Chemical Bonding: A Detailed Overview

Introduction Chemical bonding is the attractive force that holds atoms together in a molecule or crystal. Atoms bond to achieve a more stable electronic configuration, typically a noble gas configuration (octet rule). The stability gained through bonding is reflected in the release of energy during bond formation.

I. Kössel-Lewis Theory of Chemical Combination

This theory explains chemical bonding based on the tendency of atoms to achieve a stable octet (eight electrons) in their outermost shell, or a duet for hydrogen. This can happen in two primary ways:

1. Transfer of electrons (Ionic Bond): Occurs between a metal (electropositive) and a non-metal (electronegative). The metal atom loses electrons to form a cation, and the non-metal atom gains electrons to form an anion. The electrostatic attraction between these oppositely charged ions forms the ionic bond.
   • Example: Formation of NaCl. Na (electron donor) loses 1 electron to become Na+. Cl (electron acceptor) gains 1 electron to become Cl-. Na+ and Cl- are held together by electrostatic forces.
   • Lewis Electron-Dot Symbols: Valence electrons are represented by dots around the element symbol.
     • Li (2s^1) -> Li.
     • F (2s^2 2p^5) -> :F:
2. Mutual sharing of electrons (Covalent Bond): Occurs typically between two non-metal atoms. Each atom contributes electrons to form shared pairs, resulting in both atoms achieving a stable octet (or duet for H).
   • Single Bond: Sharing of two electrons (one pair). E.g., Cl2 (Cl-Cl)
   • Double Bond: Sharing of four electrons (two pairs). E.g., O2 (O=O)
   • Triple Bond: Sharing of six electrons (three pairs). E.g., N2 (N≡N)

Exceptions to the Octet Rule:

• Incomplete Octet: Central atom has fewer than eight electrons. E.g., BeCl2 (Be has 4 electrons), BF3 (B has 6 electrons).
• Odd-Electron Molecules: Molecules with an odd number of valence electrons cannot satisfy the octet rule for all atoms. E.g., NO (Nitrogen has 7 electrons in its valence shell).
• Expanded Octet: Elements in the third period and beyond can accommodate more than eight electrons due to the availability of vacant d-orbitals. E.g., PCl5 (P has 10 electrons), SF6 (S has 12 electrons).

Coordinate Covalent Bond (Dative Bond): A type of covalent bond where both shared electrons are contributed by only one atom (the donor), while the other atom (the acceptor) accepts the pair. Represented by an arrow (->). Once formed, it behaves like a regular covalent bond.

• Example: H3N -> BF3. Nitrogen donates a lone pair to boron.

Writing Lewis Structures:

1. Calculate total valence electrons. Add negative charges for anions, subtract positive charges for cations.
2. Draw a skeletal structure, connecting atoms with single bonds. The least electronegative atom is usually central.
3. Distribute remaining electrons as lone pairs to satisfy octets of outer atoms (except H).
4. If the central atom lacks an octet, convert lone pairs from outer atoms into multiple bonds.

Formal Charge: Formal Charge = (Valence electrons on free atom) – (Non-bonding electrons) – (1/2 * Bonding electrons)

• The sum of formal charges in a molecule/ion equals its overall charge.
• Preferred Lewis structures minimize formal charges and place negative formal charges on more electronegative atoms.

II. Polarity of Bonds and Molecules

Electronegativity: The tendency of an atom to attract a shared pair of electrons towards itself in a covalent bond.

• In homonuclear diatomic molecules (e.g., H2, O2), electrons are shared equally.
• In heteronuclear diatomic molecules (e.g., HCl, HF), the more electronegative atom pulls the shared pair closer, creating partial negative (δ-) and partial positive (δ+) charges. This is a polar covalent bond.

Dipole Moment (μ): A measure of bond polarity or overall molecular polarity. It’s a vector quantity. μ = q * r (where q is charge separation, r is distance).

• Unit: Debye (D). 1 D = 10^-18 esu cm = 3.33 * 10^-30 C m.
• Bond Moment: Dipole moment of an individual bond.
• Molecular Dipole Moment: Vector sum of all bond moments in a molecule.
  • Non-polar molecules can have polar bonds if the bond moments cancel due to symmetry (e.g., CO2, BF3, CCl4).
  • H2O and NH3 are polar due to their bent/pyramidal shapes and non-cancelling bond moments/lone pair contributions.

Percent Ionic Character: Percent ionic character = (μ_observed / μ_ionic (calculated for 100% ionic)) * 100

III. Resonance

When a single Lewis structure cannot adequately describe the properties of a molecule (e.g., bond lengths are intermediate between single and double bonds), resonance is invoked. The actual structure is a resonance hybrid of two or more contributing (resonating) structures, which differ only in the placement of electrons, not atoms.

