Class 11 Chemistry (Chemistry Part-I) Chapter 4 Chemical Bonding And Molecular Structure

Key Concept: Ionic Bond Formation, Real-world Application

a) NaCl

[Solution Description] Ionic bonds typically form between metals and nonmetals, where metals lose electrons to achieve noble gas configurations, and nonmetals gain those electrons. Na forms $Na^+$ by losing one electron, achieving a noble gas configuration similar to neon. Cl gains this electron to become $Cl^-$. This exchange creates an electrostatic attraction that stabilizes the ionic bond in NaCl.

For Na:

$\text{Na} &\rightarrow \text{Na}^+ + e^-$

For Cl:

$\text{Cl} + e^- &\rightarrow \text{Cl}^-$

Thus, NaCl is more likely to form a stable ionic bond than other combinations such as Mg or Al with Cl due to the simpler charge balance and electrostatic stabilization.

Key Concept: Bond Enthalpy and Stability, Resonance Impact

a) $\text{Benzene, C}_6\text{H}_6$

[Solution Description] Resonance increases stability by distributing electron density over multiple structures, which affects bond characteristics such as bond order and bond enthalpy.

Compound A ($\text{Benzene, C}_6\text{H}_6$): Benzene possesses significant resonance stabilization, with alternating single and double bonds leading to a delocalized $\pi$-electron cloud. The resulting bond order is intermediate between a single and double bond, approximately 1.5, leading to enhanced stability and higher bond enthalpy than isolated single or double bonds.

Compound B ($\text{Ethylene, C}_2\text{H}_4$): Ethylene does not have resonance; it only exhibits a fixed double bond without any stabilization from resonance.

Compound C ($\text{Formaldehyde, HCHO}$): Lacks resonance structures; primarily possesses a stable structure with a distinct double bond.

Compound D ($\text{Ozone, O}_3$): Exhibits resonance but is less stabilized compared to benzene. Its resonance forms affect bond length and order, but it remains less stable overall compared to benzene.

Therefore, benzene is most affected by resonance regarding bond order and bond enthalpy, enhancing its stability.

Key Concept: Comparative Analysis, Resonance and Stability

a) $N_2$

[Solution Description]

Nitrogen gas ($N_2$) has a bond order of 3, which is typically very strong, as represented by a triple bond.

Carbon monoxide ($CO$) also has a high bond strength due to resonance stabilization; however, it involves partial charges which slightly weaken the bond compared to nitrogen’s consistent triple bond.

Ethyne ($C_2H_2$) features a carbon-carbon triple bond with significant bond strength but less than $N_2$ due to different atomic interactions.

Ozone ($O_3$) has an average bond order between a single and double bond due to resonance, leading to weaker bond strength overall compared to others.

Among these, $N_2$’s triple bond without any charge or resonance weakening makes it the strongest.

Key Concept: Hybridization and Geometry, Comparative Overlap

c) It results in stronger C-H bonds due to more s-character.

[Solution Description]

In ethyne $(C_2H_2)$, each carbon undergoes $sp$ hybridization, forming two $sp$ hybrid orbitals. These orbitals overlap with the 1s orbitals of hydrogen to form $C-H$ sigma bonds. The sp hybridization leads to greater s-character (50%) compared to other hybridizations like $sp^3$ or $sp^2$, which have lower s-character (25% and 33.3%, respectively). The higher s-character makes the $C-H$ bonds in ethyne stronger due to increased electron density between the nuclei. Thus, the bond strength is enhanced by the linear geometry and strong axial overlap.

Key Concept: Molecular Orbital Diagrams, Advanced Reasoning

c) It is paramagnetic with a bond order of 2.5.

[Solution Description] The given configuration corresponds to $C_2^+$. Calculating the bond order requires knowing the counts of bonding and antibonding electrons:

– Bonding: $\sigma_{1s}, \sigma_{2s}, \pi_{2p_x}, \pi_{2p_y}, \sigma_{2p_z}$ contribute a total of 10 electrons.

– Antibonding: $\sigma^*_{1s}, \sigma^*_{2s}, \pi^*_{2p_x}$ contribute a total of 5 electrons.

The bond order is then:

$\text{Bond order} = \frac{1}{2}(10 – 5) = 2.5.$

Given there’s one unpaired electron in $\pi^*_{2p_x}$, this species is paramagnetic. Hence, it should have peculiar properties associated with both high bond order and magnetism.

Key Concept: Real-World Application, Bond Order and Reactivity

a) More reactive due to lower bond order

[Solution Description]

The bond order of a molecule typically suggests its stability and reactivity: higher bond orders usually mean greater stability and less reactivity.

Given:

– Bond order of $Li_2$ = 1

– Hypothetical $Li_3$ has a bond order of 1.5

Since $Li_3$ has a higher bond order than $Li_2$, it should hypothetically be more stable and thus less reactive.

Hence, $Li_2$ would be more reactive in comparison to $Li_3$.

Key Concept: Resonance Impact, Advanced Misconceptions

d) Because benzene exists as a resonance hybrid, not alternating between distinct canonical forms.

[Solution Description] Benzene is often described using two equivalent Kekulé structures involving alternating single and double bonds. However, it does not exist as a dynamic equilibrium between these structures. Instead, benzene is a resonance hybrid, with all C-C bonds having equal length due to delocalized electrons. This lowers its energy compared to any individual Kekulé form, providing additional stabilization.

Key Concept: Comparative Analysis, Advanced Reasoning

b) $PCl_5$ is trigonal bipyramidal, $IF_5$ is square pyramidal

[Solution Description] $PCl_5$ has 5 bond pairs and no lone pairs, adopting a trigonal bipyramidal shape. $IF_5$, on the other hand, has 5 bond pairs and 1 lone pair, resulting in a square pyramidal shape due to increased electron repulsion by the lone pair. This changes the spatial arrangement compared to $PCl_5$.

Key Concept: Formal Charge and Resonance, Advanced Exceptions

a) Structure with +1 on N, -1 on one O

[Solution Description]

To determine the stability of resonance structures, calculate the formal charges:

– For structures with the least separation of charges: One double bond between N and O provides better charge distribution.

– Assigning electrons based on bonds: N = 5, three O’s = 6 each.

– Structure with the lowest formal charges (close to zero) is preferred.

The structure where nitrogen bears +1 and one oxygen bears -1 while others are neutral is optimal:

– Total charge matches the ion charge (-1).

