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

Isoelectronic species contain the same total number of electrons, making their ionic radius dependent on the effective nuclear charge pulling on those electrons. As the positive charge on a cation increases, the effective nuclear charge also increases, which significantly pulls the electron cloud inward and reduces the overall size. Therefore, the assertion that a smaller positive charge leads to a smaller ionic radius is incorrect, as a lower positive charge actually corresponds to a larger ionic radius in comparison to its more highly charged counterparts.

The first ionization enthalpy represents the energy required to remove an electron from a neutral atom, while electron gain enthalpy represents the energy change when that same electron is added back to the resulting cation. Because these two processes are exact opposites, their enthalpy changes are equal in magnitude but opposite in sign. For $\mathrm{Na}$, removing an electron involves the process $\mathrm{Na} \rightarrow \mathrm{Na}^{+} + \mathrm{e}^{-}$ with an energy absorption of 5.1 eV. Reversing this reaction to represent the addition of an electron to the sodium ion, $\mathrm{Na}^{+} + \mathrm{e}^{-} \rightarrow \mathrm{Na}$, requires changing the sign of the enthalpy value, resulting in -5.1 eV.

To maximize the ratio of the cation radius to the anion radius, expressed as $\frac{r_c}{r_a}$, one must identify the compound composed of the largest possible cation and the smallest possible anion. Comparing the ions present in the options, $\text{Cs}^+$ possesses the largest ionic radius among the available cations, while $\text{F}^-$ represents the smallest ionic radius among the available anions. Consequently, the combination of these two ions in $\text{CsF}$ results in the highest numerical value for this ratio.

Atomic size is primarily determined by effective nuclear charge, which increases across a period causing the radius to decrease, and by the number of principal energy levels, which increases down a group causing the radius to grow. Within the third period, atomic radii follow the decreasing sequence Cl < P < Mg. Because calcium resides in the fourth period directly below magnesium, it contains an extra electron shell that gives it a larger atomic radius than the others, establishing the overall trend $Cl < P < Mg < Ca$.

Isoelectronic species possess the same total number of electrons, meaning their ionic size is determined by the effective nuclear charge, $Z_{eff}$, which increases as the atomic number increases. A higher atomic number exerts a stronger electrostatic attraction on the electron cloud, pulling it closer to the nucleus and resulting in a smaller ionic radius. Consequently, because these ions share 18 electrons but vary in proton count, the species with the lowest atomic number exhibits the largest radius, creating a consistent trend of decreasing size from sulfur to calcium.

The magnitude of electron gain enthalpy is primarily determined by atomic size, where smaller atoms experience significant interelectronic repulsion that hinders the addition of an extra electron. Although smaller period 2 elements generally have lower electron gain enthalpies than their period 3 counterparts, fluorine retains a higher value than sulfur because it belongs to group 17, which has a stronger inherent tendency to gain electrons. Chlorine possesses the highest electron gain enthalpy, establishing the increasing order $O < S < F < Cl$.

The ionization energy of an atom is significantly influenced by the stability of its electronic configuration, where half-filled subshells require more energy to disrupt due to their inherent symmetry. Among the options provided, $Ne[3s^{2}3p^{3}]$ contains a half-filled $3p$ orbital, providing the extra stability that makes it more difficult to remove an electron than the configurations found in $Ne[3s^{2}3p^{1}]$ or $Ne[3s^{2}3p^{2}]$. Although $Ar[3d^{10}4s^{2}4p^{3}]$ also has a half-filled outer orbital, the presence of the $3d^{10}$ electrons exerts a shielding effect that reduces the effective nuclear charge experienced by the outer valence electrons, making them more easily removed compared to the stable $Ne[3s^{2}3p^{3}]$ structure.

The electron gain enthalpy of halogens is expected to decrease down the group, but fluorine deviates from this pattern due to its exceptionally small atomic size. This compact size creates significant interelectronic repulsion for an incoming electron, which reduces the energy released during electron capture compared to that of chlorine. Therefore, the actual order of electron gain enthalpy is $Cl_2 > F_2 > Br_2 > I_2$, which contradicts the arrangement presented for this property.

Ionization enthalpy is primarily determined by how tightly the nucleus holds onto the outermost electron, a force that diminishes as the atomic size increases and the screening effect becomes more prominent. In the configuration $1s^{2}2s^{2}2p^{5}3s^{1}$, the valence electron occupies the third principal energy shell, placing it further from the nucleus and subjecting it to greater shielding than electrons residing in the second shell. Consequently, this increased distance and effective screening drastically reduce the energy required to dislodge the electron compared to the other listed configurations.

