Chemical Bonding And Molecular Structure – NEET UG ChemistryPractice Questions & PYQs

The geometric shape of a molecule is fundamentally determined by the hybridization state of its central atom, which dictates the spatial arrangement of the bonded atoms. In the case of ${CO}_{2}$, the central carbon undergoes $sp$ hybridization, resulting in a linear geometry. Other arrangements, such as the angular structure of ${SO}_{2}$ ($sp^2$ hybridized), the planar geometry of ${CO}_{3}^{2-}$ ($sp^2$ hybridized), and the tetrahedral shape of ${SO}_{4}^{2-}$ ($sp^3$ hybridized), arise from their distinct hybridization patterns.

Lattice energy is determined by the strength of the electrostatic force between ions, which is inversely proportional to the inter-ionic distance, expressed as $U \propto \frac{1}{r_+ + r_-}$. Since every compound listed involves an alkali metal cation paired with a fluoride anion, they all share identical charges, meaning the ionic size becomes the primary variable affecting this distance. As you descend Group 1, the ionic radius increases, so sodium-being the smallest cation among the choices-allows the ions to pack most closely together, which results in the highest lattice energy for sodium fluoride.

The stability of a diatomic molecule requires a positive bond order, meaning that a bond order of zero signifies an unstable system. Furthermore, these physical properties are inversely related, such that an increase in bond order results in a decrease, rather than an increase, in bond length. Consequently, both provided statements contradict established chemical principles.

Dipole moments depend on the vector sum of bond dipoles influenced by molecular geometry, where $\mathrm{H}_2\mathrm{O}$ has a value of 1.85, $\mathrm{NH}_3$ has 1.47, and $\mathrm{CHCl}_3$ has 1.04, making the provided order accurate. Regarding lone pairs on the central Xenon atom, $\mathrm{XeF}_2$ possesses 3, $\mathrm{XeF}_4$ has 2, and $\mathrm{XeO}_3$ has 1, meaning the listed sequence is incorrect. Similarly, bond length correlates inversely with bond strength, which places the actual order as $\mathrm{N}-\mathrm{O} > \mathrm{C}-\mathrm{H} > \mathrm{O}-\mathrm{H}$, contradicting the suggested sequence. Finally, bond enthalpy is directly related to bond order, and because $\mathrm{N}_2$ has a bond order of 3, $\mathrm{O}_2$ has 2, and $\mathrm{H}_2$ has 1, the descending order of bond enthalpy is consistent with the trend in the question.

The hybridization and molecular geometry of xenon compounds are determined by applying VSEPR theory, where the steric number is calculated as the sum of sigma bonding electron pairs and lone pairs around the central xenon atom. In $\mathrm{XeO}_3$, the xenon atom undergoes $\mathrm{sp}^3$ hybridization with three oxygen atoms and one lone pair, resulting in a pyramidal shape, while $\mathrm{XeF}_2$ exhibits $\mathrm{sp}^3\mathrm{d}$ hybridization with two fluorine atoms and three lone pairs, creating a linear geometry. $\mathrm{XeOF}_4$ displays $\mathrm{sp}^3\mathrm{d}^2$ hybridization with five bonding pairs and one lone pair, leading to a square pyramidal structure, whereas $\mathrm{XeF}_6$ utilizes $\mathrm{sp}^3\mathrm{d}^3$ hybridization to accommodate seven electron pairs in a distorted octahedral geometry.

Intramolecular hydrogen bonding occurs when a hydrogen atom covalently bonded to a highly electronegative atom interacts with a lone pair on a nearby electronegative atom located within the same molecule, frequently resulting in a stable ring structure. In the specific case of ortho-nitrophenol, the spatial arrangement allows the hydrogen of the hydroxyl group to be closely positioned to the oxygen atoms of the adjacent nitro group, facilitating this internal attraction. Other substances like $HF$ undergo intermolecular hydrogen bonding, where the interaction bridges separate molecules rather than occurring internally.

The bonding nature between carbon atoms reflects the hybridization and orbital overlap within each molecule, with ethane featuring a single $\sigma$ bond while ethene consists of one $\sigma$ bond and one $\pi$ bond. Ethyne incorporates a triple bond characterized by one $\sigma$ bond and two $\pi$ bonds, which contrasts with the unique electronic structure of the $\mathrm{C}_2$ molecule containing only two $\pi$ bonds. These distinct bond compositions arise from the varying degrees of unsaturation required to maintain the molecular geometry of each carbon-based species.

