Chemical Bonding and Molecular Structure

Why bonds form? Atoms achieve stable octet (Lewis 1916). Inert gases have full outer shells → low reactivity.

TYPES OF BONDS:

1. IONIC BOND (electrostatic, complete e⁻ transfer)

• Between metal (loses e⁻) and non-metal (gains e⁻). E.g., NaCl, MgO.
• Properties: high m.p., conduct in molten/aqueous state, soluble in polar solvents.
• Lattice energy: energy to form 1 mol of ionic solid from gaseous ions. Higher = stronger bond.
• Born-Haber cycle to calculate lattice energy.

2. COVALENT BOND (e⁻ sharing)

• Between non-metals (e.g., H₂, Cl₂, H₂O).
• Single (1 pair), double (2 pairs), triple (3 pairs).
• Bond order = number of shared pairs.
• Bond length: C-C (154) > C=C (134) > C≡C (120) pm.
• Bond energy: C≡C > C=C > C-C.
• Polar covalent: unequal sharing due to electronegativity difference (e.g., H-Cl).

3. COORDINATE / DATIVE BOND

• Both electrons supplied by one atom. E.g., NH₃ → BF₃ adduct, NH₄⁺.

4. HYDROGEN BOND

• Weak (5-30 kJ/mol) attraction between H attached to F/O/N and a lone pair of another F/O/N.
• Explains: high b.p. of water, H₂O₂, HF; ice less dense than water; DNA base pairing.

5. METALLIC BOND

• "Sea of delocalized electrons" around positive metal cores. Explains: conductivity, malleability, ductility, luster.

LEWIS STRUCTURES — Rules

1. Count total valence electrons.
2. Draw skeleton (least EN atom usually central, except H).
3. Place bonds + lone pairs to satisfy octet.
4. If short, form double/triple bonds.
5. Check formal charges (minimize them).

VSEPR (Valence Shell Electron Pair Repulsion):

Predicts shape from number of e⁻ pairs (bonding + lone) around central atom.

Lone-pair distortion: lone pairs repel more than bond pairs.

• CH₄: 109.5° (perfect tetra). NH₃: 107° (1 LP). H₂O: 104.5° (2 LP).

HYBRIDIZATION:

Mixing atomic orbitals to form equivalent hybrid orbitals.

• sp (50% s, 50% p) — linear, 180°.
• sp² — trigonal planar, 120°.
• sp³ — tetrahedral, 109.5°.
• sp³d — TBP.
• sp³d² — octahedral.

Example — ethyne (C₂H₂): each C is sp; 2 sp orbitals form 1 σ to other C + 1 σ to H. Unhybridized 2p form two π bonds (C≡C triple).

MOLECULAR ORBITAL THEORY (MOT):

• Atomic orbitals combine to form bonding (σ, π) and antibonding (σ*, π*) MOs.
• Bond order = (Nb − Na) / 2.
• O₂: bond order 2, paramagnetic (2 unpaired e⁻ in π*).
• N₂: bond order 3, diamagnetic.
• He₂: bond order 0 → doesn't form.

FAJANS RULES — covalent character in ionic bond:

• Smaller cation → more polarizing → more covalent.
• Larger anion → more polarizable.
• Higher cation charge → more covalent.
• Cation with pseudo-noble-gas configuration (Cu⁺, Ag⁺, Zn²⁺) is more polarizing.

DIPOLE MOMENT (μ) = Q × d (Debye).

• BF₃ (μ=0, trigonal planar — vectors cancel).
• NH₃ (μ=1.46 D — pyramidal, doesn't cancel).
• H₂O (μ=1.85 D — bent).
• CO₂ (μ=0 — linear, cancels).

EXAM HOOKS:

• Strongest covalent: H-F. Strongest H-bond: F-H...F.
• Bent shape: 2 lone pairs (H₂O, OF₂); pyramidal: 1 LP (NH₃).
• sp³d hybridization → TBP (PCl₅, SF₄).
• Among NH₃, PH₃, AsH₃, SbH₃, BiH₃ → NH₃ has highest b.p. (H-bond).

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry by treating bonding pairs and lone pairs as regions of electron density that repel each other and arrange to maximize separation.

Steric number = number of bonded atoms + number of lone pairs on the central atom.

