Classification Of Elements – JEE Main ChemistryPractice Questions & PYQs

Acidic strength for group 16 hydrides depends primarily on the ease with which the $\mathrm{H}-\mathrm{E}$ bond breaks to release a proton, where a weaker bond correlates to higher acidity. Because the $\mathrm{H}-\mathrm{Te}$ bond is weaker and requires less energy to dissociate compared to the $\mathrm{H}-\mathrm{Se}$ bond, $\mathrm{H}_2\mathrm{Te}$ releases protons more readily and is therefore a stronger acid than $\mathrm{H}_2\mathrm{Se}$, a fact supported by the $\mathrm{pK}_{a_1}$ values of 2.6 for $\mathrm{H}_2\mathrm{Te}$ and 3.89 for $\mathrm{H}_2\mathrm{Se}$. Consequently, the claim that $\mathrm{H}_2\mathrm{Se}$ is more acidic is inaccurate, while the assertion that $\mathrm{H}_2\mathrm{Se}$ possesses a higher bond dissociation enthalpy is correct.

Ionization enthalpy is primarily determined by the effective nuclear charge and the stability of the electronic configuration, where half-filled $d$-orbitals exhibit extra stability due to symmetry and exchange energy. The $\mathrm{Mn}^{2+}$ ion has a stable $3d^{5}$ configuration, making the removal of an electron extremely difficult and requiring high energy, while $\mathrm{Fe}^{2+}$ has a $3d^{6}$ configuration and reaches the stable $3d^{5}$ state upon losing one electron. Consequently, the ionization enthalpy of $\mathrm{Mn}^{2+}$ is higher than that of $\mathrm{Fe}^{2+}$, rendering the stated relationship $\mathrm{Mn}^{2+} < \mathrm{Fe}^{2+}$ incorrect.

The foundation of the modern periodic table lies in the discovery that the properties of elements are periodic functions of their atomic numbers rather than their atomic weights, rendering the first statement incorrect. Similarly, while elements with similar outer electronic configurations display analogous chemical behaviors, they are organized into the same group-the vertical columns of the table-rather than the same period.

The number of elements contained within a period is governed by the energy level being filled during the electronic filling process. Because each atomic orbital can accommodate a maximum of two electrons, the capacity of an energy level is exactly twice the number of its available orbitals. Therefore, the number of elements in a period is double the number of atomic orbitals available in that energy level, making the final assertion false.

Elements located at the extreme left of the periodic table, specifically Group 1 elements, characterize themselves by forming basic oxides rather than acidic ones. Conversely, non-metals positioned at the extreme right, such as those in Group 17, produce oxides that react with water to yield acidic solutions. This fundamental difference in chemical nature makes the first statement false while confirming the accuracy of the second statement.

Removing an electron from an isolated gaseous atom requires overcoming the force of attraction between the nucleus and the valence electron, making the process inherently endothermic. Regarding ionization energy trends, beryllium possesses a stable, fully filled $2s^2$ subshell that holds its electrons more tightly than the higher energy $2p$ electron found in boron, which explains why the first ionization energy of boron is lower than that of beryllium. These two principles confirm the validity of the statements regarding ionization energy.

Group 14 elements exhibit a wide range of melting points dictated by their internal bonding, with carbon (atomic number 6) and silicon (atomic number 14) forming robust giant covalent networks that require substantial thermal energy to disrupt. Moving to the metallic members of this group, the established trend of $\mathrm{C} > \mathrm{Si} > \mathrm{Pb} > \mathrm{Sn}$ indicates that tin (atomic number 50) possesses a lower melting point than lead (atomic number 82). Therefore, among the elements provided, the one with atomic number 50 results in the lowest melting point.

The inert pair effect determines the relative stability of oxidation states in group 14 elements, where the stability of the $+2$ oxidation state increases down the group while the $+4$ oxidation state becomes less stable. Given that elements $A$ and $B$ are Sn and Pb respectively, $Sn^{2+}$ acts as a reducing agent because it readily oxidizes to the more stable $Sn^{4+}$ state. In contrast, $Pb^{4+}$ functions as a strong oxidizing agent because it easily undergoes reduction to the more stable $Pb^{2+}$ state.