– Minimized charge separations lead to enhanced stability.

Hence, the best structure maintains a charge pattern close to neutrality with minimized number of charged atoms.

Key Concept: Comparative Analysis, Hybridization and Geometry Correlation

b) $CH_4$ due to complete octet

[Solution Description] $BeCl_2$ features sp hybridization, forming linear molecules with only four electrons around beryllium, creating an unstable electron-deficient state. In contrast, $CH_4$ displays $sp^3$ hybridization resulting in a tetrahedral geometry where carbon achieves a complete octet, providing greater stability compared to $BeCl_2$. Thus, $CH_4$ is more stable due to its complete octet and effective hybridization.

Key Concept: Lewis Structure

d) H-C-H (with single bonds around C)

[Solution Description] In a Lewis structure, each hydrogen atom needs one electron to complete its shell, and carbon needs four electrons. Hence, the Lewis structure for $CH_4$ involves carbon at the center with single bonds extending to four hydrogen atoms. Each hydrogen atom shares an electron pair with carbon.

Key Concept: Real-world Application, Influence of Expanded Octet

c) $PCl_5$ shows high reactivity because axial bonds break easily.

[Solution Description]

$PCl_5$ is known to exhibit both axial and equatorial positions since it has a trigonal bipyramidal geometry because of its $sp^3d$ hybridization. Axial bonds are longer and weaker than equatorial bonds due to increased repulsion along the axis.

These differences in bond lengths lead to varying reactivity patterns. $PCl_5$ tends to dissociate more readily at the axial position, making it more reactive under certain conditions compared to $PCl_3$, where such positional flexibility does not exist.

Thus, $PCl_5$’s structure and bond characteristics contribute to its higher reactivity compared to $PCl_3$.

Key Concept: Complex Lewis Structures, Formal Charge Calculation

c) O=S-($O^-$)-O-S=O

[Solution Description]

To find the correct Lewis structure for $SO_4^{2-}$ with minimized formal charges:

1. Count total valence electrons: Sulfur has 6 and each oxygen has 6, plus 2 extra electrons for the 2- charge gives $6 + (4 \times 6) + 2 = 32$ electrons.

2. Decide on a skeletal structure: Place sulfur in the center bonded to four oxygens.

3. Distribute electrons: Start by forming single bonds between S and each O, using 8 electrons.

4. Assign remaining electrons to complete octets on oxygen atoms: 24 electrons remaining; distribute 6 each to form octets around all four oxygens.

5. Check and minimize formal charges:

– Assign formal charges using:

$$\text{Formal charge} = \text{valence electrons} – (\text{non-bonding electrons} + \frac{1}{2} \times \text{bonding electrons})$$

– Ideal should have minimum formal charges overall. The formal charge on sulfur becomes zero when it forms double bonds with two oxygens.

Therefore, the optimal structure has two S=O double bonds and two $S-O^{-}$ single bonds.

Key Concept: Real-world Application, Spatial Arrangement

d) $NH_3$

[Solution Description] To determine the solubility in water, we must consider the polarity of each molecule. Water is a polar solvent, so polar molecules with a high net dipole moment will dissolve more readily. For example:

– **$CHCl_3$** has a tetrahedral shape with different atoms attached to the carbon atom, resulting in a significant dipole moment.

– **$BF_3$** has a trigonal planar geometry where all bond dipoles cancel out, resulting in no net dipole moment.

– **$NH_3$** has a trigonal pyramidal shape leading to a high net dipole moment due to lone pair effects.

– **$CCl_4$** has a tetrahedral shape with identical atoms, thus all bond dipoles cancel out, giving it a zero net dipole moment.

The most soluble molecule in water is therefore $NH_3$, as it possesses a substantial dipole moment.

Key Concept: Bond Length and Chemical Environment, Resonance

c) Resonance in benzene equalizes bond lengths

[Solution Description]

Benzene’s structure involves resonance, where its electrons are delocalized over the entire ring. This leads to equalization of bond lengths, making them all intermediate between single and double bonds. In contrast, cyclohexane contains distinct C–C single bonds.

Specifically:

– In benzene, the resonance gives rise to an average bond order of 1.5 for the C–C bonds, leading to bond lengths that are shorter than a typical single bond but longer than a double bond.

– Cyclohexane has regular C–C single bonds with no resonance effect.

Hence, resonance in benzene results in altered chemical environments compared to cyclohexane, affecting bond lengths significantly.

Key Concept: Lone Pair Influence

a) Both Assertion and Reason are true, and Reason is the correct explanation of Assertion.

[Solution Description] Ammonia $(NH_3)$ has a trigonal pyramidal shape with a lone pair on the nitrogen atom. This lone pair contributes to the overall dipole moment by adding to the vector sum of the N-H bond dipoles, enhancing the dipole moment. In contrast, $NF_3$ also has a similar shape, but the lone pairs on fluorine atoms are highly electronegative and oppose the direction of the N-F bond dipoles, reducing the overall dipole moment. Hence, both assertion and reason are true, and the reason correctly explains why $NH_3$ has a higher dipole moment compared to $NF_3$.

Key Concept: Molecular Orbital Types

c) $\sigma$

[Solution Description] $\sigma$ molecular orbitals are defined as those formed by the head-on overlap of atomic orbitals along the bond axis.

Key Concept: Energy Considerations

a) Energy levels are too similar

[Solution Description] Hybridisation involving 4s and 3d orbitals is possible because these orbitals have relatively close energy levels in transition metals, allowing them to form stable hybrid orbitals like $sp^3d$ or $sp^3d^2$.

Key Concept: Hybridization Impact

d) Changes it to sp

[Solution Description] In acetylene $(C_2H_2)$, the carbon atoms form one sigma bond and two pi bonds with each other. The presence of these pi bonds results in sp hybridization as each carbon atom uses only one p orbital for the pi bond formation while the other two p orbitals combine with an s orbital to create two sp hybrid orbitals.

Key Concept: Covalent vs. Van der Waals

c) Less than 188 pm

[Solution Description] Generally, the covalent radius is smaller than the van der Waals radius because it considers only half the distance between nuclei of two atoms bonded together. It must be less than 188 pm.

Key Concept: Theoretical Application, Influence of Environment

a) Due to resonance stabilization enhancing bond strength

[Solution Description]

Resonance stabilization contributes to higher overall bond strength because electrons are delocalized across different structures, resulting in lower energy states and stronger effective bonding.