Lone pairs distort geometry. Lone-pair repulsion > bond-pair repulsion. Examples:

• NH₃ (3 bonds + 1 LP): tetrahedral electron geometry, but trigonal pyramidal molecular shape, bond angle 107° (compressed from 109.5°).
• H₂O (2 bonds + 2 LPs): tetrahedral electron geometry, bent molecular shape, bond angle 104.5°.

Lone-pair preference rule for steric 5 and 6: lone pairs occupy equatorial positions in trigonal bipyramidal (less crowded) and adjacent (cis) positions in octahedral.

Why bonds form? Atoms achieve stable octet (Lewis 1916). Inert gases have full outer shells → low reactivity.

TYPES OF BONDS:

1. IONIC BOND (electrostatic, complete e⁻ transfer)

• Between metal (loses e⁻) and non-metal (gains e⁻). E.g., NaCl, MgO.
• Properties: high m.p., conduct in molten/aqueous state, soluble in polar solvents.
• Lattice energy: energy to form 1 mol of ionic solid from gaseous ions. Higher = stronger bond.
• Born-Haber cycle to calculate lattice energy.

2. COVALENT BOND (e⁻ sharing)

• Between non-metals (e.g., H₂, Cl₂, H₂O).
• Single (1 pair), double (2 pairs), triple (3 pairs).
• Bond order = number of shared pairs.
• Bond length: C-C (154) > C=C (134) > C≡C (120) pm.
• Bond energy: C≡C > C=C > C-C.
• Polar covalent: unequal sharing due to electronegativity difference (e.g., H-Cl).

3. COORDINATE / DATIVE BOND

• Both electrons supplied by one atom. E.g., NH₃ → BF₃ adduct, NH₄⁺.

4. HYDROGEN BOND

• Weak (5-30 kJ/mol) attraction between H attached to F/O/N and a lone pair of another F/O/N.
• Explains: high b.p. of water, H₂O₂, HF; ice less dense than water; DNA base pairing.

5. METALLIC BOND

• "Sea of delocalized electrons" around positive metal cores. Explains: conductivity, malleability, ductility, luster.

LEWIS STRUCTURES — Rules

1. Count total valence electrons.
2. Draw skeleton (least EN atom usually central, except H).
3. Place bonds + lone pairs to satisfy octet.
4. If short, form double/triple bonds.
5. Check formal charges (minimize them).

VSEPR (Valence Shell Electron Pair Repulsion):

Predicts shape from number of e⁻ pairs (bonding + lone) around central atom.

Lone-pair distortion: lone pairs repel more than bond pairs.

• CH₄: 109.5° (perfect tetra). NH₃: 107° (1 LP). H₂O: 104.5° (2 LP).

HYBRIDIZATION:

Mixing atomic orbitals to form equivalent hybrid orbitals.

• sp (50% s, 50% p) — linear, 180°.
• sp² — trigonal planar, 120°.
• sp³ — tetrahedral, 109.5°.
• sp³d — TBP.
• sp³d² — octahedral.

Example — ethyne (C₂H₂): each C is sp; 2 sp orbitals form 1 σ to other C + 1 σ to H. Unhybridized 2p form two π bonds (C≡C triple).

MOLECULAR ORBITAL THEORY (MOT):

• Atomic orbitals combine to form bonding (σ, π) and antibonding (σ*, π*) MOs.
• Bond order = (Nb − Na) / 2.
• O₂: bond order 2, paramagnetic (2 unpaired e⁻ in π*).
• N₂: bond order 3, diamagnetic.
• He₂: bond order 0 → doesn't form.

FAJANS RULES — covalent character in ionic bond:

• Smaller cation → more polarizing → more covalent.
• Larger anion → more polarizable.
• Higher cation charge → more covalent.
• Cation with pseudo-noble-gas configuration (Cu⁺, Ag⁺, Zn²⁺) is more polarizing.

DIPOLE MOMENT (μ) = Q × d (Debye).

• BF₃ (μ=0, trigonal planar — vectors cancel).
• NH₃ (μ=1.46 D — pyramidal, doesn't cancel).
• H₂O (μ=1.85 D — bent).
• CO₂ (μ=0 — linear, cancels).

