Chapter 10: Chemical Bonding II – Molecular Shapes and Valence Bond Theory (CHEM 1000 Study Notes)

Chapter 10: Chemical Bonding II – Molecular Shapes and Valence Bond Theory

Learning Objectives

• Use VSEPR theory to predict and explain molecular geometries.

• Explain the fundamental basis of valence bond theory.

• Write hybridization schemes for the formation of sp, sp2, sp3, sp3d, and sp3d2 hybrid orbitals; predict geometric shapes of molecules in terms of the pure and hybrid orbitals used in bonding.

• Use valence bond theory to predict molecular geometries and bonding schemes.

• Describe multiple bonds between second period elements in terms of the overlap of sp, sp2, sp3 and pure 2p orbitals to form σ bonds, and the sidewise overlap of p orbitals to form π bonds.

• Propose plausible bonding schemes from Lewis structures or from experimental information about molecules (bond lengths, bond angles, etc.).

Shapes of Molecules

VSEPR Theory and Electron Groups

The Valence Shell Electron Pair Repulsion (VSEPR) Theory is used to predict the three-dimensional structure of molecules based on the repulsion between electron groups around a central atom.

• Electron groups include lone pairs, single bonds, multiple bonds, and single electrons.

• Electron groups repel each other and assume orientations that minimize repulsion.

• Geometry is usually described around the central atom.

Electron Group Geometries

Two Electron Groups: Linear Geometry

When two electron groups are present, they arrange themselves as far apart as possible, resulting in a linear geometry with a bond angle of 180°.

• Example: BeCl2, CO2, acetylene (C2H2)

$\text{Bond angle} = 180^\circ$

Three Electron Groups: Trigonal Planar Geometry

Three electron groups arrange in a plane with 120° bond angles, forming a trigonal planar geometry.

• Example: BF3, formaldehyde (CH2O)

$\text{Bond angle} = 120^\circ$

Four Electron Groups: Tetrahedral Geometry

Four electron groups arrange in three dimensions with bond angles of approximately 109.5°, forming a tetrahedral geometry.

• Example: CH4 (methane)

$\text{Bond angle} = 109.5^\circ$

Five Electron Groups: Trigonal Bipyramidal Geometry

Five electron groups arrange in a trigonal bipyramidal geometry with two axial positions (180° apart) and three equatorial positions (120° apart).

• Example: PCl5

$\text{Bond angles} = 90^\circ, 120^\circ, 180^\circ$

Six Electron Groups: Octahedral Geometry

Six electron groups arrange symmetrically in an octahedral geometry with 90° bond angles.

• Example: SF6

$\text{Bond angle} = 90^\circ$

Worked Example: Geometry of NO3- Anion

• Draw Lewis structure: Count valence electrons (N: 5, O: 6 × 3, charge: +1) = 24 electrons.

• Count electron groups around central atom (N): 3

• Identify geometry: Trigonal planar

Effect of Lone Pairs on Geometry

Electron Geometry vs. Molecular Geometry

Electron geometry refers to the arrangement of all electron groups (including lone pairs) around the central atom. Molecular geometry refers to the arrangement of only the atoms (excluding lone pairs).

• Lone pairs count as electron groups and can alter the molecular geometry.

Three Electron Groups, One Lone Pair: Bent Geometry

• Example: SO2

• Electron geometry: Trigonal planar

• Molecular geometry: Bent

• Approximate bond angle: <120°

Four Electron Groups, One Lone Pair: Trigonal Pyramidal Geometry

• Example: NH3 (ammonia)

• Electron geometry: Tetrahedral

• Molecular geometry: Trigonal pyramidal

• Approximate bond angle: <109.5°

Four Electron Groups, Two Lone Pairs: Bent Geometry

• Example: H2O (water)

• Electron geometry: Tetrahedral

• Molecular geometry: Bent

• Approximate bond angle: <109.5°

Effects of Lone Pairs on Bond Angles

Lone pairs are more spread out in space than bonding pairs, causing bond angles to be compressed from their ideal values.

