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

Nitrogen sits at the top of group 15, which grants it the highest ionization enthalpy in the group because the energy required for electron removal decreases as you descend the column. Turning to molecular behavior, the $F_2$ molecule functions as the strongest oxidising agent due to its intense electronegativity, while the $O_2$ molecule exhibits paramagnetism resulting from two unpaired electrons in its antibonding orbitals. Finally, among the noble gases of group 18, xenon is the most reactive element because its lower ionization energy facilitates chemical bond formation that the other noble gases generally resist.

Electronegativity is fundamentally governed by atomic trends, increasing as one moves left to right across a period due to rising effective nuclear charge and decreasing as one moves down a group due to the increased distance of valence electrons from the nucleus. Since silicon lies below carbon within the same group, it exhibits lower electronegativity, while moving across the second period from carbon through nitrogen and oxygen to fluorine shows a consistent rise in this property. Combining these observations leads to the clear increasing order of $Si < C < N < O < F$.

Atoms within the same period experience a reduction in size as we move from left to right because the addition of protons increases the effective nuclear charge, which is not fully offset by the additional electrons entering the same principal energy level. This strengthened electrostatic pull draws the electron cloud closer to the nucleus, resulting in a progressive contraction of atomic radii across the row. Following this established periodic trend, the elements are arranged in a decreasing order of atomic size as $\mathrm{Li} > \mathrm{Be} > \mathrm{B} > \mathrm{C}$.

Ionization enthalpy generally increases across a period because the increasing effective nuclear charge pulls electrons closer to the nucleus, yet specific stability patterns cause periodic deviations. Specifically, while Lithium starts the trend at 520 kJ mol-1, the ionization energy of Beryllium (899 kJ mol-1) is unexpectedly higher than that of Boron (801 kJ mol-1) because Beryllium's fully-filled 2s subshell is more stable than Boron's 2p configuration. Following this exception, Carbon (1086 kJ mol-1) and Nitrogen (1402 kJ mol-1) continue the general trend, with Nitrogen holding the highest value due to the stability of its half-filled 2p subshell arrangement.

When atoms attain a stable noble gas configuration, they often form isoelectronic species possessing identical numbers of electrons, where ionic size is fundamentally determined by the ratio of nuclear charge to electron count. Within such a series, the size of an ion increases as the negative charge increases because a lower number of protons are available to pull on the same electron cloud, resulting in weaker electrostatic attraction. Nitrogen, possessing the lowest atomic number of 7 among the given elements, must gain three electrons to form the $\mathrm{N}^{3-}$ anion, which gives it the highest negative charge compared to $\mathrm{O}^{2-}$, $\mathrm{F}^{-}$, and $\mathrm{Na}^{+}$. This significant disparity in nuclear pull causes the electron cloud of $\mathrm{N}^{3-}$ to expand more than that of any other ion in the group, making it the largest.

$\mathrm{NO}$ acts as a neutral oxide that does not react with acids or bases, whereas $\mathrm{CO}_2$ is categorized as an acidic oxide because it yields an acidic solution in water. $\mathrm{CaO}$ represents a typical basic oxide that interacts with acids, while $\mathrm{ZnO}$ displays amphoteric character by reacting with both acidic and basic reagents. Placing these compounds in the order $\mathrm{NO}$, $\mathrm{CO}_2$, $\mathrm{CaO}$, and $\mathrm{ZnO}$ provides the required progression from neutral to acidic to basic to amphoteric.

The acidic or basic nature of an oxide is determined primarily by the oxidation state of the central atom and the relative metallic or non-metallic character of that element, with higher oxidation states consistently driving acidic behavior. Compounds such as $\mathrm{Mn}_2\mathrm{O}_7$, $\mathrm{SO}_2$, and $\mathrm{TeO}_3$ are definitively acidic because they exhibit significant non-metallic properties or possess elevated oxidation states that facilitate the acceptance of electron pairs in chemical reactions. In contrast, other groupings fail this classification because they incorrectly categorize amphoteric oxides, such as $\mathrm{In}_2\mathrm{O}_3$ and $\mathrm{SnO}_2$, or misidentify the nature of gases like $\mathrm{NO}_2$, which acts as an acidic oxide rather than a neutral one.

Ionization enthalpy generally increases across a period as the effective nuclear charge rises, pulling electrons closer to the nucleus and making their removal more energy-intensive. However, specific electronic stability can disrupt this trend; for instance, Nitrogen possesses a highly stable half-filled $2p^3$ subshell, which requires more energy to remove an electron than the $2p^4$ configuration of Oxygen. Consequently, the ionization energy is lowest for Carbon with its $2p^2$ configuration, followed by Oxygen, then the more stable Nitrogen, and finally Fluorine, which exhibits the highest value due to its $2p^5$ configuration and intense nuclear attraction.

Isoelectronic species are defined as those possessing an identical total number of electrons. Calculating the electron count for each ion by adjusting the atomic number based on the charge reveals that $O^{2-}$ and $F^{-}$ both contain 10 electrons, just as the pair $Na^{+}$ and $Mg^{2+}$ each contain 10 electrons. Similarly, $Fe^{3+}$ and $Mn^{2+}$ are isoelectronic with 23 electrons each. In contrast, $Fe^{2+}$ has $26 - 2 = 24$ electrons, while $Mn^{2+}$ has $25 - 2 = 23$ electrons, meaning they do not form an isoelectronic pair.