The boiling points of Group 16 hydrides are determined primarily by a combination of molecular mass and intermolecular attractive forces. Although boiling points generally rise with increasing molar mass for $\mathrm{H_2S}$, $\mathrm{H_2Se}$, and $\mathrm{H_2Te}$, water ($\mathrm{H_2O}$) exhibits an anomalously high boiling point because its molecules form extensive hydrogen bonds that are significantly stronger than the dispersion forces found in its counterparts. This strong intermolecular interaction elevates the boiling point of $\mathrm{H_2O}$ above the rest of the series, resulting in the overall order of $\mathrm{H_2O} > \mathrm{H_2Te} > \mathrm{H_2Se} > \mathrm{H_2S}$.

A molecule exhibits a non-zero dipole moment only when it possesses an asymmetric geometry that prevents the cancellation of individual bond dipoles generated by electronegativity differences between bonded atoms. In the case of $\mathrm{CCl}_4$, $\mathrm{CO}_2$, and $\mathrm{BF}_3$, the highly symmetric shapes-tetrahedral, linear, and trigonal planar, respectively-ensure that all bond dipoles cancel each other out completely, resulting in a net dipole of zero. Conversely, $\mathrm{HI}$ consists of a hydrogen atom and an iodine atom with different electronegativities arranged in a diatomic structure, which creates a permanent polarity and results in a non-zero dipole moment.

The molecule $\mathrm{PCl}_5$ adopts a trigonal bipyramidal structure characterized by $sp^3 d$ hybridization, featuring bond angles of 90 degrees for axial-equatorial positions and 120 degrees for equatorial-equatorial positions. Within this configuration, there are five $\mathrm{P}-\mathrm{Cl}$ sigma bonds, yet the two axial bonds are longer than the three equatorial bonds due to differing spatial repulsions, which makes the assertion that all bond lengths are identical incorrect.

Bond enthalpy represents the energy required to dissociate one mole of bonds in gaseous molecules, serving as a quantitative measure of bond strength. The nitrogen molecule (\text{N}_2$) contains a triple bond, making it the strongest among these with an enthalpy of 946.0 kJ mol-1, whereas the double bond \in the oxygen molecule (\text{O}_2$) requires 498 kJ mol-1 to break. The single bonds in the hydrogen molecule (\text{H}_2$) and hydrogen chloride (\text{HCl}$) are comparatively weaker, exhibiting enthalpy values of 435.8 kJ mol-1 and 431.0 kJ mol-1, respectively. By identifying these bond types and their corresponding energies, it is clear that $\text{HCl}$ corresponds to 431.0 kJ mol-1, $\text{N}_2$ corresponds to 946.0 kJ mol-1, $\text{H}_2$ corresponds to 435.8 kJ mol-1, and $\text{O}_2$ corresponds to 498 kJ mol-1.

Resonance describes the delocalization of electrons within a molecule or ion that cannot be represented by a single Lewis structure. While the $\mathrm{O}_3$ molecule only possesses two distinct resonance structures, the carbonate ion, $\mathrm{CO}_3^{2-}$, is characterized by three equivalent canonical forms where the double bond is shared equally among the three oxygen atoms attached to the central carbon. In other cases, $\mathrm{BF}_3$ maintains a symmetric trigonal planar geometry that causes the dipoles of the three $\mathrm{B-F}$ bonds to cancel out completely for a net dipole moment of zero, and $\mathrm{NH}_3$ exhibits a significantly higher dipole moment than $\mathrm{NF}_3$ because the bond dipoles reinforce the lone pair in $\mathrm{NH}_3$, whereas they oppose it in $\mathrm{NF}_3$.

The molecular geometry of these compounds is determined by the arrangement of electron pairs around the central atom, which includes both bonding pairs and lone pairs according to VSEPR theory. Nitrogen in $\mathrm{NH}_3$ is $sp^3$ hybridised with one lone pair, leading to a trigonal pyramidal shape. For $\mathrm{BrF}_5$, the $sp^3 d^2$ hybridised central atom has five bonding pairs and one lone pair, creating a square pyramidal geometry, while $\mathrm{XeF}_4$ also adopts $sp^3 d^2$ hybridisation but with two lone pairs, resulting in a square planar structure. Lastly, $\mathrm{SF}_6$ exhibits $sp^3 d^2$ hybridisation with six bonding pairs and no lone pairs, giving it an octahedral shape.

