Molecular Geometry, Polarity, and Intermolecular Forces of Attraction

Chapter 4: Molecular Geometry, Polarity, and Intermolecular Forces of Attraction

Three-Dimensional Shapes of Molecules

Molecules are not flat; they possess complex three-dimensional structures that influence their physical and chemical properties. The spatial arrangement of atoms in a molecule determines its reactivity, polarity, and interactions with other molecules.

• Chirality: Some molecules, such as those found in spearmint and caraway, have the same molecular formula but different three-dimensional arrangements, leading to different smells and biological activities.

• Structural Representation: Molecular structures can be represented in various ways, including Lewis dot structures, ball-and-stick models, and space-filling models.

• Example: Ethanol (C2H5OH) can be depicted using all three models to illustrate bonding and spatial arrangement.

Electron Geometry and Molecular Geometry

The shape of a molecule is predicted using the Valence-Shell Electron-Pair Repulsion (VSEPR) Theory. This theory states that electron groups around a central atom arrange themselves to minimize repulsion, determining the molecule's geometry.

• Stepwise Process:

  1. Draw the Lewis structure.

  2. Determine the electron geometry (arrangement of electron groups).

  3. Determine the molecular geometry (arrangement of atoms).

• Electron Geometry: Considers all electron groups (bonding and lone pairs) around the central atom.

• Molecular Geometry: Considers only the positions of atoms (ignoring lone pairs).

Common Electron and Molecular Geometries

Electron geometry is based on the number of electron groups (bonding and lone pairs) around the central atom. Molecular geometry is derived by considering only the positions of atoms.

Total Electron Groups	Bonding Groups	Lone Pairs	Electron Geometry	Molecular Shape	Approximate Bond Angle	Example
2	2	0	Linear	Linear	180°	CO2, HCN
3	3	0	Trigonal planar	Trigonal planar	120°	BF3, H2CO
3	2	1	Trigonal planar	Bent	~120°	SO2
4	4	0	Tetrahedral	Tetrahedral	109.5°	CH4
4	3	1	Tetrahedral	Trigonal pyramidal	~107°	NH3
4	2	2	Tetrahedral	Bent	~104.5°	H2O

Additional info: Bond angles decrease as the number of lone pairs increases due to greater repulsion by lone pairs compared to bonding pairs.

Representing Molecular Geometries on Paper

To depict three-dimensional molecular shapes on two-dimensional paper, chemists use specific notations:

• Straight line (–): Bond in the plane of the paper.

• Hashed line (---): Bond projecting into the paper (away from the viewer).

• Wedge (▲): Bond projecting out of the paper (toward the viewer).

This notation helps visualize the spatial arrangement of atoms, such as in methane (CH4).

Molecular Polarity

Polarity determines how molecules interact with each other and with other substances. It is influenced by both the polarity of individual bonds and the overall molecular geometry.

• Electronegativity: The ability of an atom in a covalent bond to attract electrons toward itself. The greater the difference in electronegativity between two atoms, the more polar the bond.

• Bond Polarity:

  • Nonpolar covalent: Electrons shared equally (e.g., H2, Cl2).

  • Polar covalent: Electrons shared unequally (e.g., H–F).

  • Ionic: Electrons transferred (e.g., NaCl).

• Determining Bond Type by Electronegativity Difference:

  • Nonpolar covalent: Difference < 0.5

  • Polar covalent: 0.5 ≤ Difference < 2.0

  • Ionic: Difference ≥ 2.0

• Molecular Polarity: A molecule is polar if it contains polar bonds arranged asymmetrically, resulting in a net dipole moment. If the geometry is symmetrical and all bonds are identical, the molecule is nonpolar.

• Examples:

  • CO2: Nonpolar (linear, dipoles cancel)

  • H2O: Polar (bent, dipoles do not cancel)

  • BF3: Nonpolar (trigonal planar, dipoles cancel)

  • NH3: Polar (trigonal pyramidal, net dipole)

Intermolecular Forces of Attraction

Intermolecular forces are the forces of attraction or repulsion between neighboring particles (atoms, molecules, or ions). They are generally much weaker than covalent bonds but play a crucial role in determining the physical properties of substances, such as melting and boiling points.

• Types of Intermolecular Forces:

  1. London Dispersion Forces: Present in all molecules, especially significant in nonpolar molecules. Caused by temporary dipoles due to electron movement. Strength increases with molecular size and shape.

  2. Dipole-Dipole Forces: Occur between polar molecules with permanent dipoles. Stronger than dispersion forces.

  3. Hydrogen Bonding: A special, strong type of dipole-dipole interaction occurring when hydrogen is bonded to highly electronegative atoms (F, O, N). Responsible for unique properties of water and the structure of biological molecules like DNA.

• Effect on Physical Properties: Stronger intermolecular forces lead to higher melting and boiling points.

• Example: H2O (hydrogen bonding) is a liquid at room temperature, while CO2 (only dispersion forces) is a gas.

Summary Table: Types of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Example
London Dispersion	All molecules (especially nonpolar)	Weakest	F2, CH4
Dipole-Dipole	Polar molecules	Intermediate	HCl, CH2O
Hydrogen Bonding	H bonded to F, O, or N	Strongest (of intermolecular forces)	H2O, NH3

Key Learning Objectives

• Predict the shapes of molecules using VSEPR theory.

• Determine whether a molecule is polar or nonpolar based on its geometry and bond polarity.

• Identify the types of intermolecular forces present in a compound.

• Relate the strength of intermolecular forces to melting and boiling points.