EXAM HOOKS:

• Strongest covalent: H-F. Strongest H-bond: F-H...F.
• Bent shape: 2 lone pairs (H₂O, OF₂); pyramidal: 1 LP (NH₃).
• sp³d hybridization → TBP (PCl₅, SF₄).
• Among NH₃, PH₃, AsH₃, SbH₃, BiH₃ → NH₃ has highest b.p. (H-bond).

Why bonds form? Atoms achieve stable octet (Lewis 1916). Inert gases have full outer shells → low reactivity.

TYPES OF BONDS:

1. IONIC BOND (electrostatic, complete e⁻ transfer)

• Between metal (loses e⁻) and non-metal (gains e⁻). E.g., NaCl, MgO.
• Properties: high m.p., conduct in molten/aqueous state, soluble in polar solvents.
• Lattice energy: energy to form 1 mol of ionic solid from gaseous ions. Higher = stronger bond.
• Born-Haber cycle to calculate lattice energy.

2. COVALENT BOND (e⁻ sharing)

• Between non-metals (e.g., H₂, Cl₂, H₂O).
• Single (1 pair), double (2 pairs), triple (3 pairs).
• Bond order = number of shared pairs.
• Bond length: C-C (154) > C=C (134) > C≡C (120) pm.
• Bond energy: C≡C > C=C > C-C.
• Polar covalent: unequal sharing due to electronegativity difference (e.g., H-Cl).

3. COORDINATE / DATIVE BOND

• Both electrons supplied by one atom. E.g., NH₃ → BF₃ adduct, NH₄⁺.

4. HYDROGEN BOND

• Weak (5-30 kJ/mol) attraction between H attached to F/O/N and a lone pair of another F/O/N.
• Explains: high b.p. of water, H₂O₂, HF; ice less dense than water; DNA base pairing.

5. METALLIC BOND

• "Sea of delocalized electrons" around positive metal cores. Explains: conductivity, malleability, ductility, luster.

LEWIS STRUCTURES — Rules

1. Count total valence electrons.
2. Draw skeleton (least EN atom usually central, except H).
3. Place bonds + lone pairs to satisfy octet.
4. If short, form double/triple bonds.
5. Check formal charges (minimize them).

VSEPR (Valence Shell Electron Pair Repulsion):

Predicts shape from number of e⁻ pairs (bonding + lone) around central atom.

Lone-pair distortion: lone pairs repel more than bond pairs.

• CH₄: 109.5° (perfect tetra). NH₃: 107° (1 LP). H₂O: 104.5° (2 LP).

HYBRIDIZATION:

Mixing atomic orbitals to form equivalent hybrid orbitals.

• sp (50% s, 50% p) — linear, 180°.
• sp² — trigonal planar, 120°.
• sp³ — tetrahedral, 109.5°.
• sp³d — TBP.
• sp³d² — octahedral.

Example — ethyne (C₂H₂): each C is sp; 2 sp orbitals form 1 σ to other C + 1 σ to H. Unhybridized 2p form two π bonds (C≡C triple).

MOLECULAR ORBITAL THEORY (MOT):

• Atomic orbitals combine to form bonding (σ, π) and antibonding (σ*, π*) MOs.
• Bond order = (Nb − Na) / 2.
• O₂: bond order 2, paramagnetic (2 unpaired e⁻ in π*).
• N₂: bond order 3, diamagnetic.
• He₂: bond order 0 → doesn't form.

FAJANS RULES — covalent character in ionic bond:

• Smaller cation → more polarizing → more covalent.
• Larger anion → more polarizable.
• Higher cation charge → more covalent.
• Cation with pseudo-noble-gas configuration (Cu⁺, Ag⁺, Zn²⁺) is more polarizing.

DIPOLE MOMENT (μ) = Q × d (Debye).

• BF₃ (μ=0, trigonal planar — vectors cancel).
• NH₃ (μ=1.46 D — pyramidal, doesn't cancel).
• H₂O (μ=1.85 D — bent).
• CO₂ (μ=0 — linear, cancels).

EXAM HOOKS:

• Strongest covalent: H-F. Strongest H-bond: F-H...F.
• Bent shape: 2 lone pairs (H₂O, OF₂); pyramidal: 1 LP (NH₃).
• sp³d hybridization → TBP (PCl₅, SF₄).
• Among NH₃, PH₃, AsH₃, SbH₃, BiH₃ → NH₃ has highest b.p. (H-bond).

