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Chapter 4: Chemical Bonding and Molecular Structure

1. Lewis Dot Structures

• Lewis dot structures represent the valence electrons of an atom.

• Notation involves the element symbol surrounded by dots denoting valence electrons.

2. Covalent Bonds

• Formed by mutual sharing of electrons to complete octets/duplets.

• Single Covalent Bonds: One electron pair shared (e.g., H2).

• Double Covalent Bonds: Two electron pairs shared (e.g., O2).

• Triple Covalent Bonds: Three electron pairs shared (e.g., N2).

3. Octet Rule

• Developed by Kossel and Lewis (1916).

• Atoms combine by transferring or sharing valence electrons to achieve an octet.

4. Limitations of the Octet Rule

• Incomplete octet: Central atom has fewer than eight electrons (e.g., LiCl, BeCl2, BCl3).

• Odd-electron molecules: Molecules with an odd number of electrons don’t satisfy the octet rule (e.g., NO, NO2).

• Expanded octet: Atoms beyond the third period can have more than eight electrons (e.g., PF5, SF6).

• Does not account for molecular shape.

5. Ionic Bonds (Electrovalent Bonds)

• Result from electron transfer between atoms (electropositive to electronegative).

• More likely between elements with low ionization enthalpy and high electron gain enthalpy.

• Most ionic compounds consist of cations from metals and anions from non-metals.

6. Formation of Ionic Bonds

• M(g) → M+(g) + e ; Ionization Enthalpy

• X(g) + e → X−(g) ; Electron Gain Enthalpy

• M+(g) + X−(g) → MX(s) ; Formation of Ionic Compound

7. Bond Length

• Defined as the equilibrium distance between nuclei of two bonded atoms.

8. Bond Angle

• Angle between orbitals containing bonding electron pairs around the central atom.

• Indicates orbital distribution and helps determine molecular shape.

9. Bond Enthalpy

• Energy required to break one mole of specific bonds between two atoms in a gaseous state.

• Measured in kJ/mol.

10. Bond Order

• Number of bonds between two atoms in a molecule (e.g., Bond Order of O2 = 2).

• Generally, as bond order increases, bond enthalpy increases and bond length decreases.

11. Resonance

• Concept explaining that single Lewis structures may not fully describe a molecule.

• Multiple structures (canonical forms) are used to describe the hybrid more accurately.

12. Polarity of Bonds

• In heteronuclear molecules (e.g., HCl), shared electron pairs are displaced towards the more electronegative atom, resulting in a polar covalent bond.

13. Dipole Moment

• Result of polarization in a molecule, defined as the product of charge and distance between centers of positive and negative charge.

• Expressed as: Dipole moment (μ) = charge (Q) × distance (r).

14. VSEPR Theory

• Molecular shape is determined by valence shell electron pairs around the central atom.

• Electron pairs repel each other, minimizing repulsion and maximizing distance.

• Multiple bonds count as one electron pair.

• Order of repulsive interactions: Lone pair-lone pair > Lone pair-bond pair > Bond pair-bond pair.

Geometry of Molecules under VSEPR Theory

• Bond pairs | Lone pairs | Geometry Classes | Examples

• 2 | 0 | Linear | BeCl2

• 3 | 0 | Trigonal Planar | BCl3

• 2 | 1 | Bent | SO2

• 4 | 0 | Tetrahedral | CH4

• 3 | 1 | Trigonal Pyramidal | NH3, PH3

• 2 | 2 | V-shaped | H2O

• 5 | 0 | Trigonal Bipyramidal | PCl5

• 4 | 1 | See Saw | SF4

• 3 | 2 | T-shaped | ClF3

• 2 | 3 | Linear | XeF2

• 5 | 1 | Square Pyramidal | ClF5, IF5

• 4 | 2 | Square Planar | XeF4

• 6 | 0 | Distorted Octahedral | XeF6

15. Hybridization

• Intermixing of orbitals of different energies to form new orbitals of equivalent energies and shapes.

• Hybrid orbitals formed are equal to the number of atomic orbitals hybridized.

• Hybridized orbitals are effective for stable bonding and oriented to minimize repulsion.

16. Types of Hybridization

• sp Hybridization: Mixing of one s and one p orbital (e.g., BeCl2).

• sp2 Hybridization: Mixing of one s and two p orbitals (e.g., BCl3).

• sp3 Hybridization: Mixing of one s and three p orbitals (e.g., CH4).

17. Molecular Orbitals

• Result from the linear combination of atomic orbitals, describing electron probability distribution in a molecule.

18. Bonding Molecular Orbitals

• Formed by the addition overlap of atomic orbitals (same sign lobes) with lower energy promoting bonding.

19. Anti-bonding Molecular Orbitals

• Result from the subtraction overlap of atomic orbitals (opposite sign lobes) with higher energy, not favoring bonding.

20. Bond Order in Molecular Orbitals

• Defined as B.O. = 1/2 (Nb - N).

• Helps estimate the stability of a molecule; stability if Nb > N, instability if N > Nb, and instability if Nb = N.

21. Sigma (σ) Molecular Orbitals

• Formed from the overlap of s orbitals or the head-to-head overlap of s and p orbitals.

22. Pi (π) Molecular Orbitals

• Formed by lateral overlap of two parallel p orbitals.

23. Conditions for the Combination of Atomic Orbitals

• Combine if:

  1. Same or nearly the same energy.

  2. High symmetry about the molecular axis.

  3. Maximum overlap of orbitals.

24. Energy Level Diagrams for Molecular Orbitals

• The order of energies for O2 and F2 is: σ(1s) < σ*(1s) < σ(2s) < σ*(2s) < σ(2p) < π(2p)

• B, C2, N2 exhibit a different order where π(2p) is higher than σ(2p).