In-Depth Notes on Chemical Bonding and Lewis Structures

Chemical Bonding I: The Lewis Model covers the fundamental concepts of chemical bonds used in molecules.

Three Types of Bonds:
1. Ionic Bonding: Involves electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions), forming a crystal lattice.
- Cations: Typically derived from metals with low ionization energy.
- Anions: From nonmetals; can be monoatomic (e.g., Cl⁻, O²⁻) or polyatomic (e.g., NO₃⁻, SO₄²⁻).
- Examples: NaCl, MgO, KNO₃, CuSO₄, Li₂CO₃, Ca₃(PO₄)₂.
1. Covalent Bonding: Sharing of electron pairs between two nonmetals to achieve an octet (eight electrons).
- Exceptions exist for some elements (e.g., H requires only 2 electrons).
1. Metallic Bonding: Metals share electrons in a 'sea' of electrons surrounded by positively charged nuclei, contributing to conductivity.

Valence Electrons: Electrons in the outermost principal energy level, represented by dots in Lewis structures.
- Electron Configuration for Selected Elements (Noble gas notation):
  - Li: 1s² 2s¹
  - Be: 1s² 2s²
  - B: 1s² 2s² 2p¹
  - C: 1s² 2s² 2p²
  - N: 1s² 2s² 2p³
  - O: 1s² 2s² 2p⁴
  - F: 1s² 2s² 2p⁵
  - Ne: 1s² 2s² 2p⁶

Lewis Structures & Ionic Bonding:
- Example of LiBr:
  - Formation: Li∙ + :Br: → Li⁺[:Br:]
- Example of MgO:
  - Formation: ∙Mg∙ + :O: → Mg²⁺[:O:]

Lattice Energy: Energy involved in forming one mole of ionic solid from its gaseous ions, always exothermic (negative value).
- Born-Haber Cycle: A thermochemical cycle for calculating lattice energy.
- Consider CsF:
  - Formation: Cs⁺(g) + F⁻(g) → CsF(s)
- Important process steps include sublimation, ionization energy, and electron affinity adjustments for ions formed.

Steps in the Born-Haber Cycle:
- Reactions show transformations from solid states to ions.
- CsF Cycle:
1. Sublimation of Cs, ΔH₁ = 76.5 kJ/mol (endothermic).
2. Ionization of Cs, ΔH₂ = 375.7 kJ/mol (endothermic).
3. Formation of F(g), ΔH₃ = 79.4 kJ/mol (endothermic).
4. Electron affinity for F, ΔH₄ = -328.2 kJ/mol (exothermic).
5. Total lattice energy for CsF, ΔH₅ = -756.9 kJ/mol (exothermic).

Formation Reaction for CaO:
- Ca⁺(s) + ½ O₂(g) → CaO(s).
- 8 steps including sublimation and ionization.
- Final total lattice energy for CaO = -3498.7 kJ/mol.

Coulomb’s Law: Related to lattice energy.
- Formula:
  $E = \frac{1}{4\pi \varepsilon} \frac{q<em>1 q</em>2}{r}$
- Trends include size of ions affecting lattice energy due to distance.
- Charge product on ions impacts energies.

Strength of electrostatic forces leads to high melting points.
- Example: CsF has m.p. 703°C, CaO has m.p. 2613°C.
- Non-directional attraction in ionic solids.

Lewis Structures: Helps visualize electron sharing in covalent bonds.
- Relevance of single, double, and triple bonds.

Bonded Pairs: Shared electrons forming bonds.
- Lone Pairs: Non-bonding electrons on an atom.
- Example: Water has two bonded pairs and two lone pairs.

Carbon's Valence Electrons:
- Forms stable covalent bonds due to four valence electrons.
- Examples of carbon compounds: choose bonds for achieving octet:
- Methane (CH₄), Acetaldehyde (H₂CO), Carbon dioxide (CO₂).

Definition: Carbon's ability to bond with itself leading to extensive compound variety (organic chemistry).
- Examples include octane, benzene, and other hydrocarbons.

Electronegativity (EN): Atom's ability to attract electrons in a bond. Linus Pauling's scale (F being highest).
- Trends: EN increases across a period, decreases down a group.

Electronegativity Difference (ΔEN): Determines bond type.
- Nonpolar covalent: ΔEN = 0.0-0.4; Polar covalent: ΔEN = 0.4-2.0; Ionic: ΔEN > 2.0.

Dipole Moment (μ): Quantifies bond polarity, calculated by formula: $μ = qr$ .
- Percent ionic character is computed to determine how ionic a bond is based on observed dipoles.

Steps:
1. Write skeletal structure.
2. Count total valence electrons.
3. Distribute electrons to satisfy octets.
4. Form double/triple bonds if needed.

Definition: Multiple valid Lewis structures for the same molecule.
- Guidelines ensure the same number of electrons and fixed atom positions.

Formula: $q = v - ℓ - ½ b$
- Where (q) is the formal charge, (v) is valence e⁻, (ℓ) is lone pair e⁻, and (b) is bonded pair e⁻.

Odd-electron species lead to free radicals (e.g., NO, NO₂).
Incomplete octets (< 8 e⁻) seen in some Group 13 and Group 2 elements (e.g., BF₃).
Expanded octets allow for more than 8 e⁻ in elements from periods 3-5 (e.g., PF₅).