Molecular Shapes, VSEPR Theory, and Molecular Polarity: Study Notes

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Chemical Bonding II: Molecular Shapes, VSEPR Theory, and Molecular Orbital Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory is a fundamental model used to predict the shapes of molecules based on the repulsions between electron groups around a central atom. The arrangement of these groups determines the geometry and bond angles of the molecule.

Electron groups include lone pairs, single, double, and triple bonds.
Electron groups repel each other due to coulombic forces, seeking maximum separation for stability.
The shape of the molecule is determined by the number and type of electron groups.

Electron groups and repulsions around a central atom

Basic Electron Group Geometries

There are five basic arrangements of electron groups around a central atom, each corresponding to a specific geometry and ideal bond angles.

Linear Geometry

Two electron groups around the central atom.
Groups are positioned opposite each other, resulting in a bond angle of 180°.
Example: BeCl2, CO2.

Lewis structure of BeCl2 Linear geometry of BeCl2 Linear geometry of CO2

Trigonal Planar Geometry

Three electron groups around the central atom.
Groups form a flat triangle with bond angles of 120°.
Example: BF3, formaldehyde (CH2O).

Lewis structure of BF3 Trigonal planar geometry of BF3 Lewis structure of formaldehyde Bond angles in formaldehyde

Comparison of Linear and Trigonal Planar Geometries

Comparison of linear and trigonal planar geometries

Tetrahedral Geometry

Four electron groups around the central atom.
Groups form a tetrahedron with bond angles of 109.5°.
Example: CH4.

Tetrahedral geometry Trigonal bipyramidal geometry Trigonal bipyramidal model Octahedral geometry Octahedral model

Effect of Lone Pairs on Molecular Geometry

Lone pairs exert greater repulsion than bonding pairs, causing bond angles to deviate from ideal values and altering molecular geometry.

Lone pairs are more spread out, as they are attracted to only one nucleus.
Bonding pairs are attracted to two nuclei, making them less repulsive.
Repulsion order: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

Bonding vs. lone pair electron distribution Effect of lone pairs on bond angles Bond angle distortion from lone pairs

Examples of Lone Pair Effects

Trigonal pyramidal: Four electron groups, one lone pair (e.g., NH3).
Bent: Four electron groups, two lone pairs (e.g., H2O).

Seesaw geometry Seesaw geometry model T-shaped geometry Linear geometry with lone pairs Square pyramidal geometry Square pyramidal model Square planar geometry

Summary Table: Electron and Molecular Geometries

The following table summarizes the relationship between electron groups, bonding groups, lone pairs, electron geometry, molecular geometry, and bond angles.

Electron Groups	Bonding Groups	Lone Pairs	Electron Geometry	Molecular Geometry	Approx. Bond Angles	Example
2	2	0	Linear	Linear	180°	CO2
3	3	0	Trigonal planar	Trigonal planar	120°	BF3
3	2	1	Trigonal planar	Bent	<120°	SO2
4	4	0	Tetrahedral	Tetrahedral	109.5°	CH4
4	3	1	Tetrahedral	Trigonal pyramidal	<109.5°	NH3
4	2	2	Tetrahedral	Bent	<109.5°	H2O
5	5	0	Trigonal bipyramidal	Trigonal bipyramidal	120°, 90°	PCl5
6	6	0	Octahedral	Octahedral	90°	SF6

Representing Three-Dimensional Shapes on Paper

To depict molecular geometry, chemists use conventions for bonds:

Straight line: bond in the plane of paper
Solid wedge: bond coming out of the page
Hashed wedge: bond going into the page

Bond representation conventions

Molecules with Multiple Central Atoms

Complex molecules may have several interior atoms, each with its own geometry. The overall shape is described by analyzing each central atom in sequence.

Multiple central atoms in a molecule Glycine structure with four interior atoms Glycine molecular geometry

Molecular Polarity

Molecular polarity arises from the presence of polar bonds and the overall shape of the molecule. Polarity affects properties such as solubility and intermolecular interactions.

A polar bond occurs when electrons are unequally shared between atoms of different electronegativities.
The net dipole moment depends on the vector sum of individual bond dipoles.
If dipoles cancel, the molecule is nonpolar; if they add, the molecule is polar.

Polar bond and net dipole moment in HCl No net dipole moment in CO2 Net dipole moment in H2O

Determining Molecular Polarity: Steps

Draw the Lewis structure and determine molecular geometry.
Identify polar bonds and draw vectors toward more electronegative atoms.
Sum the vectors to determine if there is a net dipole moment.

Vector Addition in Molecular Polarity

Vector addition is used to determine the net dipole moment in molecules.

Vectors pointing in opposite directions cancel (nonpolar).
Vectors pointing in the same direction add (polar).
Three identical polar bonds in trigonal planar geometry cancel (nonpolar).

Vector addition example 1 Vector addition example 2 Vector addition example 3 Vector addition example 4

Molecular Polarity and Solubility

Polar molecules dissolve well in polar solvents like water, while nonpolar molecules do not. Some molecules have both polar and nonpolar regions, affecting their solubility and biological function.

Polar molecules dissolving in water Polar and nonpolar regions affecting solubility

Valence Bond Theory and Hybridization

Valence bond theory explains chemical bonding as the overlap of atomic or hybrid orbitals, with spin-pairing of electrons. Hybridization adjusts the number and orientation of orbitals to match observed molecular shapes.

Hybrid orbitals are formed by mixing atomic orbitals (e.g., s and p).
The type of hybridization depends on the electron group geometry (see Table below).
Hybridization schemes: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), sp3d (trigonal bipyramidal), sp3d2 (octahedral).

Electron Groups	Electron Geometry	Hybridization Scheme
2	Linear	sp
3	Trigonal planar	sp2
4	Tetrahedral	sp3
5	Trigonal bipyramidal	sp3d
6	Octahedral	sp3d2

Types of Bonds: Sigma (σ) and Pi (π)

σ bond: Formed by head-on overlap of orbitals along the internuclear axis; stronger and allows free rotation.
π bond: Formed by side-by-side overlap of unhybridized p orbitals; weaker and restricts rotation.

Molecular Orbital (MO) Theory

MO theory applies quantum mechanics to molecules, combining atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.

Bonding molecular orbitals (constructive combination) are lower in energy.
Antibonding molecular orbitals (destructive combination) are higher in energy.
Bond order =
Paramagnetic molecules have unpaired electrons in MO diagrams; diamagnetic molecules have all electrons paired.

Summary and Applications

VSEPR theory predicts molecular shapes based on electron group repulsions.
Molecular geometry and polarity influence physical and chemical properties.
Valence bond and MO theories provide deeper understanding of bonding and molecular behavior.