• Represented by a double-headed arrow (<->).
• Example: O3, SO2, CO3^2-

IV. Ionic Bond Formation and Lattice Energy

Formation of Solid Ionic Compound: The formation of an ionic compound (e.g., NaCl) from its constituent elements is a multi-step process, which can be analyzed using the Born-Haber Cycle.

1. Sublimation of metal: Na(s) -> Na(g) (ΔH1)
2. Dissociation of non-metal: 1/2 Cl2(g) -> Cl(g) (ΔH2)
3. Ionization of metal: Na(g) -> Na+(g) + e- (ΔH3, Ionization Energy)
4. Electron gain by non-metal: Cl(g) + e- -> Cl-(g) (ΔH4, Electron Affinity)
5. Formation of crystal lattice: Na+(g) + Cl-(g) -> NaCl(s) (ΔH5, Lattice Energy) The overall enthalpy change of formation (ΔHf) = ΔH1 + ΔH2 + ΔH3 + ΔH4 + ΔH5.

• For a stable ionic compound, ΔHf must be negative.
• Lattice Energy: The energy released when one mole of an ionic crystal is formed from its gaseous ions, or the energy required to completely separate one mole of a solid ionic compound into its gaseous constituent ions. A higher lattice energy indicates greater stability.
  • Factors affecting lattice energy: Charge on ions (higher charge -> higher lattice energy), size of ions (smaller ions -> higher lattice energy).

Fajan’s Rules: Explain the partial covalent character in ionic compounds.

• Small cation, large anion: Favors covalent character (cation has high polarizing power, anion is easily polarizable).
• High charge on cation/anion: Favors covalent character.
• Cation with pseudo-noble gas configuration ((n-1)d^10 ns^0): Has higher polarizing power than a cation with noble gas configuration (ns^2 np^6) of similar size and charge. (e.g., Cu+ vs Na+)

V. Valence Shell Electron Pair Repulsion (VSEPR) Theory

Predicts the shapes of molecules based on minimizing repulsion between electron pairs (both bonding and lone pairs) in the valence shell of the central atom.

• Main Postulates:
  • Electron pairs arrange themselves to be as far apart as possible.
  • Repulsion order: Lone pair-Lone pair (lp-lp) > Lone pair-Bond pair (lp-bp) > Bond pair-Bond pair (bp-bp).
  • Multiple bonds are treated as a single “super-pair” for VSEPR purposes.

VI. Valence Bond (VB) Theory

Describes covalent bonds as resulting from the overlap of atomic orbitals.

• Sigma (σ) Bond: Formed by head-on (end-to-end) overlap of atomic orbitals (s-s, s-p, p-p (along internuclear axis)). Electron density is concentrated along the internuclear axis.
• Pi (π) Bond: Formed by sideways overlap of parallel p-orbitals (or d-orbitals). Electron density is concentrated above and below the internuclear axis. A double bond consists of one σ and one π bond. A triple bond consists of one σ and two π bonds.

Hybridization: The concept of mixing atomic orbitals of similar energy to form new, degenerate hybrid orbitals that are more effective in forming bonds.

• sp^3 Hybridization: One s and three p orbitals mix to form four sp^3 hybrid orbitals. (Tetrahedral geometry, 109.5° bond angle). E.g., CH4, NH3, H2O.
• sp^2 Hybridization: One s and two p orbitals mix to form three sp^2 hybrid orbitals, with one unhybridized p orbital. (Trigonal planar geometry, 120° bond angle). E.g., BF3, C2H4.
• sp Hybridization: One s and one p orbital mix to form two sp hybrid orbitals, with two unhybridized p orbitals. (Linear geometry, 180° bond angle). E.g., BeCl2, C2H2.
• Hybridization involving d-orbitals: Occurs for elements in third period and beyond.
  • dsp^2 / sp^2d: Four hybrid orbitals (Square planar geometry, 90° bond angle). E.g., [Ni(CN)4]^2-
  • sp^3d / dsp^3: Five hybrid orbitals (Trigonal bipyramidal geometry). E.g., PCl5.
  • sp^3d^2 / d^2sp^3: Six hybrid orbitals (Octahedral geometry, 90° bond angle). E.g., SF6, [Co(NH3)6]^3+.

VII. Molecular Orbital (MO) Theory

Explains bonding by combining atomic orbitals to form molecular orbitals (MOs), which electrons occupy.

• Atomic Orbitals (AOs): Monocentric (around one nucleus).
• Molecular Orbitals (MOs): Polycentric (around multiple nuclei), belonging to the molecule as a whole.
• Types of MOs:
  • Bonding Molecular Orbitals (BMOs): Formed by constructive interference of AOs, lower in energy, stabilize the molecule.
  • Antibonding Molecular Orbitals (ABMOs): Formed by destructive interference of AOs, higher in energy, destabilize the molecule. Denoted with an asterisk (*).
• Rules for filling MOs: Aufbau principle, Pauli’s exclusion principle, Hund’s rule.