In Molecule A, resonance leads to stronger bonds, making them harder to break compared to Molecule B where there is no additional resonance stabilization.

Thus, Molecule A requires more energy due to the contribution of resonance to bond enthalpy.

Key Concept: Molecular Orbital Theory, Hydrogen Bonding

c) Molecular orbitals allow strong hydrogen bonding networks.

[Solution Description]

$H_2O$ involves molecular orbitals formed from atomic orbitals. Its bent shape and lone pairs lead to significant hydrogen bonding interactions with other molecules.

According to molecular orbital theory, this results in a high degree of overlap between the hydrogen-bond acceptor’s nonbonding orbitals and the donor’s bonded hydrogen orbitals, stabilizing water’s liquid structure at room temperature.

Key Concept: Overlap Types

c) Parallel overlap of d orbitals

[Solution Description] Delta bonds ($\delta$) are formed by the overlap of two parallel d-orbitals. This type of bonding occurs where there is an extensive region of overlap along two nodal planes perpendicular to the bond axis.

Key Concept: Bond Angle Variation

c) $H_2O$

[Solution Description]

Lone pair repulsions can cause bond angles to decrease. In $CH_4$, the molecule has no lone pairs, resulting in a tetrahedral structure with bond angles of $109.5^{\circ}$. $NH_3$ has one lone pair, causing its bond angles to reduce slightly to approximately $107^{\circ}$. $H_2O$ has two lone pairs that further compress the bond angle to around $104.5^{\circ}$. $BF_3$, being trigonal planar, has no lone pairs affecting bond angles, which remain at $120^{\circ}$. Thus, $H_2O$ has the smallest bond angle due to lone pair repulsion.

Key Concept: Electronic Configuration Analysis

a) $O_2$

[Solution Description]

A molecule is paramagnetic if there are unpaired electrons in its molecular orbitals.

For $O_2$:

The electronic configuration is $(σ1s)^2 (σ^∗1s)^2 (σ2s)^2 (σ^∗2s)^2 (σ2pz)^2 (π2px=π2py)^4 (π^∗2px=π^∗2py)^2$. Since there are two unpaired electrons in (π^∗2p), $O_2$ is paramagnetic.

For $N_2$:

The electronic configuration is $(σ1s)^2 (σ^∗1s)^2 (σ2s)^2 (σ^∗2s)^2 (π2px=π2py)^4 (σ2pz)^2$. All electrons are paired, so $N_2$ is diamagnetic.

Thus, $O_2$ is paramagnetic due to the presence of unpaired electrons.

Key Concept: Bonding Patterns, Comparative Electronegativity

a) $\text{NH}_3 < \text{HF} < \text{H}_2\text{O}$ [Solution Description] Hydrogen bonding strength depends on both the electronegativity of the atom bonded to hydrogen and the availability of lone pairs for accepting hydrogen bonds. Fluorine (HF) is the most electronegative, followed by oxygen $(\text{H}_2\text{O})$, and then nitrogen $(\text{NH}_3)$. However, the number of hydrogen bonds also plays a role; water can form four hydrogen bonds per molecule while HF, although strongly electronegative, typically forms fewer due to having one hydrogen and being capable of only one hydrogen bond donation. Therefore, considering both factors combined, water is still the strongest overall hydrogen bonder in many contexts. Ranking should consider this holistic view.

Key Concept: Hybridization Concept

a) sp

[Solution Description] In $BeCl_2$, the central beryllium (Be) atom forms two sigma bonds with chlorine atoms. When a molecule is linear and involves two bond pairs, the hybridization is sp. This occurs because one s orbital and one p orbital from beryllium mix to form two equivalent sp hybrid orbitals that arrange linearly. Thus, the hybridization is sp.

Key Concept: Advanced Calculation, Influence of Molecular Shape

c) 2.39 D

[Solution Description] The dipole moment of a polyatomic molecule like $H_2O$ is calculated considering both the magnitude of individual bond dipoles and the angle between them. Water’s two O-H bond dipoles each have a moment of 1.85 Debye. Using vector addition:

$$\mu_{\text{net}} = 2 \times 1.85 \, \text{D} \times \cos\left(\frac{104.5^\circ}{2}\right)$$

Calculate:

$$\mu_{\text{net}} = 2 \times 1.85 \, \text{D} \times \cos(52.25^\circ) \approx 2.39 \, \text{D}$$

Therefore, the net dipole moment of water is approximately $2.39 \, \text{D}$.

Key Concept: Bond Order Basics

b) 2

[Solution Description] To find the bond order, use the formula $\text{Bond order} = \frac{1}{2}(N_b – N_a)$. Here, $N_b = 8$ (bonding electrons), $N_a = 4$ (antibonding electrons). Calculate:

$$\text{Bond order} = \frac{1}{2}(8 – 4) = \frac{1}{2}(4) = 2.$$

The bond order of the molecule is 2.

Key Concept: Advanced Bonding Concepts, Formal Charge Calculations

a) One oxygen with formal charge of $0$, others with $-1$.

[Solution Description] For $CO_3^{2-}$, each oxygen can form either a single or a double bond with carbon, distributing varying formal charges. Determining the formal charge involves assigning electrons within bonds: shared electrons are split equally. The most stable structure will minimize formal charges across the molecule. With a double bond between carbon and one of the oxygens, that oxygen carries a charge of $0$, while the others carry $-1$, keeping total charge consistent with $2-$.

Key Concept: Resonance

c) $O_3$

[Solution Description] Resonance occurs when a molecule can be represented by two or more valid Lewis structures that differ only in the positions of electrons, not atoms. Ozone ($O_3$) shows resonance because it can have double bonds located between different pairs of adjacent oxygen atoms.

Key Concept: Real-world Applications, Ionic Bond Formation

b) Three moles

[Solution Description] To solve this, consider the reaction: $Al + Cl_2 \rightarrow AlCl_3$.

– Each Al atom loses 3 electrons to form $Al^{3+}$ ion.

– Chlorine gains these electrons to form $Cl^-$ ions.

– Balanced equation: $2Al + 3Cl_2 \rightarrow 2AlCl_3$ indicates that two moles of aluminum transfer 6 moles of electrons (since each Al transfers 3 electrons).

– Thus, per mole of Al, 3 moles of electrons are transferred.