The IUPAC systematic nomenclature for superheavy elements assigns names based on the Latin roots of the digits in their atomic numbers, where the sequence follows prefixes like un for 1, nil for 0, and tri for 3. Evaluating the provided list confirms that Unnilunium (101), Unniltrium (103), and Unnilhexium (106) correctly map to Mendelevium, Lawrencium, and Seaborgium, respectively. The pairing of Unununnium with Darmstadtium is incorrect, as Unununnium represents the atomic number 111, which is officially named Roentgenium.

The chemical nature of oxides is determined by their reactivity toward acids and bases. Metallic oxides like $BaO$ act as basic oxides, while non-metallic oxides like $Cl_2O_7$ behave as acidic oxides. $Al_2O_3$ is distinctly amphoteric, meaning it demonstrates both acidic and basic characteristics, whereas $CO$ remains inert toward both types of reagents, identifying it as a neutral oxide. By aligning these chemical properties with their respective categories, we observe that $CO$ corresponds to neutral (ii), $BaO$ to basic (i), $Al_2O_3$ to amphoteric (iv), and $Cl_2O_7$ to acidic (iii).

The chemical nature of an oxide is primarily determined by the metallic character and the oxidation state of the constituent element. Sodium oxide, $\mathrm{Na}_2\mathrm{O}$, functions as a basic oxide characteristic of alkali metals, whereas aluminum oxide, $\mathrm{Al}_2\mathrm{O}_3$, is classified as amphoteric because it reacts with both acids and bases. Dinitrogen monoxide, $\mathrm{N}_2\mathrm{O}$, remains chemically inert toward acidic or basic conditions, making it a neutral oxide, while the non-metallic oxide dichlorine heptoxide, $\mathrm{Cl}_2\mathrm{O}_7$, is a strong acid.

Ionization enthalpy generally rises as we move across a period because the increasing nuclear charge pulls electrons more tightly toward the nucleus. However, this trend experiences specific anomalies because fully filled and half-filled valence subshells offer extra stability, making those electrons significantly harder to remove. Specifically, beryllium's stable $2s^2$ configuration and nitrogen's stable half-filled $2p^3$ configuration require more energy for ionization than boron and oxygen, respectively, resulting in the order Li < B < Be < C < O < N < F < Ne.

The acidity of Group 2 metal oxides is inversely related to their basic character, which increases consistently as you move down the periodic table. Since basic strength follows the trend $BeO < MgO < CaO < BaO$, Beryllium oxide stands out as the most acidic member of this set because it is amphoteric, whereas all the other oxides listed are strictly basic in nature.

The trend in atomic radii for group 13 elements deviates from the standard expected pattern due to the poor shielding effect of the filled d-orbitals in gallium, which increases the effective nuclear charge and draws valence electrons closer to the nucleus. While boron is the smallest at 85 pm, gallium measures only 135 pm, making it unexpectedly smaller than the 143 pm observed for aluminum. Consequently, the atomic size follows the relationship $B < Ga < Al < In < Tl$ as the size continues to expand with indium at 167 pm and thallium at 170 pm.

The classification of an element in the periodic table depends on its valence electronic structure, which dictates its group assignment. By distributing the 114 electrons according to the Aufbau principle, we place the noble gas core of radon as the baseline and complete the orbitals to arrive at the electronic configuration $[\mathrm{Rn}] 7s^2 5f^{14} 6d^{10} 7p^2$. Since this configuration features four valence electrons in the $7s$ and $7p$ subshells, following the general $ns^2 np^2$ pattern, the element belongs to the carbon family of group 14.

Electron gain enthalpy, which represents the energy released when an atom gains an electron, typically increases in magnitude down a group, yet fluorine serves as a notable exception to this rule. Because the incoming electron in fluorine is added to a compact energy level ($n=2$), it faces intense repulsion from the valence electrons already present, significantly reducing the energy released compared to chlorine. This anomaly ensures that chlorine possesses a higher electron gain enthalpy than fluorine, which directly contradicts the sequence $I < Br < Cl < F$.

These species are isoelectronic, meaning they share the same total number of electrons, so the variation in their size depends entirely on the effective nuclear charge. As the atomic number increases, the nucleus exerts a stronger electrostatic force of attraction on the fixed number of electrons, effectively pulling the electron cloud inward and decreasing the radius according to the relationship $r \propto 1/Z$. Because calcium possesses the highest atomic number among the three, followed by potassium and then argon, it exerts the strongest pull on its electron cloud, resulting in the increasing order of $\mathrm{Ca}^{2+} < \mathrm{K}^{+} < \mathrm{Ar}$.

When ions are isoelectronic, meaning they share the same number of electrons, their relative size is determined entirely by the magnitude of the nuclear charge. As the atomic number increases, the nucleus pulls the electron cloud closer with greater force, causing the ionic radius to decrease. Because the true order of ionic radii for these species is $\mathrm{N}^{3-} > \mathrm{O}^{2-} > \mathrm{F}^{-} > \mathrm{Na}^{+} > \mathrm{Mg}^{2+} > \mathrm{Al}^{3+}$, none of the specific sequences provided in the options correctly represent this trend.

The electron gain enthalpy of elements follows a periodic trend where it becomes more negative when moving from left to right across a period and less negative when moving down a group. Because ${Ca}$ is a group 2 element with a stable, fully-filled s-orbital, it has the least negative electron gain enthalpy of the group, followed by ${Al}$ which resides in group 13. Within period 2, the electron gain enthalpy becomes progressively more negative as the effective nuclear charge increases, resulting in the sequence from ${C}$ to ${O}$ to ${F}$, where ${F}$ is the most negative. This establishes the overall order as ${Ca} < {Al} < {C} < {O} < {F}$.