NCERT Solutions: Class 11 Chemistry Chapter 4 - Chemical Bonding and Molecular Structure

Atoms combine because the resulting species has lower energy than the separated atoms. Ionic bonds form by electron transfer followed by electrostatic attraction between ions. Covalent bonds form by sharing electron pairs so that atoms attain stable valence-shell configurations.

Lewis symbols show valence electrons. Mg is group 2, Na group 1, B group 13, O group 16, N group 15 and Br group 17. Therefore their Lewis symbols contain 2, 1, 3, 6, 5 and 7 dots respectively.

Sulphur has 6 valence electrons and gains 2 electrons to form S2- with 8 valence electrons. Aluminium has 3 valence electrons and loses them to form Al3+. Hydrogen has 1 electron and H- has 2 electrons, the helium-like duplet.

The rule is significant because it relates bonding to stable noble-gas configurations and helps draw Lewis structures. Its limitations include electron-deficient molecules such as BeH2, BF3 and AlCl3; odd-electron molecules such as NO and NO2; and hypervalent molecules such as PF5, SF6 and H2SO4 where the central atom has more than eight valence electrons.

A metal should lose electrons easily, so low ionization enthalpy helps. A non-metal should accept electrons readily, so a large negative electron gain enthalpy helps. The oppositely charged ions should release substantial energy on forming the crystal lattice, so high lattice enthalpy stabilizes the ionic compound.

VSEPR arranges electron pairs to minimise repulsion. BeCl2 has two bond pairs and no lone pair on Be, so linear. BCl3 has three bond pairs, so trigonal planar. SiCl4 has four bond pairs, so tetrahedral. AsF5 has five bond pairs, so trigonal bipyramidal. H2S has two bond pairs and two lone pairs on S, so bent. PH3 has three bond pairs and one lone pair on P, so trigonal pyramidal.

Electron pair repulsions follow lp-lp > lp-bp > bp-bp. NH3 has one lone pair and three bond pairs, giving a pyramidal molecule with bond angle about 107 degrees. H2O has two lone pairs and two bond pairs, so greater lone-pair repulsion reduces the H-O-H angle to about 104.5 degrees.

Higher bond order means more bonding electron pairs or greater net bonding interaction between atoms. Therefore triple bonds are generally stronger than double bonds, which are stronger than single bonds; bond length usually decreases as bond order increases.

At the equilibrium bond distance, attractive and repulsive forces balance and the molecule has minimum potential energy. This internuclear distance is called bond length.

In each canonical form of carbonate, one C-O bond is double and two C-O bonds are single with negative charge on oxygen atoms. Since any oxygen can form the double bond, three equivalent structures are possible. The real ion is not any one structure; it is a hybrid in which all C-O bonds are identical with bond order about 1.33 and the charge is delocalised over the three oxygen atoms.

Canonical resonance forms must have the same skeletal arrangement of atoms and differ only in the distribution of electrons. In the two shown structures of H3PO3, hydrogen is attached differently, changing atom connectivity. Therefore they are not resonance forms of the same molecule.

Potassium loses one electron, while sulphur gains two, so two K atoms transfer one electron each to S. Calcium loses two electrons and oxygen gains two. Aluminium loses three electrons and nitrogen gains three. The resulting oppositely charged ions are held by ionic bonds.

In CO2, the C=O dipoles are equal and opposite along a straight line. In H2O, lone pairs on oxygen make the molecule bent; the two O-H dipoles are at an angle and their vector sum is non-zero. This difference in dipole moment is consistent with linear CO2 and bent H2O.

A non-zero dipole moment indicates a polar molecule. Dipole moment values help distinguish linear and bent or symmetrical and unsymmetrical structures, estimate percentage ionic character of bonds, and compare polarity in related molecules.

Electronegativity is a relative, dimensionless property of bonded atoms and depends on bonding environment. Electron gain enthalpy is a thermodynamic quantity for isolated gaseous atoms, measured in kJ mol^-1.

In HCl, chlorine is more electronegative than hydrogen, so the shared electron pair is displaced toward chlorine. Chlorine becomes partially negative and hydrogen partially positive, making the H-Cl bond polar covalent.

Acetic acid contains a methyl group bonded to a carboxyl carbon. The carboxyl carbon forms one double bond to oxygen and one single bond to the hydroxyl oxygen; the hydroxyl oxygen is bonded to hydrogen. Each oxygen has two lone pairs in the Lewis structure.

Methane has four bond pairs and no lone pair around carbon. VSEPR predicts a tetrahedral arrangement for minimum bond-pair repulsion. Valence bond theory also gives sp3 hybridisation with orbitals directed toward the corners of a tetrahedron. Square planar geometry would not give equivalent maximum-overlap sp3 bonds for carbon.