Key Concept: Hybridisation and Molecular Geometry, Overlap and Bond Strength

d) $sp^2$-$sp^3$

[Solution Description] The structure of urea shows that each nitrogen is bonded to the carbon atom through a single bond. Carbon forms four bonds: two with nitrogen atoms, one double bond with oxygen, and has no lone pairs. Therefore, carbon’s hybridisation will be $sp^2$. The C-N sigma bond involves overlap between an $sp^2$ orbital from carbon and an $sp^3$ orbital from nitrogen.

Key Concept: Synthesis and Design, Comparative Analysis

b) Six-membered rings minimize angle strain and enhance stability.

[Solution Description] Six-membered rings are generally more stable due to lesser angle strain compared to five-membered rings. Although both configurations can establish intramolecular hydrogen bonds, six-membered rings tend to adopt a chair conformation that minimizes steric hindrance and maximizes hydrogen bonding efficiency, providing better stability.

Key Concept: Hybridisation and Geometry

c) Linear

[Solution Description]

In the case of sp hybridisation, one s orbital and one p orbital combine to form two equivalent sp hybrid orbitals. These orbitals are arranged linearly at 180 degrees to each other to minimize repulsion, resulting in a linear molecular geometry.

Key Concept: Hybridization and Bond Formation

d) sp orbitals form one $\sigma$ bond; $\pi$ bonds from p orbitals

[Solution Description] In acetylene, each carbon atom undergoes sp hybridization, forming two sp hybrid orbitals. One of these orbitals forms a $\sigma$ bond with another carbon, and the remaining ones form sigma bonds with hydrogen atoms. The unhybridized p orbitals overlap to form two $\pi$ bonds, completing the triple bond between the carbon atoms.

Key Concept: Real-world Application, Role of Electrostatic Attraction

c) Ionic bonding with strong electrostatic attraction

[Solution Description] Elements from Group 2 typically form +2 cations, and those from Group 16 form -2 anions. The binary compound resulting from these groups will have strong electrostatic attractions due to the full charge transfer. High melting points and conducting properties upon melting suggest a strong lattice structure characteristic of ionic compounds. Therefore, the compound has ionic bonds and is stable due to strong electrostatic attractions.

Key Concept: Hybrid Structure Analysis, Resonance in Complex Molecules

a) Identical C-O bond lengths arise from electron delocalization, reducing angular strain.

[Solution Description]

The carbonate ion has three resonance forms, each displaying one C=O double bond and two C-O single bonds in different placements. Resonance causes the actual structure to be a hybrid where all C-O bonds have partial double bond character, leading to identical bond lengths shorter than a typical single bond but longer than a double bond. Additionally, the bond angles approach $120^\circ$, resembling trigonal planar geometry due to even electron distribution.

Key Concept: Symmetry Requirement

d) $3d_{xy}$

[Solution Description] The combination of orbitals requires similar symmetry properties. A $3d_{xy}$ orbital has specific orientation and symmetry, so an orbital with similar characteristics, like another d orbital (specifically $3d_{xy}$), is needed for effective overlap.

Key Concept: Dipole Moment Units

b) $5.00346 \times 10^{-30} \, C \, m$

[Solution Description] The conversion factor between Debye and Coulomb meter is given as $1 \, D = 3.33564 \times 10^{-30} \, C \, m$. To convert $1.5 \, D$ to $C \, m$, we multiply: $1.5 \times 3.33564 \times 10^{-30} = 5.00346 \times 10^{-30} \, C \, m$

Key Concept: Orbital Symmetry

b) 1s and 2s

[Solution Description] For effective combination and formation of molecular orbitals, the combining atomic orbitals must have compatible symmetry. Both 1s and 2s orbitals are spherically symmetrical, allowing them to combine effectively.

Key Concept: Comparative Stability, Bond Length and Order

a) $H_2$

[Solution Description] To determine the molecule with the lowest energy state, we must identify the one with the highest bond order, as it indicates stronger bonding and lower potential energy. The bond order can be determined by using Molecular Orbital Theory:

$H_2$: Bond order = $\frac{1}{2}(2 – 0) = 1$

$He_2$: Bond order = $\frac{1}{2}(2 – 2) = 0$

$Li_2$: Bond order = $\frac{1}{2}(2 – 0) = 1$

$Be_2$: Bond order = $\frac{1}{2}(2 – 2) = 0$

Among these, $H_2$ and $Li_2$ have a bond order of 1, but considering stability through effective overlap and enthalpic considerations, $H_2$ is more stable due to its shorter bond length and higher bond dissociation energy than $Li_2$. Therefore, $H_2$ has the lowest energy state.

Key Concept: Bonding Orbitals

b) $sp^2$ and p

[Solution Description] In $BF_3$, each boron atom uses its $sp^2$ hybrid orbitals to overlap with the p orbitals of fluorine atoms, forming sigma bonds. Specifically, the overlap occurs between the boron’s $sp^2$ hybrid orbitals and the unhybridized 2p orbitals of fluorine.

Key Concept: Orbital Involvement

c) 2py and 2px

[Solution Description]

The carbon atoms in $C_2H_2$ undergo sp hybridisation, which leaves two unhybridised p orbitals on each carbon atom, specifically the 2px and 2py orbitals. These unhybridised orbitals overlap sidewise to form $\pi$ bonds.

Key Concept: Resonance Structures

c) They represent the same molecule

[Solution Description] Resonance structures are different ways of drawing the same molecule that differ only in the positions of electrons, not atoms. They have the same molecular formula, arrangement of atoms, and net charge. However, they do not exist independently; rather, the true structure is a resonance hybrid.

– Option A: Incorrect because resonance structures represent different electron placements, not energies.

– Option B: Incorrect as bond lengths are averaged in the resonant structure.

– Option C: Correct because resonance structures are just different representations of the same molecule.

– Option D: Incorrect as atomic positions remain constant.

Therefore, the correct answer is that they represent the same molecule.

Key Concept: Character Percentage

b) 50%

[Solution Description] An sp hybrid orbital results from the combination of one s orbital and one p orbital. This means there are equal contributions from s and p orbitals, hence the s-character is 50%.

Key Concept: Basic Concept

a) Both Assertion and Reason are true, and Reason is the correct explanation of Assertion.

[Solution Description] In $SF_6$, sulfur utilizes its available 3d orbitals to accommodate more than eight electrons, allowing for an expanded octet configuration. This results in $sp^3d^2$ hybridization to form six equivalent S-F bonds. Therefore, both the assertion and reason are true, with the reason correctly explaining the assertion.