These species are all isoelectronic, meaning they each possess eighteen electrons, so their relative sizes are determined primarily by the strength of the nuclear attraction. As the atomic number increases from sulfur to calcium, the number of protons rises, causing the nucleus to pull the shared electron cloud more tightly inward, which progressively decreases the ionic radius. Because the sulfide ion has the lowest atomic number and thus the weakest nuclear hold, it expands to the largest size, while the calcium ion experiences the strongest attraction and is the most compact. This inverse relationship between nuclear charge and ionic radius results in the specific order $Ca^{2+} < K^{+} < Ar < Cl^{-} < S^{2-}$.

The proton affinity of an ionic species is governed by the availability of its lone pair to donate to a proton, a tendency dependent on the atom's electronegativity and size. Across a period, basicity decreases as electronegativity increases because more electronegative atoms hold their electrons more tightly, which explains why $NH_2^-$ is a stronger base than $F^-$. As you move down a group, basicity declines with increasing atomic mass due to a decrease in charge density, which confirms that $F^-$ is more basic than $I^-$ and $HO^-$ is more basic than $HS^-$. Among the given species, $NH_2^-$ possesses the greatest proton affinity because it benefits from nitrogen's relatively lower electronegativity compared to the other atoms.

Bond dissociation energy does not follow a simple periodic trend because of the unusually weak nature of the fluorine-fluorine bond. While the values for the other halogens are 57 kcal/mol for $Cl_2$, 45.5 kcal/mol for $Br_2$, and 35.6 kcal/mol for $I_2$, the $F_2$ molecule possesses a surprisingly low bond dissociation energy of 38 kcal/mol. This anomaly occurs because the small size of the fluorine atoms leads to significant inter-electronic repulsion between non-bonding electrons in the 2p-orbitals, which destabilizes the bond and causes it to be weaker than the $Cl-Cl$ and $Br-Br$ bonds.

Ionic radius is determined by the electrostatic attraction between the nucleus and the valence electrons, a magnitude quantified as the effective nuclear charge ($Z_{eff}$). Because a higher effective nuclear charge exerts a stronger inward pull on the electron cloud, the size of the ion decreases, establishing an inverse proportionality expressed as $r \propto \frac{1}{Z_{eff}}$.

Isoelectronic species possess the same number of electrons, meaning their size is primarily governed by the effective nuclear charge exerted by the protons in the nucleus. As the atomic number increases across the series from oxygen to aluminum, the increasing positive charge of the nucleus attracts the electron shell more strongly, causing the ionic radius to shrink consistently from ${O}^{2-}$ down to ${Al}^{3+}$. This trend results in the specific order ${O}^{2-} > {F}^{-} > {Na}^{+} > {Mg}^{2+} > {Al}^{3+}$, where the rising number of protons effectively pulls the shared electron density closer toward the center of the ion.

The first ionization potential usually rises as you move left to right across a period because increasing nuclear charge holds electrons more tightly. A notable exception occurs between beryllium $(2s^2)$ and boron $(2s^2 2p^1)$ because the filled $s$-orbital in beryllium creates a stable configuration that demands more energy to disrupt than the removal of a lone $p$-electron from boron. Since this difference makes it easier to ionize boron than beryllium, the actual order of ionization potential is $\mathrm{Li} < \mathrm{B} < \mathrm{Be} < \mathrm{C}$.

The first ionization potential generally increases across a period as atomic size shrinks and effective nuclear charge rises, though unique electronic stabilities create notable deviations from this trend. Beryllium, with its fully filled $2s^2$ subshell, requires more energy for ionization than boron, just as nitrogen's stable half-filled $2p^3$ subshell makes it more resistant to electron loss than oxygen. Integrating these stability effects into the overall periodic trend yields the progression $\mathrm{B} < \mathrm{Be} < \mathrm{C} < \mathrm{O} < \mathrm{N}$.

The strength of a chemical bond is heavily influenced by the inter-electronic repulsion between the non-bonding valence electrons of the bonded atoms. Because fluorine has a small atomic size, its non-bonding 2p electrons are packed closely together, resulting in significant repulsion that weakens the bond. In contrast, chlorine is larger, allowing its 3p electrons to remain further apart and minimizing this repulsive effect. Consequently, the bond energy of ${\mathrm{F}}_{2}$ is measured at 158.5 kJ/mole, whereas the bond energy of ${\mathrm{Cl}}_{2}$ is higher at 242.6 kJ/mole, contradicting the statement that the bond energy of fluorine is greater.