Chapter 9: Molecular Geometry and Bonding Theories – Study Notes

Chapter 9: Molecular Geometry and Bonding Theories

9.1 Molecular Shapes

The three-dimensional shape of a molecule is crucial in determining its physical and chemical properties. While Lewis structures show the connectivity and types of bonds, they do not represent the actual spatial arrangement of atoms.

• Molecular shape refers to the arrangement of atoms in space, not just their connectivity.

• The size and shape of a molecule, along with bond strength and polarity, influence properties such as reactivity, polarity, phase, color, magnetism, and biological activity.

• Example: In CCl4, Lewis structures show all atoms in a plane, but the actual geometry is tetrahedral, with Cl atoms at the corners of a tetrahedron.

9.2 Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

The VSEPR model is used to predict the shapes of molecules based on the repulsion between electron domains (regions where electrons are likely to be found) around a central atom.

• Electron domain: A region where electrons are most likely found (includes bonding pairs and lone pairs).

• Bonding pair: Electrons shared between two atoms (single, double, or triple bonds count as one domain each).

• Non-bonding pair (lone pair): Electrons localized on one atom.

• Only electron domains around the central atom are counted for geometry prediction.

Electron Domain Geometries

Number of Electron Domains	Electron Domain Geometry	Predicted Bond Angles	Example
2	Linear	180°	CO2
3	Trigonal planar	120°	BF3
4	Tetrahedral	109.5°	CH4
5	Trigonal bipyramidal	120°, 90°, 180°	PCl5
6	Octahedral	90°, 180°	SF6

Molecular Geometry

The molecular geometry describes the arrangement of only the atoms (not lone pairs) in a molecule. Lone pairs affect the geometry by repelling bonding pairs, altering bond angles.

Electron Domains	Bonding Domains	Nonbonding Domains	Molecular Geometry	Example
3	3	0	Trigonal planar	BF3
3	2	1	Bent	SO2
4	4	0	Tetrahedral	CH4
4	3	1	Trigonal pyramidal	NH3
4	2	2	Bent	H2O
5	5	0	Trigonal bipyramidal	PCl5
5	4	1	Seesaw	SF4
5	3	2	T-shaped	ClF3
5	2	3	Linear	XeF2
6	6	0	Octahedral	SF6
6	5	1	Square pyramidal	IF5
6	4	2	Square planar	XeF4

Steps for Predicting Molecular Geometry (VSEPR)

1. Draw the Lewis structure of the molecule or ion.

2. Count the number of electron domains around the central atom.

3. Determine the electron domain geometry.

4. Determine the molecular geometry by considering only the arrangement of bonded atoms.

Examples

• BF3: 3 electron domains, trigonal planar geometry, bond angles 120°.

• CCl4: 4 electron domains, tetrahedral geometry, bond angles 109.5°.

• PCl5: 5 electron domains, trigonal bipyramidal geometry, bond angles 120°, 90°, 180°.

• XeF2: 5 electron domains (3 lone pairs), linear molecular geometry, bond angle 180°.

Predicting Bond Angles

Bond angles are determined by the electron domain geometry. Lone pairs compress bond angles relative to the ideal geometry.

• Tetrahedral: 109.5°

• Trigonal planar: 120°

• Linear: 180°

• Trigonal bipyramidal: 120° (equatorial), 90° (axial), 180°

• Octahedral: 90°, 180°

9.3 Molecular Shape and Molecular Polarity

The polarity of a molecule depends on both the polarity of its bonds and its molecular geometry. Bond dipoles are vector quantities; their sum determines the overall molecular dipole moment.

• If bond dipoles are equal in magnitude and opposite in direction, they cancel, resulting in a non-polar molecule (e.g., CO2).

• If bond dipoles do not cancel, the molecule is polar (e.g., H2O, NH3).

• Symmetry plays a key role: molecules with symmetrical shapes (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral) are non-polar if all outer atoms are the same.

• Asymmetrical shapes (bent, trigonal pyramidal, seesaw, T-shaped, square pyramidal) are always polar.