Key Concept: Real-world Application, Advanced Octet Rule

d) $XeF_2$

[Solution Description] To solve this question, we need to identify which molecule has expanded valence shells beyond the typical octet. $XeF_2$ typically involves xenon forming bonds with fluorine, where xenon can expand its valency because it is in period 5 of the periodic table.

Key Concept: Orbital Overlap

d) $sp^3-sp^3$

[Solution Description] In the ethane molecule, each carbon atom undergoes $sp^3$ hybridization. One $sp^3$ hybrid orbital from one carbon overlaps with one $sp^3$ hybrid orbital from the other carbon to form a $sp^3-sp^3$ sigma bond. This is characteristic of the single bond between two carbons in alkanes.

Key Concept: Odd-Electron Molecules

a) Both Assertion and Reason are true, and Reason is the correct explanation of Assertion.

[Solution Description]

Since NO has an odd number of electrons, namely 11, it cannot distribute these to complete an octet around both nitrogen and oxygen atoms. This means the assertion that NO does not satisfy the octet rule is true. Additionally, NO having an unpaired electron indeed results in its paramagnetic nature; hence, the reason provided accurately explains the assertion.

Key Concept: Complex Shape Analysis, Advanced Hybridization

c) Seesaw, $sp^3d$

[Solution Description] To find the molecular shape and hybridization, we use VSEPR theory. The arrangement involves 5 regions of electron density (4 bond pairs + 1 lone pair), leading to trigonal bipyramidal electronic geometry. With one lone pair, the shape becomes seesaw. For hybridization, 5 orbitals are involved: s, p, px, py, d, which corresponds to $sp^3d$ hybridization.

Key Concept: Molecular Orbital Theory and Real-World Application, Molecular Orbital Energy Level Comparison

b) Due to lower energy of $\pi2p$ orbitals in NO than atomic p

[Solution Description]

In the molecular orbital diagram for NO, the bonding orbitals are filled first before filling any antibonding orbitals, which results in a net stabilization relative to separate N and O atoms.

Since NO forms by combining atomic orbitals of N and O, the overlap leads to a situation where bonding orbitals are lowered in energy more than antibonding orbitals are raised. This creates an overall lower potential energy state.

Compared with separated atoms, NO’s occupied bonding orbitals contribute to stabilizing the molecule. The presence of an unpaired electron in the antibonding orbital gives NO its paramagnetic property but doesn’t significantly destabilize it due to higher bond order (2.5).

Thus, the molecular orbital approach directly explains why NO is more stable than individual N or O atoms.

Key Concept: Geometry Indication

c) Trigonal planar

[Solution Description] In $sp^2$ hybridisation, one s and two p orbitals mix to form three $sp^2$ hybrid orbitals, leading to a trigonal planar geometry.

Key Concept: Overlapping Extent

b) Sigma bonds have more overlapping than pi bonds

[Solution Description] Sigma bonds are formed by end-to-end overlap along the internuclear axis, allowing for more extensive overlap than pi bonds, which are formed by sidewise overlap with parallel axes. Therefore, sigma bonds have a greater extent of overlapping compared to pi bonds.

Key Concept: Real-World Application

b) It affects how quickly the salt dissolves in water

[Solution Description] Lattice enthalpy affects the solubility of ionic compounds. Salts used on icy roads must dissolve readily in water to lower the freezing point efficiently and melt ice. Higher lattice enthalpy means stronger forces holding the ions together, leading to lower solubility.

Key Concept: Bond Energy, Advanced Overlap Concepts

c) Sigma bonds result from end-to-end sp-sp overlap and have higher bond energy

[Solution Description] In acetylene ($C_2H_2$), each carbon atom is sp hybridized and forms a sigma bond by head-on overlapping of sp-sp orbitals. This overlap occurs along the internuclear axis, providing strong bonding due to substantial orbital overlap leading to high bond energy. Sigma bonds exhibit greater overlapping than pi bonds, resulting in increased bond strength and stability.

Key Concept: Orbital Overlap

a) 1s and 1s

[Solution Description] In the formation of an $H_2$ molecule, two 1s atomic orbitals from each hydrogen atom overlap to form a sigma bond. This overlap is crucial for the bond formation in these molecules.

Key Concept: Complex Energy Calculations, Real-world Application

b) $-2912 kJ/mol$

[Solution Description] The total energy change involves adding the ionization enthalpy, the negative of the electron gain enthalpy, and the lattice enthalpy. This accounts for forming the ions and then creating the solid ionic structure:

$$\text{Total Energy Change} = (738\, \text{kJ/mol}) + (-(-141)\, \text{kJ/mol}) + (-3791\, \text{kJ/mol})$$

$$= 738 + 141 – 3791 = -2912\, \text{kJ/mol}$$

Thus, the total stability energy is $-2912 kJ/mol$.

Key Concept: Hybridization Basics

d) $sp^3d$

[Solution Description]

The molecule $\text{PCl}_5$ involves phosphorus and chlorine atoms. To understand the hybridization, we consider the central atom, which is phosphorus. Phosphorus has 5 valence electrons and forms bonds with 5 chlorine atoms. The required hybrid orbitals to accommodate these 5 pairs are $sp^3d$. Therefore, the hybridization of phosphorus in $\text{PCl}_5$ is $sp^3d$.

Key Concept: Symmetry Requirement

b) Same Symmetry

[Solution Description] For atomic orbitals to combine effectively to form molecular orbitals, they must have the same symmetry around the bond axis and similar energy levels to maximize overlap.

Key Concept: Bond Parameters, Polar Molecules

c) Assertion is true, but Reason is false.

[Solution Description]

The assertion is true because the difference in electronegativity between atoms causes one atom to attract shared electrons more strongly, creating a dipole moment and contributing to molecular polarity.

The reason is false. While a greater difference in electronegativity can lead to stronger ionic character, it does not directly correlate with bond enthalpy. Bond enthalpy depends on bond strength, which is influenced by various factors including but not limited to electronegativity differences, bond length, and atom size.

Thus, while both statements are related to bonds, they do not logically explain each other as described in the Assertion and Reason format.

Key Concept: Predictive Analysis, Molecular Orbital Diagram Interpretation

c) Bond order increases to 2.5; Remains paramagnetic

[Solution Description] When $O_2$ loses an electron, the configuration changes to $(σ1s)^2 (σ*1s)^2 (σ2s)^2 (σ*2s)^2 (σ2pz)^2 (π2px)^2 (π2py)^2 (π*2px)^0 (π*2py)^1$. The bond order becomes $Bond\ order = \frac{1}{2}(10 – 5) = 2.5$. With only one unpaired electron left, it remains paramagnetic.