Atomic radii decrease across a period due to increased effective nuclear charge, while they increase down a group because additional electron shells are filled. Between the elements provided in the incorrect option, magnesium possesses an extra electron shell compared to beryllium, making it the larger atom. Because the actual periodic trend is $\mathrm{Be} < \mathrm{Mg}$, the sequence starting with $\mathrm{Be} > \mathrm{Mg}$ is fundamentally flawed.

Ionization energy represents the energy required to remove electrons from an atom, and a significant, abrupt jump in these successive values indicates that an atom has transitioned from losing valence electrons to removing an electron from a more stable, inner shell. By observing the ratios between consecutive ionization steps, we can identify exactly where this transition occurs, as the removal of a core electron requires much higher energy than the removal of valence electrons.

Comparing the consecutive ionization energies of 800, 2427, 3658, 25024, and 32824 kJ/mol reveals the progression of energy needed to strip away each electron. The ratio of the second to the first energy is 3.03, the ratio of the third to the second is 1.51, and the ratio of the fourth to the third is 6.84, while the fifth to the fourth is 1.31. The most dramatic increase occurs between the third and fourth steps, which is represented by the equation $\frac{IE_4}{IE_3} = \frac{25024}{3658} = 6.84$. This massive leap confirms that the first three electrons were valence electrons, while the fourth electron came from an inner, more tightly bound shell. Since the element has three electrons in its outermost shell, it must be located in Group 13.

Elements in the modern periodic table are organized into horizontal rows known as periods, which are determined by the principal quantum number of their valence electrons. While Iridium, Osmium, and Platinum all belong to the 6th period, Palladium is positioned in the 5th period. This placement indicates that Palladium occupies a different energy level than the other three transition metals.

The first ionization enthalpy follows a periodic trend where it increases across a period and decreases down a group, which can be summarized by the relationship $IE \propto 1/r$ where $r$ is the distance from the nucleus to the valence electron. Among the given options, atomic numbers 19 and 87 belong to Group 1, while 32 and 35 reside in groups 14 and 17, respectively, which naturally exhibit higher ionization enthalpies than the alkali metals in the same period. Within Group 1, the energy required decreases as the principal quantum number increases, and because atomic number 87 is in Period 6 while atomic number 19 is in Period 4, the element with atomic number 87 has the most loosely held valence electron and the lowest overall first ionization enthalpy.

Ionization energy trends depend on a competition between effective nuclear charge and atomic shielding, which occasionally causes deviations from simple periodic expectations. In the case of lead and tin, the presence of filled d- and f-orbitals in lead creates poor shielding, resulting in a stronger nuclear pull on valence electrons and confirming that $IE_1(\text{Pb}) > IE_1(\text{Sn})$, with respective values of 715 kJ mol⁻¹ and 708 kJ mol⁻¹, making Statement I true. Conversely, when comparing germanium and silicon, the significant increase in atomic size for Ge outweighs the gain in nuclear charge, causing its ionization energy of 761 kJ mol⁻¹ to be lower than that of Si at 786 kJ mol⁻¹, which confirms that Statement II is false.

Moving from left to right across a period, the atomic radius decreases because the increasing effective nuclear charge pulls the valence electrons more strongly toward the nucleus. Since magnesium lies to the left of aluminum, magnesium actually possesses a larger atomic radius than aluminum, which makes the statement $r_{\mathrm{Mg}} < r_{\mathrm{Al}}$ a violation of the standard periodic trend.

Atomic properties such as covalent radius, ionic radius, and first ionization energy follow a regular downward trend in group 17 because of the consistent increase in atomic size. Electron affinity serves as the sole exception to this pattern, as the concentrated electronic cloud of the fluorine atom causes significant inter-electronic repulsion upon the addition of an electron, which alters its position relative to chlorine. This anomaly results in the observed affinity order of $\mathrm{Cl} > \mathrm{F} > \mathrm{Br} > \mathrm{I}$, differentiating it from the other properties that change predictably across the group.

The ionic character of a compound is determined by the electronegativity difference between its constituent elements, which is directly proportional to the magnitude of the electron gain enthalpy. Because element $E$ is consistent throughout, the ionic character of the products $EA$, $EB$, $EC$, and $ED$ increases as the magnitude of the electron gain enthalpy of the reacting elements becomes larger. Given the values of -328, -349, -325, and -295 $\mathrm{kJ\,mol^{-1}}$ for elements $A, B, C$, and $D$ respectively, the magnitudes follow the order $349 > 328 > 325 > 295$, which results in the ionic character sequence $EB > EA > EC > ED$.

