Classification Of Elements And Periodicity In Properties – NEET UG ChemistryPractice Questions & PYQs

Electron affinity measures the energy released when a gaseous atom gains an electron, and although this value typically decreases moving down Group 17, Chlorine serves as a notable exception by having a higher affinity than Fluorine. Fluorine is extremely small, causing an incoming electron to face significant interelectronic repulsion within its compact $2p$ subshell, which counteracts the attraction from the nucleus. Because Chlorine has a larger atomic radius, the electrons in its $3p$ subshell are more spread out, resulting in less repulsion and allowing the atom to release more energy upon gaining an electron than the smaller Fluorine atom.

These ions are all isoelectronic, meaning they contain an identical number of electrons, so the variation in their size depends entirely on the nuclear charge. As the number of protons in the nucleus increases, the electrostatic attraction toward the electron shell becomes stronger, pulling the shell closer and reducing the ionic radius. Consequently, $Al^{3+}$ exhibits the most significant contraction among the group due to its highest positive charge, making it the smallest.

Determining an element's position begins by detailing its electronic configuration, which for atomic number 33 is $1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 3d^{10} 4s^{2} 4p^{3}$. Since the outermost shell contains two electrons in the 4s subshell and three in the 4p subshell, the atom possesses a total of five valence electrons. Elements exhibiting this specific arrangement in their valence shell are classified in the fifth group, which is historically designated as Group V A.

Pauling's electronegativity scale is founded on the principle that the distribution of shared electrons in a chemical bond depends on the relative attraction of the bonded atoms for those electrons. When there is a significant difference, denoted as $\Delta EN$, in the electronegativity values of two atoms, the electron density shifts toward the more electronegative element, resulting in a bond with substantial polar character. Consequently, by evaluating these values, we can determine the degree of ionic nature within a bond and predict the overall polarity of a molecule.

Atomic volume represents the volume occupied by one gram of an element, and its trend across a period is governed by the interplay between atomic radius and density. Moving from left to right, the atomic radius initially contracts due to an increasing effective nuclear charge, which causes the atomic volume to decrease. As the period progresses toward the noble gases, the influence of atomic mass and structural packing changes, leading to a subsequent increase in atomic volume.

Non-metals possess a strong affinity for electrons because they seek to fill their valence shells to achieve a stable electronic configuration. This characteristic process, represented by reactions such as $X + e^{-} \rightarrow X^{-}$, results in the formation of negatively charged ions. Because they readily attract electrons in this manner, non-metals are fundamentally classified as being electronegative.

The ionization potential represents the energy necessary to remove the most loosely bound electron from a neutral gaseous atom, and it is significantly influenced by the stability of the electronic configuration. Beryllium features a completely filled $2s^{2}$ subshell, which provides a high degree of stability and requires an ionization energy of 9.32 eV. In contrast, Boron possesses a $1s^{2} 2s^{2} 2p^{1}$ configuration where the single electron in the $2p$ orbital is effectively shielded by the inner shells, making it easier to remove and resulting in a lower ionization potential of 8.29 eV.

Electron gain enthalpy represents the energy released when a neutral atom gains an electron, where a more negative value signifies a greater tendency to attract an electron. Although atomic size typically dictates that smaller atoms release more energy upon adding an electron, fluorine is a notable exception due to its extremely compact electron cloud, which generates significant inter-electronic repulsion that limits the energy released to -328 kJ/mol. Conversely, chlorine provides a more favorable configuration for the incoming electron, resulting in a higher energy release of -349 kJ/mol, followed by bromine at -325 kJ/mol and iodine at -295 kJ/mol, establishing the observed trend as $\text{Cl} > \text{F} > \text{Br} > \text{I}$.