• Order of repulsion: lone pair/lone pair > lone pair/bonded pair > bonded pair/bonded pair

• Bond angles in NH3 and H2O are less than 109.5° due to lone pair compression.

Five Electron Groups: Lone Pair Effects

One Lone Pair: Seesaw Geometry

• Example: SF4

• Electron geometry: Trigonal bipyramidal

• Molecular geometry: Seesaw

• Lone pair occupies equatorial position to minimize repulsion.

• Approximate bond angles: <120° and/or <90°

Two Lone Pairs: T-Shaped Geometry

• Example: BrF3

• Electron geometry: Trigonal bipyramidal

• Molecular geometry: T-shaped

• Lone pairs occupy equatorial positions.

• Approximate bond angles: <90°

Three Lone Pairs: Linear Geometry

• Example: XeF2

• Electron geometry: Trigonal bipyramidal

• Molecular geometry: Linear

• Lone pairs occupy all equatorial positions.

• Bond angle: 180°

Six Electron Groups: Lone Pair Effects

One Lone Pair: Square Pyramidal Geometry

• Example: BrF5

• Electron geometry: Octahedral

• Molecular geometry: Square pyramidal

• Lone pair can occupy any position.

• Approximate bond angles: <90°

Two Lone Pairs: Square Planar Geometry

• Example: XeF4

• Electron geometry: Octahedral

• Molecular geometry: Square planar

• Lone pairs occupy positions 180° apart.

• Bond angle: 90°

Summary Table: VSEPR Geometries and Bond Angles

Electron Groups	Lone Pairs	Electron Geometry	Molecular Geometry	Approx. Bond Angles	Example
2	0	Linear	Linear	180°	CO2
3	0	Trigonal planar	Trigonal planar	120°	BF3
3	1	Trigonal planar	Bent	<120°	SO2
4	0	Tetrahedral	Tetrahedral	109.5°	CH4
4	1	Tetrahedral	Trigonal pyramidal	<109.5°	NH3
4	2	Tetrahedral	Bent	<109.5°	H2O
5	0	Trigonal bipyramidal	Trigonal bipyramidal	90°, 120°, 180°	PCl5
5	1	Trigonal bipyramidal	Seesaw	<120°, <90°	SF4
5	2	Trigonal bipyramidal	T-shaped	<90°	BrF3
5	3	Trigonal bipyramidal	Linear	180°	XeF2
6	0	Octahedral	Octahedral	90°	SF6
6	1	Octahedral	Square pyramidal	<90°	BrF5
6	2	Octahedral	Square planar	90°	XeF4

Predicting Shapes of Larger Molecules

For larger molecules (especially organic compounds), multiple interior (central) atoms may be present. VSEPR theory can be applied to each interior atom to determine its local geometry.

• Example: Glycine (NH2CH2COOH) – VSEPR is used to determine the geometry around each nitrogen and carbon atom.

Drawing Tetrahedral Structures

Tetrahedral geometries are common in organic molecules, especially for carbon atoms with four single bonds. Three-dimensional representations use:

• Straight line: Bond in plane of paper

• Hashed wedge: Bond going into the page

• Solid wedge: Bond coming out of the page

Worked Example: Methanol (CH3OH)

• Draw Lewis structure: C (4), O (6), H (1 × 4) = 14 valence electrons

• Predict geometry:

  • C: 4 electron groups, 0 lone pairs → tetrahedral (109.5°)

  • O: 4 electron groups, 2 lone pairs → electron geometry: tetrahedral; molecular geometry: bent (<109.5°)

Key Takeaways

• VSEPR theory is a powerful tool for predicting molecular shapes based on electron group repulsions.

• Lone pairs affect both geometry and bond angles, often compressing angles below ideal values.

• Practice is essential for mastering molecular geometry predictions.

Additional info: Hybridization schemes and valence bond theory are mentioned in objectives but not detailed in these slides. For full context, students should review hybrid orbital formation and sigma/pi bonding in the textbook.