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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:
Same or nearly the same energy.
High symmetry about the molecular axis.
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).