Lewis Structure	Electron Domain Geometry	Molecular Geometry	Polarity, Bond Angles
C2H2	Linear	Linear	Non-polar, 180°
XeF4	Octahedral	Square planar	Non-polar, 180°
SO3	Trigonal planar	Trigonal planar	Non-polar, 120°
IF5	Octahedral	Square pyramidal	Polar, <90°

Summary Table: Polarity and Geometry

Always Non-polar (if all outer atoms are the same)	Always Polar
Linear Trigonal bipyramidal Tetrahedral Trigonal planar Square planar	Square pyramidal Seesaw T-shaped Trigonal pyramidal Bent

9.4 Covalent Bonding and Orbital Overlap

Valence-bond theory explains covalent bonding as the overlap of atomic orbitals from two atoms, resulting in a shared electron pair.

• Bonding occurs when atomic orbitals overlap, allowing electrons to be shared.

• There is an optimum bond length where the potential energy is minimized (balance between attractive and repulsive forces).

• Example: In H2, two 1s orbitals overlap to form a sigma bond.

9.5 Hybrid Orbitals

To explain observed molecular geometries, atomic orbitals on a central atom can mix to form hybrid orbitals with new shapes and orientations.

• sp hybridization: Linear geometry (e.g., BeCl2).

• sp2 hybridization: Trigonal planar geometry (e.g., BF3).

• sp3 hybridization: Tetrahedral geometry (e.g., CH4).

• sp3d hybridization: Trigonal bipyramidal geometry (e.g., PCl5).

• sp3d2 hybridization: Octahedral geometry (e.g., SF6).

Hybridization can be determined by the number of electron domains around the central atom (number of hybrid orbitals = number of electron domains).

Examples

• NH3: 4 electron domains, sp3 hybridization (trigonal pyramidal geometry).

• SO32-: 4 electron domains, sp3 hybridization (tetrahedral electron domain geometry).

• XeF2: 5 electron domains, sp3d hybridization (linear molecular geometry).

• SF6: 6 electron domains, sp3d2 hybridization (octahedral geometry).

9.6 Multiple Bonds

Multiple bonds involve more than one shared pair of electrons between two atoms. There are two types of covalent bonds:

• Sigma (σ) bond: Electron density is concentrated along the internuclear axis (directly between the nuclei). All single bonds are sigma bonds.

• Pi (π) bond: Electron density is above and below the internuclear axis, formed by the sideways overlap of unhybridized p orbitals. Present in double and triple bonds.

Double bonds consist of one sigma and one pi bond; triple bonds consist of one sigma and two pi bonds.

Example: Formaldehyde (CH2O)

• Electron domain geometry around C: Trigonal planar (120° bond angles).

• Hybridization around C: sp2.

• One unhybridized p orbital on C forms the pi bond with O.

• C-H bonds: sigma bonds (sp2 on C overlaps with 1s on H).

• C=O bond: one sigma (sp2 on C with sp2 on O) and one pi (unhybridized p orbitals on C and O).

Example: Acetonitrile (CH3CN)

• Hybridization at central C: sp (linear geometry).

• Number of sigma bonds: 5; number of pi bonds: 2.

• C-H and C-C bonds: sigma bonds; C≡N bond: one sigma and two pi bonds.

Summary Table: Sigma and Pi Bonds

Bond Type	Number of Sigma Bonds	Number of Pi Bonds
Single	1	0
Double	1	1
Triple	1	2

Key Equations:

• Bond angles for ideal geometries:

  • Linear: $180^\circ$

  • Trigonal planar: $120^\circ$

  • Tetrahedral: $109.5^\circ$

  • Trigonal bipyramidal: $120^\circ$, $90^\circ$, $180^\circ$

  • Octahedral: $90^\circ$, $180^\circ$

Additional info: These notes are based on lecture slides and textbook images, with tables and examples reconstructed for clarity and completeness. All molecular geometries and hybridizations are determined using VSEPR and valence bond theory principles.