Order of Filling MOs (Energy Levels):

• For O2, F2, Ne2: σ 1s < σ1s < σ 2s < σ2s < σ 2pz < (π 2px = π 2py) < (π2px = π2py) < σ*2pz
• For Li2, Be2, B2, C2, N2: (due to s-p mixing) σ 1s < σ1s < σ 2s < σ2s < (π 2px = π 2py) < σ 2pz < (π2px = π2py) < σ*2pz

Molecular Characteristics from MO Theory:

1. Stability: Stable if Nb (bonding electrons) > Na (antibonding electrons).
2. Bond Order (BO): BO = 1/2 (Nb – Na). Positive BO indicates a stable molecule. Higher BO means stronger, shorter bond.
3. Magnetic Nature:
   • Diamagnetic: All electrons are paired.
   • Paramagnetic: Contains one or more unpaired electrons.

50 Important MCQs on Chemical Bonding

1. Which of the following elements can form a compound with an expanded octet? (a) Carbon (b) Nitrogen (c) Oxygen (d) Sulfur
2. According to VSEPR theory, the shape of the XeF4 molecule is: (a) Tetrahedral (b) Square pyramidal (c) Square planar (d) Octahedral
3. The type of hybridization in the central atom of PCl5 (gaseous phase) is: (a) sp^3 (b) sp^3d (c) sp^3d^2 (d) dsp^3
4. Which of the following molecules has a zero dipole moment? (a) H2O (b) NH3 (c) CCl4 (d) SO2
5. Which of the following shows the correct order of bond order? (a) O2^2- < O2^- < O2 < O2^+ (b) O2^+ < O2 < O2^- < O2^2- (c) O2^2- < O2 < O2^- < O2^+ (d) O2^- < O2^2- < O2 < O2^+
6. Which of the following is an example of an odd-electron molecule? (a) CO2 (b) NO2 (c) N2O (d) CH4
7. The bond angle in water (H2O) is approximately 104.5° due to: (a) sp^2 hybridization (b) Larger size of oxygen atom (c) Lone pair-bond pair repulsions (d) Presence of hydrogen bonding
8. Which of the following compounds has a coordinate covalent bond? (a) HCl (b) NaCl (c) NH4^+ (d) CH4
9. According to Fajan’s rules, covalent character is favored by: (a) Large cation and large anion (b) Small cation and small anion (c) Small cation and large anion (d) Large cation and small anion
10. The phenomenon of two or more structures being possible for a molecule, none of which perfectly describes the molecule, is known as: (a) Isomerism (b) Hybridization (c) Resonance (d) Allotropy
11. What is the formal charge on the oxygen atom in the carbonate ion (CO3^2-), considering one C=O bond and two C-O^- bonds? (a) 0 for C=O, -1 for C-O^- (b) -1 for C=O, 0 for C-O^- (c) +1 for C=O, -1 for C-O^- (d) 0 for all oxygen atoms
12. Which of the following species is diamagnetic? (a) O2 (b) B2 (c) C2 (d) NO
13. The lattice energy of an ionic compound primarily depends on: (a) Electronegativity difference (b) Ionization energy (c) Charge and size of ions (d) Electron affinity
14. Which of the following overlap is responsible for π (pi) bond formation? (a) s-s overlap (b) s-p overlap (c) p-p sideways overlap (d) p-p head-on overlap
15. The correct order of bond strength is: (a) Single > Double > Triple (b) Triple > Double > Single (c) Double > Single > Triple (d) Single = Double = Triple
16. What is the hybridization of carbon atoms in acetylene (C2H2)? (a) sp^3 (b) sp^2 (c) sp (d) dsp^2
17. Which theory predicts that electron pairs around the central atom arrange themselves to minimize repulsion? (a) Valence Bond Theory (VBT) (b) Molecular Orbital Theory (MOT) (c) VSEPR Theory (d) Crystal Field Theory (CFT)
18. The geometry around the central atom in SF6 is: (a) Trigonal bipyramidal (b) Square pyramidal (c) Octahedral (d) Tetrahedral
19. Which of the following has the highest dipole moment? (a) CO2 (b) BF3 (c) NH3 (d) CH4
20. The non-existence of He2 molecule can be explained by: (a) VSEPR theory (b) Molecular orbital theory (c) Valence bond theory (d) Octet rule
21. Which of the following bonds would be the most polar? (a) C-H (b) O-H (c) N-H (d) F-H
22. The percentage of p-character in sp^2 hybridized orbitals is: (a) 25% (b) 33.33% (c) 50% (d) 66.67%
23. Which of the following species has the smallest bond angle? (a) CH4 (b) NH3 (c) H2O (d) BF3