The Be atom in BeH2 is sp hybridised and the molecule has a linear H-Be-H geometry. Although each Be-H bond is polar, the two bond dipoles act in opposite directions along the same line, giving zero net dipole moment.

Both molecules are trigonal pyramidal. In NH3, the N-H bond dipoles and the lone-pair contribution act in the same general direction, increasing the net dipole. In NF3, the N-F bond dipoles oppose the lone-pair contribution, so the net dipole moment is smaller.

In sp hybridisation, one s and one p orbital mix to form two equivalent orbitals oriented 180 degrees apart. In sp2 hybridisation, one s and two p orbitals form three equivalent orbitals in one plane at 120 degrees. In sp3 hybridisation, one s and three p orbitals form four equivalent orbitals directed tetrahedrally at about 109.5 degrees.

In AlCl3, aluminium forms three sigma bonds and is trigonal planar, corresponding to sp2 hybridisation. On accepting a lone pair from Cl-, Al forms four sigma bonds in AlCl4-, which is tetrahedral and uses sp3 hybrid orbitals.

In BF3, boron is trigonal planar and sp2 hybridised. In the adduct F3B.NH3, nitrogen donates a lone pair to boron and boron becomes tetra-coordinate, so B becomes sp3 hybridised. Nitrogen in NH3 already has three bonds and one lone pair, so it remains approximately sp3.

In C2H2, H-C≡C-H has two C-H sigma bonds and one C-C sigma bond, plus two pi bonds in the triple bond. In C2H4, four C-H bonds and one C-C bond are sigma bonds, while the C=C double bond has one pi bond.

A sigma bond requires head-on overlap along the internuclear axis. With the x-axis as internuclear axis, 2py orbitals are perpendicular to the axis and overlap sideways, forming a pi-type interaction rather than a sigma bond. The other listed overlaps can occur along the internuclear axis.

Carbon with four single bonds is sp3 hybridised. Carbon in a C=C double bond or C=O carbonyl is sp2 hybridised. Applying these rules gives the listed hybridisations.

Bond pairs form covalent bonds and lie between bonded atoms. Lone pairs remain on one atom and influence shape through repulsion. In ammonia, nitrogen has three N-H bond pairs and one lone pair.

Sigma electron density is concentrated along the line joining the nuclei and can arise from s-s, s-p or p-p head-on overlap. Pi electron density lies above and below the internuclear axis from lateral p-p overlap. A double bond has one sigma and one pi bond; a triple bond has one sigma and two pi bonds.

Each hydrogen atom has one electron in a 1s orbital. As two H atoms approach, their 1s orbitals overlap and electron density increases between the two nuclei. If the electron spins are opposite, attraction overcomes repulsion and the system reaches minimum energy at the H-H bond length. This stable overlap forms the H-H sigma bond.

In LCAO, combining atomic orbitals should be close in energy so neither dominates completely. They must have the same symmetry with respect to the molecular axis so constructive/destructive combination is possible. Greater overlap gives more effective molecular orbital formation.

Each Be atom has configuration 1s2 2s2, so Be2 has eight electrons. Its MO configuration fills sigma1s, sigma*1s, sigma2s and sigma*2s equally. The number of bonding and antibonding electrons is equal, so bond order = (Nb - Na)/2 = 0. Therefore no net bond forms.

Using MO theory, bond orders are: O2+ = 2.5, O2 = 2, O2- = 1.5 and O2^2- = 1. Higher bond order means greater stability, giving the listed order. O2+ and O2- have one unpaired electron each, O2 has two unpaired electrons, and peroxide O2^2- has all electrons paired.

Atomic orbitals are wave functions and can have positive or negative phase. When orbitals of the same phase overlap, constructive interference gives bonding molecular orbitals. Opposite phases overlap destructively and give antibonding molecular orbitals.

Phosphorus forms five sigma bonds in PCl5 using sp3d hybrid orbitals arranged as a trigonal bipyramid. Three equatorial bonds lie in one plane at 120 degrees, while two axial bonds are perpendicular to that plane. Each axial bond has three 90-degree interactions with equatorial bonds, so it experiences greater repulsion and becomes longer and weaker than equatorial bonds.

Hydrogen bonding occurs when H is covalently bonded to F, O or N and becomes strongly partially positive. It can then interact with a lone pair on another electronegative atom. This interaction is directional and stronger than ordinary van der Waals forces.

In MO theory, bond order is (Nb - Na)/2. For N2, the net bonding count gives bond order 3. For O2, bond order is (10 - 6)/2 = 2. Removing one electron from antibonding O2 gives O2+ with bond order 2.5. Adding one electron to antibonding O2 gives O2- with bond order 1.5.