Key Concept: Formal Charge Application, Complex Molecule Structures

b) Structure with equal double bonds from sulfur to each oxygen, resulting in zero formal charges.

[Solution Description] Sulfur dioxide ($SO_2$) can have several possible Lewis structures. For each possible structure, calculate the formal charge using:

$$\text{Formal Charge} = \left( \text{Valence electrons} – \text{Non-bonding electrons} – \frac{\text{Bonding electrons}}{2} \right)$$

Possible structures include:

1. Double bond between S and one O, single bond with another O.

2. Two double bonds between S and both O atoms.

Calculating formal charges:

– In case 1, for S: $$6 – 2 – \frac{6+2}{2} = +1$$; for double bonded O: $$6 – 4 – \frac{4}{2} = 0$$; for single bonded O: $$6 – 6 – \frac{2}{2} = -1$$.

– In case 2, for S: $$6 – 0 – \frac{8}{2} = 0$$; for both double bonded O’s: $$6 – 4 – \frac{4}{2} = 0$$.

Case 2 shows zero formal charges across all atoms, indicating this is the most stable configuration.

Key Concept: Symmetry Requirement

c) 3pz

[Solution Description]

The combining orbitals must have the same symmetry about the molecular axis. The 3pz orbital has a definite symmetry around the z-axis, so it will combine with another p orbital along the same axis.

Key Concept: Complex Geometry Prediction, Energy Considerations

c) See-saw

[Solution Description] The given molecule has four bond pairs and one lone pair around the central atom. According to VSEPR theory, we first consider what the shape would be if there were no lone pairs; with five total electron pairs, this would be a trigonal bipyramidal arrangement. However, the presence of the lone pair causes deviation from ideal geometry due to the higher repulsion exerted by the lone pair compared to bonding pairs. In a trigonal bipyramidal structure, lone pairs preferentially occupy equatorial positions because they experience less steric hindrance than axial positions. Therefore, placing the lone pair in an equatorial position, the remaining four bonding pairs form a see-saw shape. The most critical factor here is energy minimization and maximizing distance between electron pairs which results in a stable molecular geometry.

Key Concept: Molecular Orbital Theory, Comparative Bond Energies

a) Bond order is 2, less than that of NO.

[Solution Description] The bond order can be calculated using the formula:

$$\text{Bond Order} = \frac{\text{Number of electrons in bonding orbitals} – \text{Number of electrons in antibonding orbitals}}{2}$$

For the given configuration:

Number of electrons in bonding orbitals = 8.

Number of electrons in antibonding orbitals = 4.

Therefore,

$$\text{Bond Order} = \frac{8 – 4}{2} = 2.0$$

Nitrogen monoxide (NO) has a total of 15

electrons with its bond order being typically calculated as $$\frac{10 – 5}{2} = 2.5$$. Thus, compared to NO, this hypothetical molecule’s bond order is less.

Key Concept: Basic Definition

d) The structure of organic compounds

[Solution Description] The octet rule is mainly used to understand the structures of most organic compounds and applies particularly to second-period elements. It explains how atoms arrange themselves to achieve a filled valence shell with eight electrons, similar to noble gases.

Key Concept: Electron Configuration Analysis, Exception Justification

c) $[Ne]3s^2 3p^3 3d^1$

[Solution Description]

Phosphorus in $PF_5$ extends beyond the octet rule by making five bonds with fluorine atoms, utilizing an expanded valence shell. The electronic configuration of phosphorus starts at $[Ne]3s^23p^3$. When forming $PF_5$, phosphorus hybridizes its orbitals to include one additional 3d orbital, effectively using $[Ne]3s^23p^3 3d^1$ to create five equivalent $sp^3d$ hybrid orbitals.

The ability to expand the valence shell stems from phosphorus being in the 3rd period, where d-orbitals become energetically accessible. Hence, the molecule does not conform to the octet rule because phosphorus effectively uses the d orbital to accommodate more electrons.

Key Concept: Hybridization Analysis, Electronegativity Comparison

c) $NH_3$ shows $sp^3$ hybridization and nitrogen is more electronegative.

[Solution Description]

In $NH_3$, the nitrogen atom undergoes $sp^3$ hybridization, resulting in a trigonal pyramidal shape due to one lone pair of electrons and three hydrogen atoms bonded to nitrogen. The electronegativity difference between nitrogen and hydrogen contributes to the molecule’s polarity. Nitrogen, being more electronegative than hydrogen, attracts the shared electrons towards itself more strongly, leading to a polar molecule with a significant dipole moment.

Key Concept: Hybridisation and Bond Angles

d) 109.5°

[Solution Description] An $sp^3$ hybridisation involves the mixing of one s orbital and three p orbitals, resulting in a tetrahedral geometry. The bond angles in a perfect tetrahedron are 109.5°.

Key Concept: Noble Gas Reactivity, Expanded Octet

c) Xenon

[Solution Description] Xenon (Xe) is known to form stable compounds such as $XeF_2$, $XeF_4$, and $XeF_6$ where it uses the available d-orbitals to accommodate more than 8 electrons, thereby achieving expanded octets. This demonstrates xenon’s reactivity under specific conditions despite being a noble gas.

Key Concept: Theory Application

c) Trigonal pyramidal

[Solution Description]

To predict the geometry of the $PH_3$ molecule using VSEPR theory, we must first consider its electron pair arrangement. The central atom P has 5 valence electrons and forms three single bonds with hydrogen atoms, thus utilizing 3 electrons. The remaining two electrons form a lone pair. According to VSEPR theory, this results in four regions of electron density around phosphorus: three bonding pairs and one lone pair. This configuration suggests a tetrahedral electron pair geometry, but because there is one lone pair, the molecular shape becomes “trigonal pyramidal.” This is similar to ammonia ($NH_3$), which also has a trigonal pyramidal shape.

Key Concept: Hybridization Influence

d) $180^\circ$, as it results from linear arrangement.

[Solution Description] Molecules with sp hybridization have a linear shape, where the bond angle between atoms is $180^\circ$. This occurs because sp hybridization involves mixing one s orbital with one p orbital, resulting in two equivalent orbitals arranged linearly to minimize electron repulsion.

