresponding neutral atoms, meaning the generalization that all ions are smaller than their parent elements is false.

The acidic strength of hydrides relies on the ease of proton release, which is governed by the strength of the bond between the hydrogen and the central atom. Down the group, increasing atomic size results in a longer, weaker bond, which means $\mathrm{H}_2\mathrm{Te}$ has a lower bond dissociation enthalpy compared to $\mathrm{H}_2\mathrm{S}$. This reduction in energy required to break the bond directly allows $\mathrm{H}_2\mathrm{Te}$ to ionize more effectively to release $\mathrm{H}^+$. Consequently, the lower bond dissociation enthalpy is the reason for the increased acidity of $\mathrm{H}_2\mathrm{Te}$ relative to $\mathrm{H}_2\mathrm{S}$.

The ability of an element to extend its valency to match its group number is fundamentally governed by the availability of vacant d-orbitals within its valence shell. Because second-period elements such as nitrogen, oxygen, and fluorine lack a $2d$ subshell, they cannot expand their coordination sphere to reach the higher valencies typically associated with their group position. Consequently, among the nine elements provided, only these three are restricted by this orbital limitation, meaning there are 3 elements that cannot form compounds matching their group valencies.

The chemical behavior of elements is defined by their periodic properties, which vary systematically across groups and periods based on their atomic structure. Chlorine and sulfur exhibit the highest negative electron gain enthalpies, while germanium and arsenic serve as metalloids that demonstrate characteristics of both metals and non-metals. Because francium and radium are found in the lower-left corner of the periodic table, they possess the largest atomic sizes, and fluorine and oxygen are recognized for having the highest electronegativity.

Oxygen stands out from other group 16 elements primarily because of its very small atomic radius and exceptionally high electronegativity. These distinct physical traits enable oxygen to form strong hydrogen bonds and facilitate stable $p\pi-p\pi$ multiple bonding, characteristics that the heavier members of the group do not exhibit. Furthermore, the absence of d-orbitals in its valence shell restricts its bonding capacity, which fundamentally differentiates its chemical behavior from that of sulfur and the other elements in the group.

Ionization enthalpy is governed by the interplay between increasing effective nuclear charge and the specific stability of an element's electronic configuration as one moves across the periodic table. While the atomic radius generally decreases from left to right, Nitrogen exhibits an exceptionally high ionization energy because its half-filled $2p^3$ orbital provides a stable electronic arrangement that resists the removal of an electron. Comparing these specific elements, the energy required follows the order $\mathrm{Al} < \mathrm{Si} < \mathrm{C} < \mathrm{N}$, which identifies Nitrogen as the element requiring the greatest amount of energy to ionize.

Ionization enthalpy generally increases across a period due to the rise in effective nuclear charge, though electronic configuration stability creates notable exceptions. Because Beryllium features a fully filled $2s^2$ subshell, it demonstrates a higher ionization enthalpy than Boron, whose single $2p^1$ electron is more easily removed. Similarly, the half-filled $2p^3$ configuration of Nitrogen provides additional stability, resulting in a higher ionization enthalpy than Oxygen, which has a $2p^4$ configuration; this leads to an overall increasing order of $E < C < A < B < D$.

While the p-block encompasses a diverse array of metals, non-metals, and metalloids, the d-block consists exclusively of metallic elements, rendering the first statement inaccurate. Because non-metals possess a greater effective nuclear charge, they exert a stronger attraction on their valence electrons, which necessitates higher energy for electron removal, known as ionisation enthalpy, and results in a greater propensity to attract bonding electrons, termed electronegativity, when compared to metals.

Periodic trends dictate that as you move across a period from left to right, atomic size and metallic character generally decrease, whereas moving down a group increases these properties. By applying these fundamental rules to the given elements, the established order for atomic radii and metallic character follows the sequence $B' < A' < D' < C'$, confirming the validity of the first two statements. Similarly, comparing ionic radii based on these positions yields the consistent trend $B'+ < A'+ < D'+ < C'+$. Because the size relation $D' < C' < B' < A'$ contradicts these observed patterns, that particular assertion is clearly false, leaving the remaining three statements as the accurate interpretation of the elements' properties.

Periodic trends within Group 13 elements are heavily dictated by the presence of $d$ and $f$ orbitals, which provide poor shielding and cause significant irregularities in physical and atomic properties. The atomic radius generally increases as the principal quantum number increases down the group, but due to screening effects and the contraction of orbitals, the specific trend is $Tl > In > Al > Ga > B$. Similarly, the ionic radius for the $M^{+3}$ ion follows the order $Tl > In > Ga > Al > B$ as it reflects these complex electronic shielding variations.

Melting points are determined by the strength of the atomic lattice, where Boron's unique covalent network structure results in the highest value, leading to the overall trend $B > Al > Tl > In > Ga$. The first ionization enthalpy, $\Delta_i H_1$, is influenced by the effective nuclear charge and the ease of electron removal; Boron holds its valence electrons most tightly, while poor shielding in heavier elements creates the specific order $B > Tl > Al \approx Ga > In$.

Across any period in the periodic table, the first ionisation enthalpy actually increases rather than decreases, making the assertion inherently incorrect. This trend occurs because the nuclear charge, represented by $Z$, grows as protons are added to the nucleus, creating a stronger attraction for valence electrons. Since this increase in nuclear charge effectively outweighs the constant shielding effect of inner shells, removing an electron becomes more difficult, which confirms that the stated reason is scientifically accurate.

The properties of these ions are defined by their status as an isoelectronic series, where the total number of electrons remains constant even though their nuclear charges vary. Each of these species contains $10$ electrons, satisfying the definition of being isoelectronic, but their nuclear charges differ because they belong to elements with atomic numbers ranging from $8$ to $12$, meaning they do not share the same nuclear charge. Within an isoelectronic group, the ionic radius decreases as the nuclear charge increases because a higher concentration of protons exerts a stronger electrostatic attraction on the same electron cloud. This physical relationship dictates that $\mathrm{O}^{2-}$ possesses the largest ionic radius due to its lower nuclear charge, while $\mathrm{Mg}^{2+}$ exhibits the smallest due to its higher nuclear charge.

Electronic configuration is governed by the Aufbau principle, which dictates that electrons occupy atomic orbitals starting from the lowest energy level based on the atomic number of the element. By identifying the nearest noble gas core, one can systematically determine the arrangement of valence electrons in the higher subshells.

Nitrogen, with an atomic number of 7, utilizes the helium core followed by $\mathrm{[He]}\,2\,s^{2}\,2\,p^{3}$ to account for its seven electrons. Sulfur, having an atomic number of 16, uses the neon core followed by $\mathrm{[Ne]}\,3\,s^{2}\,3\,p^{4}$. Bromine, with an atomic number of 35, builds upon the argon core and fully occupies the $3d$ orbitals before filling the $4s$ and $4p$ shells, resulting in $\mathrm{[Ar]}\,3\,d^{10}\,4\,s^{2}\,4\,p^{5}$. Krypton, possessing an atomic number of 36, completes the sequence by filling the $4p$ subshell, expressed as $\mathrm{[Ar]}\,3\,d^{10}\,4\,s^{2}\,4\,p^{6}$.