Classification Of Elements – JEE Main ChemistryPractice Questions & PYQs

The chemical nature of an oxide is largely dictated by the position of the parent element in the periodic table, where metallic oxides generally exhibit basic characteristics and non-metallic oxides tend to be acidic. Applying these periodic trends, $\mathrm{Na}_2\mathrm{O}$ is a basic metal oxide, while the high oxidation state of chlorine in $\mathrm{Cl}_2\mathrm{O}_7$ makes it an acidic non-metal oxide. Certain oxides defy these simple classifications, such as $\mathrm{Al}_2\mathrm{O}_3$, which behaves as an amphoteric oxide capable of reacting with both acids and bases, and $\mathrm{N}_2\mathrm{O}$, which is categorized as a neutral oxide.

Elements within the same group of the periodic table share an identical valence shell electronic configuration pattern, which for group 16 is defined by the general form $ns^2 np^4$. Because the element "E" is situated in period 4, its valence electrons occupy the 4th principal energy level, giving it a configuration ending in $3d^{10} 4s^2 4p^4$. Moving to the element directly above "E" in the same group requires shifting to the 3rd period, which reduces the principal quantum number $n$ by one and results in a valence configuration of $3s^2 3p^4$.

Ionization enthalpy generally increases across a period because the rise in effective nuclear charge creates a stronger attraction between the nucleus and the valence electrons. However, atoms with exceptionally stable electronic configurations, such as Magnesium with a fully filled $s$-subshell and Phosphorus with a half-filled $p$-subshell, demand more energy for ionization than the elements immediately surrounding them. This localized stability disrupts the standard trend, resulting in the sequence $\mathrm{Al} < \mathrm{Mg} < \mathrm{Si} < \mathrm{S} < \mathrm{P}$.

Ionic radius is determined by the number of principal energy levels occupied by electrons and the strength of the nuclear pull on those electrons. For isoelectronic species like $\mathrm{Al}^{3+}$, $\mathrm{Mg}^{2+}$, and $\mathrm{Na}^{+}$, which all contain 10 electrons, the radius shrinks as positive charge increases because the higher nuclear charge pulls the electron cloud closer to the nucleus. Because $\mathrm{K}^{+}$ resides in a higher principal energy level than these 10-electron ions, its radius is significantly larger regardless of its positive charge. Combining these factors yields the trend $\mathrm{Al}^{3+} < \mathrm{Mg}^{2+} < \mathrm{Na}^{+} < \mathrm{K}^{+}$, reflecting how ionic size increases as nuclear charge decreases within an energy level and as additional shells are added.

Generally, the absolute value of electron gain enthalpy decreases as one descends the halogen group due to the increase in atomic size, which would normally predict a trend of F > Cl > Br > I. However, fluorine is an exception because its small atomic radius creates substantial inter-electronic repulsion for the incoming electron, effectively reducing the energy released compared to chlorine. This shift results in absolute electron gain enthalpy values of 349 kJ/mol for chlorine, 328 kJ/mol for fluorine, 325 kJ/mol for bromine, and 295 kJ/mol for iodine, yielding the established order of Cl > F > Br > I.

Ionisation enthalpy generally increases across a period due to an increasing effective nuclear charge, though specific electronic configurations cause periodic deviations. Magnesium ([Ne] 3s^2$) has a higher first ionisation enthalpy than Aluminium ([Ne] 3s^2 3p^1$) due to the stability of its fully filled s-orbital, and Phosphorus ([Ne] 3s^2 3p^3$) has a higher enthalpy than Sulphur ([Ne] 3s^2 3p^4$) due to its stable half-filled p-subshell. Integrating these electronic stability effects with the general periodic trend results in the order $Al < Mg < S < P$.

As we traverse a period from left to right, the atomic radius decreases, causing the nucleus to exert a stronger pull on valence electrons, which inherently reduces metallic character and promotes non-metallic properties. This trend occurs because ionization enthalpy increases, making electron loss more difficult, and electron gain enthalpy also increases, making electron capture more favorable. While the assertion correctly identifies these shifting trends, the reason is incorrect because it describes electron gain enthalpy as decreasing rather than increasing.

