chapter 7 part 2

Chemical Bonding and Molecular Geometry

Resonance

Resonance is a concept used to describe molecules or ions whose structures cannot be represented by a single Lewis structure. Instead, two or more valid Lewis structures (called resonance structures) are drawn, and the actual structure is considered a hybrid of these.

• Definition: Resonance occurs when more than one valid Lewis structure can be drawn for a molecule or ion, differing only in the placement of electrons.

• Example: Ozone (O3) – The molecule can be represented as O = O – O or O – O = O, but experimentally both bonds are equal (128 pm). The true structure is a resonance hybrid.

• Notation: Resonance structures are connected by a double-headed arrow (↔).

• Example: Nitrate ion (NO3-) – Three resonance structures exist, each with different arrangements of double and single bonds.

• Other Examples: Carbonate ion (CO32-), benzene (C6H6).

Additional info: Resonance stabilizes molecules by delocalizing electrons, lowering their energy.

Molecular Geometry & Polarity

Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule. The geometry affects physical and chemical properties, including polarity.

• Bond Lengths & Angles: Determined experimentally, but can be predicted using electron pair repulsion models.

• Electron Pair Repulsion: Electron pairs in the valence shell repel each other, arranging themselves to minimize repulsion.

The VSEPR Model

The Valence-Shell Electron-Pair Repulsion (VSEPR) model predicts molecular geometry based on the repulsion between electron pairs.

• Rule 1: Double and triple bonds are treated as single bonds, but occupy more space due to higher electron density.

• Rule 2: VSEPR can be applied to any resonance structure; formal charges are not shown.

• Repulsion Order: Lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

Electron-Domain Geometry and Molecular Geometry

Electron domains include both bonding pairs and lone pairs. The arrangement of these domains determines the electron-domain geometry, while the arrangement of atoms determines the molecular geometry.

• Notation: A = central atom, B = bonded atoms, E = lone pairs.

• Common combinations: AB2, AB3, AB2E, AB4, AB3E, AB2E2, etc.

Geometries Based on Electron Groups

Electron Groups	Type	Example	Electron Geometry	Molecular Geometry	Bond Angles
2	AB2	BeCl2, CO2	Linear	Linear	180°
3	AB3	BF3, H2CO	Trigonal planar	Trigonal planar	120°
3	AB2E	SO2, NO2-	Trigonal planar	Bent	<120°
4	AB4	CH4, SiCl4	Tetrahedral	Tetrahedral	109.5°
4	AB3E	NH3	Tetrahedral	Trigonal pyramidal	107.3°
4	AB2E2	H2O	Tetrahedral	Bent	104.5°
5	AB5	PCl5	Trigonal bipyramidal	Trigonal bipyramidal	90°, 120°, 180°
5	AB4E	SF4	Trigonal bipyramidal	Seesaw	<90°, <120°, <180°
5	AB3E2	ClF3	Trigonal bipyramidal	T-shaped	<90°, <180°
5	AB2E3	XeF2, I3-	Trigonal bipyramidal	Linear	180°
6	AB6	SF6	Octahedral	Octahedral	90°, 180°
6	AB5E	BrF5	Octahedral	Square pyramidal	<90°, <180°
6	AB4E2	XeF4	Octahedral	Square planar	90°, 180°

Additional info: Lone pairs reduce bond angles compared to ideal geometries due to increased repulsion.

Geometry of Molecules with Multiple Central Atoms

For molecules with more than one central atom, analyze the geometry around each atom separately.

• Example: Ethanol (CH3CH2OH) – C1 and C2 are both tetrahedral (bond angles ≈ 109°), O is bent (bond angle ≈ 105°).

Guidelines for Applying the VSEPR Model

1. Write the Lewis structure, focusing on electron pairs around the central atom.

2. Treat double and triple bonds as single bonds for geometry prediction.

3. Predict the overall arrangement of electron pairs.

4. Predict the molecular geometry and bond angles, considering repulsion order.

Additional info: Examples include PCl3, ICl4-, CH3OH, H2NCH2COOH, C2H4, C2H2, IF4+, IF2-, PF5.

Dipole Moments

Dipole moment is a quantitative measure of the polarity of a molecule, resulting from the separation of positive and negative charges.

• Definition: The dipole moment (μ) is the product of the magnitude of the charge (Q) and the distance (r) between the charges.

• Formula: $\mu = Q \times r$

• Units: Debye (D), where 1 D = 3.336 × 10-30 C·m.

• Notation: A crossed arrow (+→) or δ+ and δ- indicate bond polarity.

• Example: HF molecule – H is δ+, F is δ-, and the molecule aligns in an electric field.

• Homodiatomic molecules: (e.g., O2, N2) have μ = 0; heteronuclear diatomics (e.g., CO, HCl) are polar.

Determining Molecular Polarity

1. Check if the molecule contains polar bonds (use vector notation towards more electronegative atom).

2. Determine if bond dipoles add to a net dipole moment (sum vectors).

3. If vectors sum to zero, molecule is nonpolar (e.g., CO2).

• Examples: H2O (polar), CCl4 (nonpolar), CHCl3 (polar), NH3 (polar), BF3 (nonpolar).

Dipole Moment and Partial Charges

Partial charges (δ+ and δ-) quantify the charge separation in a polar bond. The dipole moment can be used to estimate the magnitude of these charges.

• Calculation Example (HI molecule):

• Bond length: 1.61 Å = $1.61 \times 10^{-10}$ m

• Dipole moment: 0.44 D = $0.44 \times 3.336 \times 10^{-30}$ C·m = $1.468 \times 10^{-30}$ C·m

• Partial charge: $Q = \frac{\mu}{r} = \frac{1.468 \times 10^{-30}}{1.61 \times 10^{-10}} = 9.118 \times 10^{-21}$ C

• In units of electron charge: $\frac{9.118 \times 10^{-21}}{1.6022 \times 10^{-19}} = 0.057$ electrons

• Therefore, H = +0.057, I = -0.057 (partial charges).

Dipole Moments of Selected Molecules

Molecule	Bond Length (Å)	Dipole Moment (D)
HF	0.92	1.82
HCl	1.27	1.08
HBr	1.41	0.82
HI	1.61	0.44

Additional info: The partial charge is less than a full electron charge, indicating partial, not complete, electron transfer.