Special Intermolecular Interactions

Overview of Special Intermolecular Interactions

In the context of chemical bonding, special intermolecular interactions are comparatively weak forces between molecules or between different parts of large molecules. They are distinct from the stronger intramolecular bonds (covalent, ionic, metallic) that hold atoms together within a molecule or solid.

These interactions are crucial for:

• Condensation of gases to liquids and solids
• Physical properties such as boiling and melting points, viscosity, and solubility
• Structures of biological macromolecules (proteins, DNA, membranes)
• Recognition processes (enzyme–substrate, receptor–ligand, antigen–antibody)

In this chapter, the focus is on two particularly important types of special intermolecular interactions:

• van der Waals forces
• Hydrogen bonding

General features common to both:

• They are electrostatic in nature: they arise due to attractions between charges, partial charges, or dipoles.
• Their strengths are significantly weaker than typical covalent or ionic bonds, but collective effects (many interactions at once) can be substantial.
• They act over relatively short distances, usually comparable to or slightly larger than molecular dimensions.
• They are highly directional or geometry-dependent to varying degrees (strongly so for hydrogen bonds, less for van der Waals interactions).

The following subsections treat these interactions in detail.

Types of Special Intermolecular Interactions

Even though different names are used, the interactions can be broadly understood as arising from charge distributions within and between molecules. Important categories include:

• Dipole–dipole interactions (between permanent dipoles)
• Ion–dipole and ion–induced dipole interactions
• Dipole–induced dipole interactions
• London dispersion forces (between instantaneous dipoles in all molecules)
• Hydrogen bonds (a special case with a hydrogen atom bridging two electronegative atoms)

In introductory chemistry, the term “van der Waals forces” often serves as a collective name for all weak intermolecular forces except hydrogen bonds. In more detailed treatments, London dispersion, dipole–dipole, and dipole–induced dipole interactions are distinguished and analyzed separately.

Energetics and Relative Strength

The typical energy ranges highlight how these interactions compare:

• Covalent bond: roughly $100$–$1000\ \text{kJ mol}^{-1}$
• Ionic lattice energy: typically several $100\ \text{kJ mol}^{-1}$ per ion pair
• Hydrogen bond: roughly $10$–$40\ \text{kJ mol}^{-1}$ (varies with system)
• van der Waals (individual contribution): roughly $0.1$–$10\ \text{kJ mol}^{-1}$

Even though an individual van der Waals interaction is weak, the cumulative effect in condensed phases or in large molecules (e.g. in polymers, proteins, DNA base stacking) becomes very important.

The energy of these interactions typically decreases rapidly with distance $r$ between interacting entities. For example, for London dispersion forces between two spherical particles:
$$
E_\text{dispersion} \propto -\frac{1}{r^6}
$$
This strong distance dependence means that small changes in molecular proximity can significantly alter interaction energies.

Directionality and Geometry

Intermolecular interactions are not only about “how strong” but also about “in which directions” and “at what distances” they act.

• van der Waals forces:
• Often approximately isotropic for nonpolar, spherical molecules (e.g. noble gases).
• Still affected by molecular shape and surface area; elongated or flat molecules can have stronger cumulative interactions due to larger contact surfaces.
• Hydrogen bonds:
• Strongly directional; optimal when the atoms involved (donor atom, hydrogen, and acceptor atom) are nearly collinear.
• Geometry leads to specific structures in solids and liquids (e.g. ice crystals, secondary structures in proteins, base pairing in DNA).

Because of directionality, intermolecular interactions can stabilize particular arrangements of molecules and, in the case of macromolecules, define their three-dimensional structure.

Role in Phase Transitions and Physical Properties

Special intermolecular interactions are central to understanding why substances exist as gases, liquids, or solids under given conditions, and why different substances exhibit very different physical properties.

Boiling and Melting Points

The strength and number of intermolecular interactions largely determine:

• Boiling point: temperature at which vapor pressure equals external pressure
• Melting point: temperature at which solid and liquid are in equilibrium

In general:

• Stronger intermolecular interactions → higher boiling and melting points
• Molecules capable of hydrogen bonding typically have higher boiling points than similar-sized molecules unable to hydrogen bond.
• For nonpolar molecules of similar shape, larger molar mass (and thus more electrons) enhances London dispersion forces, also raising boiling points.

Trends often observed:

• Among alkanes, boiling point increases with chain length due to increasing surface area and dispersion forces.
• Alcohols and amines, which can form hydrogen bonds, have higher boiling points than hydrocarbons of similar molar mass.

Solubility and Miscibility

Intermolecular interactions help rationalize the principle “like dissolves like”:

• Polar solvents (e.g. water) interact strongly via dipole–dipole forces and hydrogen bonds with polar or hydrogen-bonding solutes.
• Nonpolar solvents (e.g. hexane) interact mainly through London dispersion forces and readily dissolve nonpolar solutes.

The balance between solute–solute, solvent–solvent, and solute–solvent intermolecular forces determines whether dissolution is favorable.