Atomic size follows clear periodic trends where neutral atoms are larger than their corresponding cations, and among isoelectronic ions, a higher nuclear charge effectively pulls the electron cloud closer, resulting in a smaller radius. In the third period, magnesium has a larger atomic radius than aluminum because atomic size decreases from left to right. Furthermore, both neutral atoms are significantly larger than their ionized forms, $\mathrm{Mg}^{2+}$ and $\mathrm{Al}^{3+}$. Between the two cations, the higher atomic number of aluminum means it exerts a stronger attraction on the electrons, making $\mathrm{Al}^{3+}$ the smallest species in the set. Consequently, the claim that aluminum and $\mathrm{Mg}^{2+}$ are the extremes is inaccurate, as the largest species is actually neutral magnesium and the smallest is $\mathrm{Al}^{3+}$.

The remaining statements are factually sound within the context of chemical principles. The IUPAC systematic naming convention correctly assigns the name Unnilseptium to the element with atomic number 107. Lithium and magnesium exhibit a diagonal relationship because of their similar charge densities and polarizing power, which causes them to behave similarly in various chemical reactions. Finally, in the complex ion $[\mathrm{AlCl}(\mathrm{H_2O})_5]^{2+}$, the central aluminum atom maintains an oxidation state of +3, and it exhibits a covalency of 6 by forming coordinate bonds with five water molecules and one chloride ion.

When evaluating the size of isoelectronic ions, where each species contains exactly 18 electrons, we must look to the nuclear charge to determine their relative dimensions. Ions with a lower atomic number exert a weaker attractive force on the same electron shell, causing the electron cloud to expand outward. Since the ${S}^{2-}$ ion has the lowest atomic number in this group, its nucleus pulls the electrons the least effectively, resulting in the largest ionic size compared to ${Cl}^{-}$, ${K}^{+}$, and ${Ca}^{2+}$.

Metallic character, which represents the tendency of an element to donate electrons, strengthens as you move down a group and weakens from left to right across a period. This property is inversely proportional to electronegativity, expressed as $\text{Metallic character} \propto \frac{1}{\text{Electronegativity}}$. Comparing the periodic positions of these elements, sodium and magnesium lie to the left in period 3, while silicon and phosphorus reside further right, and beryllium is situated in period 2 above magnesium. Since electronegativity values increase as $\text{Na}=0.9$, $\text{Mg}=1.2$, $\text{Be}=1.5$, $\text{Si}=1.8$, and $\text{P}=2.1$, the metallic character must follow the reverse sequence, yielding the increasing order $\mathrm{P} < \mathrm{Si} < \mathrm{Be} < \mathrm{Mg} < \mathrm{Na}$.

Successive ionization energy spikes occur when removing an electron forces a transition from a valence shell into a stable, fully filled noble gas core. An abnormally large increase between the second and third ionization energies identifies an element with exactly two valence electrons, as the third electron must be removed from a more stable, inner shell. Among the provided choices, $1 s^{2}, 2 s^{2}, 2 p^{6}, 3 s^{2}$ is the only ground-state configuration that provides these two valence electrons before entering a stable core, consistent with the observed energy gap.

The atomic number of an element is determined by the total count of its electrons, which, for the configuration $1s^2 2s^2 2p^6 3s^2 3p^3$, sums to 15. This electronic arrangement identifies the element as belonging to group VA of the periodic table. To find the atomic number of the element located immediately below it in the same group, one must account for the inclusion of the transition series by adding 18 to the original atomic number. Calculating $15 + 18 = 33$ provides the atomic number for the element positioned directly below the initial one.

When species are isoelectronic, they possess the same number of electrons, meaning the size of the ion is determined solely by the magnitude of the nuclear charge pulling on that electron shell. As the atomic number increases from ${Na}^{+}$ to ${Si}^{4+}$, the number of protons rises, causing a significantly stronger inward electrostatic pull on the electrons. This intensified attraction effectively draws the electron cloud closer to the nucleus, leading to a consistent decrease in ionic radius as the positive charge magnitude grows.