24. The total number of σ (sigma) bonds in ethane (C2H6) is: (a) 1 (b) 2 (c) 6 (d) 7
25. Which of the following processes involves the breaking of an ionic bond? (a) NaCl(s) -> NaCl(g) (b) NaCl(s) -> Na+(g) + Cl-(g) (c) NaCl(g) -> Na(g) + Cl(g) (d) NaCl(aq) -> Na+(aq) + Cl-(aq)
26. Which statement about resonance structures is TRUE? (a) Resonance structures are in rapid equilibrium with each other. (b) The true structure is an average of the resonance structures. (c) Resonance structures have different atomic arrangements. (d) Resonance structures always have different energy levels.
27. The increasing order of bond angle for H2O, H2S, H2Se, and H2Te is: (a) H2O < H2S < H2Se < H2Te (b) H2Te < H2Se < H2S < H2O (c) H2S < H2Se < H2Te < H2O (d) H2Te < H2O < H2S < H2Se
28. The highest electron density in a π bond is: (a) Along the internuclear axis (b) Above and below the internuclear axis (c) At the nucleus (d) In the hybrid orbitals
29. Which of the following is a flexidentate ligand? (a) Ethylene diamine (b) Oxalate (c) EDTA (d) Pyridine
30. The formal charge on sulfur in the sulfate ion (SO4^2-), assuming all S-O bonds are single bonds, is: (a) +2 (b) +1 (c) 0 (d) -2
31. Which of the following molecules has a linear shape? (a) SO2 (b) H2O (c) CO2 (d) NF3
32. The process of formation of an ionic compound involves the sum of various energy changes, including lattice energy, ionization energy, and electron affinity. This cycle is known as: (a) Hess’s Law (b) Born-Haber Cycle (c) Carnot Cycle (d) Standard Enthalpy Cycle
33. Which of the following is the correct order of increasing covalent character? (a) LiCl < NaCl < KCl (b) KCl < NaCl < LiCl (c) LiCl = NaCl = KCl (d) LiCl > KCl > NaCl
34. The bond order of NO^+ is: (a) 1 (b) 2 (c) 2.5 (d) 3
35. Hydrogen bonding is strongest in: (a) H2S (b) NH3 (c) H2O (d) HF
36. Which of the following factors does NOT influence the dipole moment of a molecule? (a) Electronegativity difference between bonded atoms (b) Molecular geometry (c) Presence of lone pairs (d) Atomic number of the central atom
37. The number of lone pairs on the central atom of XeF2 is: (a) 1 (b) 2 (c) 3 (d) 4
38. Which of the following contains both σ and π bonds? (a) C2H6 (b) C2H4 (c) CH4 (d) H2O
39. The structure of white phosphorus (P4) consists of: (a) Linear chain (b) Planar square (c) Tetrahedral unit (d) Cyclic ring
40. Which of the following statements about hybridization is correct? (a) It involves mixing of atomic orbitals of different energy levels. (b) It always forms bonds that are weaker than pure atomic orbitals. (c) The number of hybrid orbitals formed is equal to the number of atomic orbitals mixed. (d) Hybrid orbitals are always degenerate in energy.
41. The bond order of the F2 molecule is: (a) 0 (b) 1 (c) 1.5 (d) 2
42. Which of the following molecules has a square planar geometry? (a) SF4 (b) BrF5 (c) XeF4 (d) PCl5
43. Which of the following has the highest melting point? (a) I2 (b) CCl4 (c) SiC (d) H2O
44. In which of the following pairs is the hybridization of the central atom NOT the same? (a) BF3, BCl3 (b) NH3, PH3 (c) H2O, H2S (d) CO2, SO2
45. The concept of back bonding is observed in: (a) NH3 (b) BF3 (c) PH3 (d) H2O
46. Which of the following statements about the bond angles in NH3, H2O, and CH4 is correct? (a) CH4 > NH3 > H2O (b) H2O > NH3 > CH4 (c) NH3 > H2O > CH4 (d) CH4 < NH3 < H2O
47. The term “isoelectronic” refers to species having: (a) The same number of protons (b) The same number of neutrons (c) The same number of electrons (d) The same atomic mass
48. The stability of a molecule, according to MO theory, increases when: (a) Nb < Na (b) Nb = Na (c) Nb > Na (d) Nb + Na is high
49. Which of the following is an intermolecular force of attraction? (a) Covalent bond (b) Ionic bond (c) Hydrogen bond (d) Metallic bond
50. The bond order of the C2 molecule is: (a) 1 (b) 2 (c) 2.5 (d) 3

The Answer Keys