Why bonds form? Atoms achieve stable octet (Lewis 1916). Inert gases have full outer shells → low reactivity.

TYPES OF BONDS:

1. IONIC BOND (electrostatic, complete e⁻ transfer)

• Between metal (loses e⁻) and non-metal (gains e⁻). E.g., NaCl, MgO.
• Properties: high m.p., conduct in molten/aqueous state, soluble in polar solvents.
• Lattice energy: energy to form 1 mol of ionic solid from gaseous ions. Higher = stronger bond.
• Born-Haber cycle to calculate lattice energy.

2. COVALENT BOND (e⁻ sharing)

• Between non-metals (e.g., H₂, Cl₂, H₂O).
• Single (1 pair), double (2 pairs), triple (3 pairs).
• Bond order = number of shared pairs.
• Bond length: C-C (154) > C=C (134) > C≡C (120) pm.
• Bond energy: C≡C > C=C > C-C.
• Polar covalent: unequal sharing due to electronegativity difference (e.g., H-Cl).

3. COORDINATE / DATIVE BOND

• Both electrons supplied by one atom. E.g., NH₃ → BF₃ adduct, NH₄⁺.

4. HYDROGEN BOND

• Weak (5-30 kJ/mol) attraction between H attached to F/O/N and a lone pair of another F/O/N.
• Explains: high b.p. of water, H₂O₂, HF; ice less dense than water; DNA base pairing.

5. METALLIC BOND

• "Sea of delocalized electrons" around positive metal cores. Explains: conductivity, malleability, ductility, luster.

LEWIS STRUCTURES — Rules

1. Count total valence electrons.
2. Draw skeleton (least EN atom usually central, except H).
3. Place bonds + lone pairs to satisfy octet.
4. If short, form double/triple bonds.
5. Check formal charges (minimize them).

VSEPR (Valence Shell Electron Pair Repulsion):

Predicts shape from number of e⁻ pairs (bonding + lone) around central atom.

Lone-pair distortion: lone pairs repel more than bond pairs.

• CH₄: 109.5° (perfect tetra). NH₃: 107° (1 LP). H₂O: 104.5° (2 LP).

HYBRIDIZATION:

Mixing atomic orbitals to form equivalent hybrid orbitals.

• sp (50% s, 50% p) — linear, 180°.
• sp² — trigonal planar, 120°.
• sp³ — tetrahedral, 109.5°.
• sp³d — TBP.
• sp³d² — octahedral.

Example — ethyne (C₂H₂): each C is sp; 2 sp orbitals form 1 σ to other C + 1 σ to H. Unhybridized 2p form two π bonds (C≡C triple).

MOLECULAR ORBITAL THEORY (MOT):

• Atomic orbitals combine to form bonding (σ, π) and antibonding (σ*, π*) MOs.
• Bond order = (Nb − Na) / 2.
• O₂: bond order 2, paramagnetic (2 unpaired e⁻ in π*).
• N₂: bond order 3, diamagnetic.
• He₂: bond order 0 → doesn't form.

FAJANS RULES — covalent character in ionic bond:

• Smaller cation → more polarizing → more covalent.
• Larger anion → more polarizable.
• Higher cation charge → more covalent.
• Cation with pseudo-noble-gas configuration (Cu⁺, Ag⁺, Zn²⁺) is more polarizing.

DIPOLE MOMENT (μ) = Q × d (Debye).

• BF₃ (μ=0, trigonal planar — vectors cancel).
• NH₃ (μ=1.46 D — pyramidal, doesn't cancel).
• H₂O (μ=1.85 D — bent).
• CO₂ (μ=0 — linear, cancels).

EXAM HOOKS:

• Strongest covalent: H-F. Strongest H-bond: F-H...F.
• Bent shape: 2 lone pairs (H₂O, OF₂); pyramidal: 1 LP (NH₃).
• sp³d hybridization → TBP (PCl₅, SF₄).
• Among NH₃, PH₃, AsH₃, SbH₃, BiH₃ → NH₃ has highest b.p. (H-bond).