Key Concept: Comparative Bonding Analysis, Theoretical Existence

b) Be$_2$ has a higher bond order but remains unstable

[Solution Description] For He$_2$, the bond order is zero, calculated using its molecular orbital configuration $(\sigma 1s)^2 (\sigma^* 1s)^2$, indicating no stable bond forms. In contrast, Be$_2$ has a bond order greater than zero when considering its molecular structure. Despite both being theoretically weak, Be$_2$, having some positive bond order, implies a slight stability compared to He$_2$’s complete instability.

Key Concept: Hybridization Basics

b) sp³d²

[Solution Description] In SF₆, the central sulfur atom undergoes $sp^3d^2$ hybridization to form six equivalent hybrid orbitals that overlap with fluorine’s p orbitals.

Key Concept: Electron Cloud Distribution

c) Above and below the plane

[Solution Description] In ethene, the $\pi$ bond is formed by the sideways overlap of the unhybridised p orbitals on each carbon. The electron clouds are distributed above and below the plane of the molecule.

Key Concept: Geometry Identification

c) Tetrahedral

[Solution Description]

In $sp^3$ hybridisation, one s orbital mixes with three p orbitals to form four $sp^3$ hybridised orbitals which arrange themselves in tetrahedral geometry.

Key Concept: Directional Properties, Advanced Real-world Application

c) Their directional nature allowing specific alignment

[Solution Description] Hydrogen bonds show directional properties that enable them to stabilize the formation of $\alpha$-helices and $\beta$-sheets by forming regular patterns between the carbonyl oxygen and amide hydrogen in the protein backbone. This directional bonding leads to stable intermolecular attractions, maintaining these structural motifs.

Key Concept: Bond Stability, Octet Rule Limitations

b) Sulfur hexafluoride ($SF_6$)

[Solution Description] In sulfur hexafluoride ($SF_6$), sulfur forms six bonds, leading to a total of 12 electrons surrounding it, violating the octet rule. Despite this, the high electronegativity of fluorine stabilizes the structure through strong electrostatic interactions, making the compound stable. The calculation of electrovalency shows zero net charge, as each fluorine gains one electron while the sulfur remains neutral overall.

Key Concept: Real-World Application, Hybridization and Octet

c) Forms stable octahedral complex with $sp^3d^2$ hybridization.

[Solution Description]

$AlCl_3$ possesses an incomplete octet on the aluminum atom which renders it highly reactive with water. When $AlCl_3$ hydrates, it reacts with water to form hexahydrate $Al(H_2O)_6^{3+}$, alleviating its electron deficiency by accepting additional electron pairs from water molecules.

During this reaction, the hybridization of Al changes as follows:

– In $AlCl_3$, the aluminum atom originally has $sp^2$ hybridization due to forming three $\sigma$ bonds with chlorine atoms.

– Upon hydration forming $Al(H_2O)_6^{3+}$, the aluminum atom undergoes hybridization to $sp^3d^2$, accommodating six coordinate bonds with water molecules.

Therefore, the incomplete octet drives $AlCl_3$’s interaction with water, leading to a significant change in hybridization.

Key Concept: Comparative Analysis, Ionic Bond Strength Factors

b) MgO, because of double the ionic charges compared to LiF

[Solution Description] According to Coulomb’s Law, the force of attraction increases with higher charge magnitudes. MgO consists of +2 and -2 ions while LiF consists of +1 and -1 ions. This results in larger electrostatic forces for MgO, indicating stronger ionic interactions compared to LiF.

Key Concept: Basic Concept

c) One s and two p orbitals

[Solution Description] In $sp^2$ hybridisation, one s orbital and two p orbitals combine to form three equivalent $sp^2$ hybridised orbitals. This involves mixing one 2s orbital and two 2p orbitals.

Key Concept: Comparative Analysis, Advanced Resonance Impact

b) All bonds become equivalent with lengths between single and double bonds

[Solution Description]

The carbonate ion, $CO_3^{2-}$, can be represented by three resonance structures where the double bond between carbon and oxygen shifts among the three oxygens. In reality, due to resonance, these bonds are not purely single or double but have equal bond length characterized by a fractional bond order.

For $CO_3^{2-}$:

– Without resonance, we would expect some bonds to be single and one to be a double, leading to different lengths.

– With resonance, the bond order averages out with a value between that of a single and double bond.

Fractional bond order due to resonance leads to intermediate bond lengths in all three C-O bonds rather than distinct single or double bond lengths, thus making them shorter than a typical C-O single bond and longer than a typical C-O double bond.

Key Concept: Comparative Analysis

b) Structure II

[Solution Description] To determine the stability based on formal charges, we calculate the formal charge for each atom in different resonance structures and choose the one with the least separation of charge. Typically, the most stable structure has the lowest magnitude of formal charges and/or places negative charges on more electronegative atoms.

Key Concept: Energy Compensation

c) By the release of energy due to orbital overlap

[Solution Description] In $CH_4$, when carbon is excited from $[He] 2s^2 2p^2$ to $[He] 2s^1 2p_x^1 2p_y^1 2p_z^1$, the energy needed for this process is compensated by the release of energy during the formation of bonds. Specifically, when the sp3 hybrid orbitals overlap with hydrogen’s 1s orbitals, the energy released through these overlaps compensates the initial energy requirement for excitation.

Key Concept: Excited State Configuration

b) $3s^13p^33d^2$

[Solution Description] In $SF_6$, the ground state outer electronic configuration of sulfur is $3s^23p^4$. In the excited state, one electron from the 3p orbital is promoted to the 3d orbital, making it: $3s^13p^33d^2$.

Key Concept: Real-world Application, Factors Affecting Lattice Enthalpy

b) SrO has higher lattice enthalpy due to smaller ionic radius

[Solution Description] High-temperature refractory materials require compounds with high thermal stability, often indicated by substantial lattice enthalpy. SrO and BaO are alkaline earth metal oxides where Sr and Ba differ in ionic size.

Factors affecting lattice enthalpy include:

– Smaller ions lead to stronger electrostatic attractions between them due to decreased internuclear distance, thus increasing lattice enthalpy.

– Higher charge density also enhances lattice enthalpy.

Considering these factors, $Sr^{2+}$ has a smaller ionic radius compared to $Ba^{2+}$, leading to a higher lattice enthalpy and better thermal stability.

Hence, SrO will have a higher lattice enthalpy than BaO, making it more suitable for use in high-temperature environments.