Ionization enthalpy trends across a period generally reflect an increase in effective nuclear charge, yet specific stability patterns disrupt this steady progression. Nitrogen maintains a relatively high ionization enthalpy of 1402 kJ mol⁻¹ due to the inherent stability of its half-filled $2p^3$ subshell, whereas Oxygen, with a $2p^4$ configuration, exhibits a lower value of 1314 kJ mol⁻¹ because removing an electron facilitates the formation of a stable $2p^3$ state. While the assertion that Oxygen's ionization enthalpy is lower than both Nitrogen and Fluorine is correct, the provided reason only addresses the electronic configuration factor relevant to the comparison with Nitrogen. It fails to account for the significantly higher ionization enthalpy of Fluorine (1681 kJ mol⁻¹), which is driven by a much stronger nuclear charge and smaller atomic radius, meaning the reason does not provide a comprehensive explanation for the entire statement.

The size of an ionic species relative to its neutral atom depends on the occupancy of its electron cloud. When an iodine atom loses an electron to form a cation ($I^{+}$), the consistent nuclear charge pulls the remaining electrons closer to the nucleus, significantly reducing its size compared to the parent atom. Conversely, adding an electron to create an anion ($I^{-}$) increases electron-electron repulsion, which causes the electron cloud to expand and results in a larger radius than the neutral atom ($I$). Consequently, ordering these species by size produces the sequence $I^{-} > I > I^{+}$.

The arrangement of elements into chemical families depends on the number of electrons residing in the valence shell, as this feature dictates their periodic table position. While the configurations $[ {Xe}] 4 f^{14} 5 d^{10} 4 s^{2}$, $[ {Kr}] 4 d^{10} 5 s^{2}$, and $[ {Ar}] 3 d^{10} 4 s^{2}$ each feature two electrons in the valence shell, establishing them as members of the same group, the configuration $[ {Ne}] 3 s^{2} 3 p^{5}$ possesses seven valence electrons. This difference in electron count indicates that the latter element belongs to a distinct chemical family.

Metallic character is fundamentally defined by the ease with which an element loses electrons to form cations. Moving across a period, the effective nuclear charge ($Z_{eff}$) increases, which draws valence electrons closer to the nucleus and makes them harder to remove, leading to a decrease in metallic character. As you move down a group, the number of electron shells increases, which increases atomic size and reduces the electrostatic pull on outer electrons, thereby making them easier to lose and increasing the overall metallic nature.

Isoelectronic species must contain an identical total number of electrons, and atomic radii trend upwards by increasing magnitude as one moves down a group. Based on these principles, Ar, $K^+$, $Cl^-$, $Ca^{2+}$, and $S^{2-}$ are isoelectronic with 18 electrons, and the atomic radius of Cs at 262 pm is greater than that of Rb at 244 pm and Li at 152 pm, making statements C and E correct. Conversely, Ga and Cs both possess low melting points of 303 K and 302 K respectively, nitrogen and chlorine share an equal Pauling electronegativity of 3.0, and the first ionization enthalpy order is $Si > Mg > Al > Na$, where Mg is higher than Al due to greater $3s$ orbital penetration.

Main group elements are defined as those where the valence electrons occupy either $s$ or $p$ orbitals, encompassing elements found in groups 1, 2, and 13 through 18 of the periodic table. Comparing the provided configurations, $[\mathrm{Ne}] 3 s^1$ represents an $s$-block element and $[\mathrm{Kr}] 4 d^{10} 5 s^2 5 p^5$ represents a $p$-block element, confirming both belong to the main group. In contrast, configurations such as $[\mathrm{Ar}] 3 d^3 4 s^2$, $[\mathrm{Ar}] 3 d^{10} 4 s^1$, and $[\mathrm{Rn}] 5 f^0 6 d^2 7 s^2$ designate electrons to $d$ or $f$ subshells, which categorizes them as transition or inner transition elements rather than main group members.