The periodic table organizes elements into distinct chemical families based on their group placement, where Group 15 represents the pnictogens, Group 16 the chalcogens, Group 17 the halogens, and Group 18 the noble gases. By evaluating the specific synthetic elements within these groups, we can identify that Moscovium serves as the representative pnictogen, while Livermorium is categorized as a chalcogen. Furthermore, Tennessine fits the profile of a halogen, and Oganesson rounds out the sequence as the heaviest noble gas. This mapping establishes the direct relationship between each element symbol and its respective chemical family.

Early attempts to organize elements were based entirely on atomic mass rather than atomic number, a property that was not yet understood or utilized at that time. John Newlands' Law of Octaves proposed that elements repeat properties in groups of eight when ordered by increasing atomic mass, just as Lothar Meyer similarly identified periodic patterns by graphing physical properties against atomic mass. Because both statements falsely claim that these historical classifications utilized atomic number, both propositions are incorrect.

The first ionization enthalpy typically decreases down a group as atomic size increases, but Group 13 exhibits significant irregularities due to the poor shielding effect of intervening $d$ and $f$ orbitals. Boron holds the highest ionization enthalpy due to its small atomic size, whereas Indium represents the lowest value in the group because the poor screening provided by the $f$-subshell causes the ionization energy of Thallium to increase unexpectedly above that of Indium.

The variations in atomic and ionic size within group 13 are governed by the competition between the d-block contraction and the addition of new electron shells. The poor shielding effect of d-orbitals in gallium increases its effective nuclear charge compared to aluminum, resulting in a metallic radius for aluminum that is larger than that of gallium, which makes the first statement incorrect. In the ionic state, however, the increase in the principal quantum number becomes the dominant factor, leading to a larger ionic radius for $\mathrm{Ga}^{3+}$ compared to $\mathrm{Al}^{3+}$, thereby validating the second statement.

Periodic trends in Group 13 elements often deviate from simple downward increase or decrease due to factors like the poor shielding of d and f electrons. While atomic radius typically increases down a group, the atomic size of gallium at 135 pm is smaller than aluminum at 143 pm because of the scandide contraction, which invalidates the proposed order $\mathrm{B} < \mathrm{Al} < \mathrm{Ga} < \mathrm{In} < \mathrm{Tl}$. Similarly, density increases consistently down the group as the mass of the elements grows significantly compared to their volume, leading to the trend $\mathrm{Tl} (11.85 \;\mathrm{g}/\mathrm{ml}) > \mathrm{In} (7.31 \;\mathrm{g}/\mathrm{ml}) > \mathrm{Ga} (5.9 \;\mathrm{g}/\mathrm{ml}) > \mathrm{Al} (2.7 \;\mathrm{g}/\mathrm{ml}) > \mathrm{B} (2.35 \;\mathrm{g}/\mathrm{ml})$, which contradicts the order provided for density.

Electronegativity values reflect the tendency of an atom to attract electrons, and for this group, the values follow the order $\mathrm{B} (2.0) > \mathrm{Tl} (1.8) > \mathrm{In} (1.7) > \mathrm{Ga} (1.6) > \mathrm{Al} (1.5)$, confirming the validity of the statement written in ascending order as $\mathrm{Al} < \mathrm{Ga} < \mathrm{In} < \mathrm{Tl} < \mathrm{B}$. Regarding ionization energy, the energy required to remove the first electron follows the sequence $\mathrm{B} (801 \;\mathrm{kJ}/\mathrm{mol}) > \mathrm{Tl} (589 \;\mathrm{kJ}/\mathrm{mol}) > \mathrm{Ga} (579 \;\mathrm{kJ}/\mathrm{mol}) > \mathrm{Al} (577 \;\mathrm{kJ}/\mathrm{mol}) > \mathrm{In} (558 \;\mathrm{kJ}/\mathrm{mol})$, demonstrating that the sequence $\mathrm{In} < \mathrm{Al} < \mathrm{Ga} < \mathrm{Tl} < \mathrm{B}$ accurately reflects the experimental data.