Key Concept: Molecular Orbital Formation

b) Addition of wave functions $\psi_A + \psi_B$

[Solution Description]

To form a bonding molecular orbital, the wave functions $\psi_A$ and $\psi_B$ from two atomic orbitals must be added together. This results in constructive interference which increases electron density between the two nuclei.

Key Concept: Bond Formation Process

c) One sigma and one pi bond

[Solution Description] In a double bond, one sigma bond and one pi bond are present. The sigma bond forms first via end-to-end overlap, providing strong bonding; the pi bond complements this with sidewise overlap.

Key Concept: Comparative Analysis, Ionic Bond Exceptions

c) Ionic bonds have higher bond enthalpy compared to covalent bonds, except in cases like $NH_4^+$.

[Solution Description] Generally, ionic bonds have higher melting and boiling points, indicating stronger interactions in solid state due to lattice structures, but individual bond enthalpies in gaseous phase can vary. Covalent bonds might show higher specific bond enthalpy due to directional bonding. The exception with ammonium ($NH_4^+$), while classified as ionic, shows some covalent character affecting direct comparison. Hence, the correct statement considers general trends and exceptions.

Key Concept: Hydrogen Bonding and Molecular Geometry, Real-world Application

c) It decreases the density of ice making it float

[Solution Description] Ice is less dense than liquid water because of the way hydrogen bonds arrange the water molecules into a crystalline structure. At freezing temperatures, hydrogen bonds become more rigid, leading to an open hexagonal lattice.

This structure leaves more space between the molecules compared to the dynamic, constantly breaking and reforming hydrogen bonds in liquid water, where molecules can pack closer together. Thus, the increased volume of ice compared to the same mass of liquid water results in a lower density, allowing ice to float.

The correct understanding of this structural behavior leads to the conclusion:

Key Concept: Energy Comparison

c) Bonding molecular orbital

[Solution Description] Bonding molecular orbitals have lower energy than antibonding molecular orbitals. This is because bonding molecular orbitals result from constructive interference between atomic orbitals, leading to greater electron density between nuclei, thereby lowering the system’s total energy.

Key Concept: VSEPR Theory Application

c) Seesaw

[Solution Description] According to VSEPR theory, the geometry around a central atom with four bond pairs and one lone pair (represented by $AB_4E$) results in a seesaw shape. The presence of a lone pair affects the ideal trigonal bipyramidal arrangement by distorting it into a seesaw configuration due to stronger repulsion from the lone pair.

Key Concept: Hybridization and Shape

b) $sp^2$

[Solution Description] Trigonal planar geometry is achieved through $sp^2$ hybridization, where one s and two p orbitals mix to form three equivalent $sp^2$ hybrid orbitals that arrange themselves at 120-degree angles.

Key Concept: Advanced Bonding Concepts, Bond Length and Order

d) Same as $N_2$

[Solution Description] Each $X$ atom has 7 electrons, making $X_2$ similar to $N_2$ with a total of 14 electrons. The electronic configuration for $N_2$ indicates a bond order of 3:

$\sigma 1s^2 \sigma^* 1s^2 (\sigma 2s)^2 \sigma^* 2s^2 \pi_{2p_y}^2 \sigma_{2p_z}^2$​

Bond Order = $\frac{1}{2} (10 – 4) = 3$

A bond order of 3 suggests a very strong bond, typically resulting in shorter bond lengths when compared to lower bond orders found in diatomic molecules like $O_2$ or $H_2$. Thus, compared to these molecules, $X_2$ should have a shortest bond length similar to $N_2$.

Key Concept: Resonance Structures, Hydrogen Bonding

c) Nitrous acid ($\text{HNO}_2$)

[Solution Description]

To determine resonance, we look for structures that can be represented by more than one Lewis structure. For hydrogen bonding, hydrogen must be bonded to highly electronegative atoms like F, O, or N.

– Water ($\text{H}_2\text{O}$) has hydrogen bonding but no resonance.

– Carbon dioxide ($\text{CO}_2$) has resonance but no hydrogen.

– Nitrous acid ($\text{HNO}_2$) shows both characteristics: it can exist with different placements of double bonds between nitrogen and oxygen, indicating resonance, and the presence of OH group allows for hydrogen bonding.

– Methane ($\text{CH}_4$) has neither characteristic.

Therefore, nitrous acid ($\text{HNO}_2$) exhibits both.

Key Concept: Sigma and Pi Bonds, Hybridization

b) 3 $\sigma$ bonds, 2 $\pi$ bonds; $sp$ hybridization

[Solution Description] Acetylene ($C_2H_2$) consists of a triple bond between the two carbon atoms and a single bond between each carbon and hydrogen atom. Each carbon atom forms one sigma bond with hydrogen and shares two additional pi bonds with the other carbon atom. Therefore, there is a total of 3 $\sigma$ bonds and 2 $\pi$ bonds. For these bonds, each carbon atom undergoes sp hybridization involving the mixing of one s and one p orbital to create two linearly arranged sp orbitals for forming sigma bonds, while the leftover p orbitals participate in forming the $\pi$ bonds.

Key Concept: Real-World Applications, Advanced Magnetic Properties

b) $\text{O}_2$

[Solution Description] Paramagnetism occurs when there are unpaired electrons in molecular orbitals. For the given options, we analyze their bond order and electron configuration:

– **$\text{O}_2$:** Total of 12 valence electrons, configuration ends with two unpaired electrons in $\pi^*$ orbitals. Bond order is 2.

– **$\text{N}_2$:** Total of 10 valence electrons, all paired, bond order is 3.

– **$\text{F}_2$:** Total of 14 valence electrons, all paired, bond order is 1.

– **$\text{H}_2$:** Total of 2 valence electrons, all paired, bond order is 1.

Since $\text{O}_2$ has unpaired electrons, it exhibits paramagnetism.

Key Concept: Hybridization in Complex Molecules, Hybridization and Real-world Application

d) Lone pairs occupy axial positions, distorting shape

[Solution Description] $XeF_4$ involves xenon as the central atom, using its 5p orbitals along with available 5d orbitals for hybridization, forming $sp^3d^2$ hybridized orbitals allowing for an octahedral arrangement. However, two of these positions are occupied by lone pairs rather than bonding pairs. Thus, while the electron pair geometry remains octahedral, the molecular geometry observed becomes square planar because only four bonding pairs are directed towards fluorine atoms, minimizing repulsion mainly between bonding pairs and excluding lone pairs.