Viscosity and Surface Tension

• Viscosity (resistance to flow) increases when molecules interact more strongly and/or become more entangled. Hydrogen bonding and strong dipolar interactions tend to increase viscosity.
• Surface tension arises because molecules at the surface of a liquid experience unbalanced intermolecular attractions, pulling them inward. Liquids with stronger intermolecular forces often exhibit higher surface tension (e.g. water).

These macroscopic properties offer experimental evidence for the presence and relative strength of intermolecular interactions.

Intermolecular Interactions in Molecular Recognition and Self-Assembly

In chemical and biological systems, molecules often recognize and bind to each other in specific ways without forming permanent covalent bonds. This specificity is largely governed by special intermolecular interactions.

Molecular Recognition

Typical examples include:

• Enzyme–substrate binding
• Antigen–antibody recognition
• Receptor–ligand interactions
• Host–guest complexes in supramolecular chemistry

Key principles:

• Complementarity of shape (lock-and-key or induced-fit concepts)
• Complementarity of charge and polarity (arrangements of hydrogen bond donors/acceptors and dipoles)
• Multiple weak interactions working cooperatively (hydrogen bonds, dispersion forces, ion–dipole interactions)

A single weak interaction is usually too small to ensure specificity. However, combinations of many such interactions produce sufficiently strong and selective binding.

Self-Assembly and Supramolecular Structures

Intermolecular forces guide spontaneous organization of molecules into:

• Molecular crystals
• Liquid crystals
• Micelles and bilayers (e.g. from surfactants or lipids)
• Supramolecular polymers and cages

The driving forces include:

• Minimization of free energy through optimal packing and interaction patterns
• Maximization of favorable interactions (e.g. hydrogen bonds, dispersion stabilization in stacked aromatic rings)
• Exclusion of unfavorable contacts (e.g. hydrophobic effect in water)

Here, the reversibility of weak interactions allows dynamic structures that can respond to external stimuli such as concentration, temperature, or solvent changes.

Competition and Cooperation Between Different Intermolecular Interactions

In real systems, several types of interactions often act simultaneously and can either reinforce or counteract each other.

Examples:

• In aqueous solutions of ionic compounds, ion–dipole interactions between ions and water compete with ion–ion attractions in the crystal lattice.
• In proteins, hydrogen bonding patterns compete with interactions between polar side chains and water, and cooperate with hydrophobic interactions that drive nonpolar residues into the interior.
• In molecular crystals, hydrogen bonding may define a particular directional framework, while dispersion forces help pack the remaining parts of the molecules.

The net behavior (solubility, structure, stability) results from the balance of all these contributions.

Experimental and Theoretical Description

Intermolecular interactions are characterized and quantified using both experimental methods and theoretical models.

Experimental Approaches

Some observables that are sensitive to intermolecular interactions include:

• Boiling and melting points, heats of vaporization and fusion
• Viscosity, surface tension, and density of liquids
• Crystal structures from X-ray diffraction
• Spectroscopic changes upon association (e.g. shifts in IR, NMR, or UV/Vis spectra)
• Solubility and partition coefficients between phases

Measurements of these properties allow indirect inference of interaction strengths, distances, and geometries.

Theoretical and Computational Models

To describe and predict intermolecular interactions, chemists use:

• Classical electrostatic models for permanent charges and dipoles
• Induction and dispersion terms often modeled with empirical potentials (e.g. Lennard-Jones type)
• Quantum chemical calculations to obtain more accurate interaction energies and potential energy surfaces

A typical empirical potential for nonbonded interactions between atoms or molecules is the Lennard-Jones potential:
$$
E(r) = 4\varepsilon \left[\left(\frac{\sigma}{r}\right)^{12} - \left(\frac{\sigma}{r}\right)^6\right]
$$
where:

• $r$ is the distance between particles
• $\varepsilon$ represents the depth of the potential well (interaction strength)
• $\sigma$ characterizes the effective collision diameter (distance at zero interaction energy)

The $r^{-12}$ term models strong short-range repulsion, while the $r^{-6}$ term represents attractive dispersion forces.

These models are foundational in molecular simulations of liquids, solutions, and biomolecules.

Relevance Across Chemistry and Related Fields

Special intermolecular interactions intersect with many other topics in the course:

• Inorganic and coordination chemistry: packing of ions and complexes in crystals, solvation of ions.
• Organic chemistry: conformations of molecules, stacking of aromatic rings, intramolecular hydrogen bonds stabilizing particular shapes.
• Biochemistry: stabilization of protein and nucleic acid structures, enzyme catalysis, membrane organization.
• Materials science: mechanical properties, glass transition in polymers, behavior of liquid crystals, adsorption on surfaces.
• Analytical chemistry: separation principles in chromatography (different interactions with stationary and mobile phases), binding in sensing and recognition systems.

Understanding special intermolecular interactions thus provides an essential bridge between microscopic structure and macroscopic behavior in a wide variety of chemical systems.