Bond dissociation enthalpy typically decreases down a group as atomic size increases, but the $\mathrm{F}_2$ molecule exhibits anomalous behavior due to the extreme crowding of electrons on its small atoms. Specifically, strong electrostatic repulsions between the lone pairs of the tiny fluorine atoms make the $\mathrm{F}-\mathrm{F}$ bond significantly weaker than expected, resulting in an enthalpy of 158.8 kJ mol-1, while the $\mathrm{I}-\mathrm{I}$ bond is the weakest overall at 151.1 kJ mol-1 due to steric repulsion between large iodine atoms. This creates the observed trend where $\mathrm{Cl}_2$ stands at the top with 242.6 kJ mol-1, followed by $\mathrm{Br}_2$ at 192.3 kJ mol-1:

$\mathrm{Cl}_2 > \mathrm{Br}_2 > \mathrm{F}_2 > \mathrm{I}_2$

For isoelectronic species, the size of an ion is primarily dictated by the effective nuclear charge, $Z_{\text{eff}}$. Because $\mathrm{Na}^{+}$, $\mathrm{Mg}^{2+}$, and $\mathrm{Al}^{3+}$ all contain 10 electrons, the increasing atomic number from 11 to 13 results in a stronger electrostatic attraction between the nucleus and the electron cloud. This elevated nuclear pull shrinks the ionic radius, meaning that as the net positive charge grows, the overall ionic size decreases. Given the starting point of 1.02 Å for the sodium ion, the subsequent values for the magnesium and aluminum ions must follow this downward trend, resulting in 0.72 Å and 0.54 Å respectively.

The acidic or basic nature of metal hydroxides is determined by the electropositivity of the central atom and the ionic character of the element-oxygen bond. $\mathrm{NaOH}$ and $\mathrm{Ca}(\mathrm{OH})_2$ are basic because they readily dissociate to release $\mathrm{OH}^-$ ions, whereas boron forms the acidic compound $\mathrm{B}(\mathrm{OH})_3$ due to the significant covalent nature of its bonds. Meanwhile, both $\mathrm{Be}(\mathrm{OH})_2$ and $\mathrm{Al}(\mathrm{OH})_3$ are classified as amphoteric because they possess the ability to react with both acids and bases.

Electron gain enthalpy typically becomes less negative down a group due to the increase in atomic size, but oxygen is an anomalous exception because of its extremely small radius. The high concentration of electron density within the oxygen atom generates significant electrostatic repulsion against an incoming electron, which hinders the addition process and makes the energy release much lower than expected. Because sulfur, selenium, and tellurium possess larger atomic volumes that effectively reduce these inter-electronic forces, the observed order of electron gain enthalpy follows the sequence $S > Se > Te > O$.

Oxides of non-metals are typically acidic, whereas metal oxides behave as bases. Because boron and silicon are non-metals, their respective oxides, $\mathrm{B}_2\mathrm{O}_3$ and $\mathrm{SiO}_2$, are acidic and chemically capable of reacting with bases to form salt and water, as demonstrated by $\mathrm{B}_2\mathrm{O}_3 + 2\mathrm{NaOH} \rightarrow 2\mathrm{NaBO}_2 + \mathrm{H}_2\mathrm{O}$ and $\mathrm{SiO}_2 + 2\mathrm{NaOH} \rightarrow \mathrm{Na}_2\mathrm{SiO}_3 + \mathrm{H}_2\mathrm{O}$. The other combinations include oxides of metals like $\mathrm{CaO}$ and $\mathrm{BaO}$, which are basic, or the neutral oxide $\mathrm{N}_2\mathrm{O}$, making them incorrect selections.

Amphoteric compounds are defined by their capacity to act as either an acid or a base in a chemical reaction. Among the selected substances, $\mathrm{BeO}$ and $\mathrm{Be}(\mathrm{OH})_2$ exhibit this dual nature, while $\mathrm{BaO}$ and $\mathrm{Sr}(\mathrm{OH})_2$ are basic in character and produce alkaline solutions upon reacting with water. Therefore, the total number of amphoteric compounds present is two.