Isoelectronic species are identified by having the exact same total count of electrons within their structure. When examining $\mathrm{Ca}^{2+}$, $\mathrm{Ar}$, $\mathrm{K}^{+}$, and $\mathrm{Cl}^{-}$, we find that each entity possesses 20 electrons, satisfying the requirement for an isoelectronic set. This uniformity in electron number is the defining property that characterizes these particular species as isoelectronic.

The thermal behavior of hydrated halides is governed by the ionic or covalent character of the compound. While the hydrated chlorides and bromides of larger alkaline earth metals like $\mathrm{Ca}$, $\mathrm{Sr}$, and $\mathrm{Ba}$ are ionic and undergo simple dehydration upon heating, those of $\mathrm{Be}$ and $\mathrm{Mg}$ possess significant covalent character due to their high polarizing power, leading them to undergo hydrolysis instead. Because the first statement incorrectly suggests hydrolysis for the former group and the second statement incorrectly claims dehydration for the latter, both assertions are false.

Molecular Orbital Theory indicates that for light homonuclear diatomic molecules like $N_2$, significant s-p mixing occurs due to the relative proximity in energy between the 2s and 2p orbitals. This interaction effectively pushes the $\sigma 2 p_z$ orbital to a higher energy level relative to the degenerate $\pi 2 p_x$ and $\pi 2 p_y$ pair. Consequently, the established order of increasing molecular orbital energy is $\sigma 1 s < \sigma^* 1 s < \sigma 2 s < \sigma^* 2 s < (\pi 2 p_x = \pi 2 p_y) < \sigma 2 p_z < (\pi^* 2 p_x = \pi^* 2 p_y) < \sigma^* 2 p_z$.

Dipole moment depends on the electronegativity difference between bonded atoms and the overall molecular geometry, which dictates the net charge separation within the molecule. Because fluorine is the most electronegative element, $\mathrm{HF}$ exhibits the highest dipole moment, whereas the perfectly symmetrical tetrahedral structure of $\mathrm{CH}_4$ results in a non-polar nature. Between the remaining two, $\mathrm{NH}_3$ has a greater dipole moment than $\mathrm{H}_2 \mathrm{S}$ due to the significantly higher electronegativity of nitrogen relative to sulfur and the pyramidal geometry that prevents the individual bond polarities from canceling out.

The octet rule states that atoms typically form bonds to surround themselves with eight electrons in their valence shell, with exceptions arising in electron-deficient species and those with expanded octets. The molecules $\mathrm{NH}_3$ and $\mathrm{CCl}_4$ complete their octets with 8 valence electrons, whereas $\mathrm{AlCl}_3$ has 6 valence electrons, $\mathrm{BeCl}_2$ has 4 valence electrons, and $\mathrm{PCl}_5$ has 10 valence electrons. Since three of the listed species deviate from the requirement of exactly eight valence electrons, the count of species not satisfying the octet rule is 3.

Molecular orbital theory describes the formation of molecular orbitals based on the overlapping of atomic wave functions, where the resulting symmetry is determined by the orientation of these orbitals. Specifically, molecular orbitals derived from $2P_x$ and $2P_y$ atomic orbitals are inherently unsymmetrical around the internuclear bond axis, as they contain a nodal plane along that axis. Meanwhile, the formation of bonding molecular orbitals occurs through the constructive interference of electron waves, causing them to reinforce each other, whereas $\pi$-bonding molecular orbitals characteristically feature higher electron density positioned above and below the internuclear axis. Finally, antibonding molecular orbitals, such as the $\pi^*$ orbital, are distinguished by the presence of a node located directly between the nuclei, resulting from the destructive interference of wave functions.

The bond angles in molecules are determined by Valence Shell Electron Pair Repulsion (VSEPR) theory, which explains how lone pair-lone pair and lone pair-bond pair repulsions force bonding atoms into positions that minimize these interactions. Carbon dioxide, $\mathrm{CO}_2$, adopts a linear geometry with a bond angle of $180^{\circ}$, while the $sp^3$ hybridized ammonium ion, $\mathrm{NH}_4^+$, features no lone pairs and thus maintains a regular tetrahedral angle of $109.5^{\circ}$. As lone pairs are introduced, they exert greater repulsive forces; the nitrogen atom in $\mathrm{NH}_3$ contains one lone pair that reduces the angle to $107^{\circ}$, and the two lone pairs on the oxygen in $\mathrm{H}_2\mathrm{O}$ cause even more significant compression, resulting in an angle of $104.5^{\circ}$.