The periodic trends in ionization energy reveal the electronic structure of an atom, as a dramatic increase between successive ionization energies signifies the transition from removing a valence electron to stripping one from a stable, inert gas configuration. For element $X$, the first ionization energy is 495 kJ/mol, but it spikes to 4563 kJ/mol for the second. This drastic jump indicates that the first electron was the only valence electron, and removing the second requires breaking the stable octet of a neon-like core, which identifies $X$ as sodium.

Conversely, element $Y$ displays ionization energies of 731 kJ/mol and 1450 kJ/mol. Since this moderate increase suggests that both electrons are removed from the same valence shell without disrupting an inert gas core, this pattern corresponds to the removal of two electrons from magnesium, which has the configuration $[\text{Ne}]3s^2$.

The periodic table arranges elements into metals, non-metals, and metalloids based on their valence shell configurations and positions, which allows us to classify them directly from their atomic numbers. Arsenic, represented by $X$ with atomic number 33, occupies a position along the boundary between groups, defining it as a metalloid, whereas iodine, represented by $Y$ with atomic number 53, is a halogen characteristic of non-metals. Bismuth, denoted as $Z$ with atomic number 83, sits at the bottom of its group where larger atomic size and ionization energy trends solidify its identity as a metal.

These ions form an isoelectronic series because each species contains 18 electrons with an electronic configuration of 2, 8, 8. For isoelectronic ions, the size is governed by the magnitude of the nuclear charge, where a higher atomic number leads to a stronger attractive force between the nucleus and the electron shell, causing the ionic radius to shrink. Because the atomic numbers increase from $\mathrm{P}$ ($Z=15$) to $\mathrm{S}$ ($Z=16$), $\mathrm{Cl}$ ($Z=17$), $\mathrm{K}$ ($Z=19$), and $\mathrm{Ca}$ ($Z=20$), the inward pull on the electrons becomes progressively stronger, resulting in an ionic radius order of $\mathrm{P}^{3-} > \mathrm{S}^{2-} > \mathrm{Cl}^{-} > \mathrm{K}^{+} > \mathrm{Ca}^{2+}$.

Dmitri Mendeleev developed his periodic table during a time when the inner structure of the atom, including subatomic particles like electrons and protons, was entirely unknown to scientists. While he authored the influential textbook Principles of Chemistry, had the element with atomic number 101 named in his honor, and made significant advancements in the construction of precise barometers, the state of scientific knowledge at that time did not include an understanding of atomic structure. Attributing knowledge of atomic structure to Mendeleev at the time he proposed his periodic table is therefore historically and scientifically inaccurate.

Ionization energy generally increases when moving from left to right across a period because an increasing effective nuclear charge pulls valence electrons closer, making them harder to remove. Given the third-period trend $\mathrm{Na} < \mathrm{Al} < \mathrm{Mg} < \mathrm{Si} < \mathrm{S} < \mathrm{P} < \mathrm{Cl} < \mathrm{Ar}$, magnesium has a lower ionization energy than elements to its right like chlorine and argon, while sodium, positioned to its left, exhibits a lower ionization energy than magnesium. Consequently, argon and chlorine satisfy the requirement of having a higher ionization energy than magnesium, while sodium meets the condition of having a lower ionization energy.

Qualitative inorganic analysis organizes metal cations into groups based on their distinct chemical reactivity and precipitation behavior with specific reagents. Applying this systematic identification, $\mathrm{Mn}^{2+}$ is classified in the fourth group, while $\mathrm{As}^{3+}$ resides in the group IIB category. Furthermore, $\mathrm{Cu}^{2+}$ constitutes a cation of the group IIA, and $\mathrm{Al}^{3+}$ belongs to group III, completing the required mapping for each ion.

The thermal stability of Group 2 carbonates is directly influenced by the size of the metal cation, as larger cations more effectively stabilize the bulky carbonate anion, leading to increased stability as you move down the group. This fundamental relationship explains the high thermal stability observed in $BaCO_3$ compared to lighter group members. Additionally, the solubility of these carbonates in water decreases down the group due to the reduction in the hydration energy of the cations, which is why barium carbonate is insoluble in water.