11.3:
Intermolecular Forces
Intermolecular forces, which exist between molecules, originate from electrostatic interactions between charges, partial charges, and temporary charges.
All molecules produce temporary charges. Due to the varying distribution of electrons, a higher electron density in one region of the electron cloud results in an instantaneous dipole, or a temporary dipole.
This subsequently induces another instantaneous dipole in the neighboring molecule. The domino effect of dipoles gives rise to weak intermolecular attractive forces called dispersion forces, which exist between all molecules, whether polar or nonpolar.
Some covalent compounds, like water, for example, exhibit electron-rich and electron-poor regions caused by the atoms’ differences in electronegativity.
The uneven distribution of shared electrons and the molecular shape of the compound create permanent partial charges, resulting in a permanent dipole in the otherwise neutral compound — making it polar.
Molecules with permanent dipoles — also called polar compounds — align themselves through dipole–dipole forces, where the positive end of one molecule interacts electrostatically with the negative end of the neighboring molecule.
If a polar compound contains a hydrogen atom covalently bonded to a small, highly electronegative atom like fluorine, oxygen, or nitrogen, the atoms tend to exhibit larger partial charges.
Consequently, the hydrogen atom in F–H, O–H, or N–H bonds strongly interacts with the electronegative atom on its neighbor through a special type of dipole–dipole force called a hydrogen bond.
Notably, hydrogen bonds are stronger than dipole–dipole forces, and compounds capable of forming hydrogen bonds exhibit higher melting and boiling points.
The three intermolecular forces — dispersion, dipole–dipole, and hydrogen bonds — are, in comparison with intramolecular forces, relatively weak, with varying strengths. They collectively are classified as van der Waals forces.
While dispersion forces are present between all molecules, polar or nonpolar, dipole–dipole forces and hydrogen bonds exist only around polar molecules.
Exclusive to solutions is the ion–dipole force, which is the strongest intermolecular force. When an ionic compound like sodium chloride is dissolved in a polar solvent like water, the dissociated ions interact with the solvent's dipoles via strong ion–dipole forces.
Here, the cations associate with the negative ends of the water molecules, while the anions interact with the positive ends.
11.3:
Intermolecular Forces
Atoms and molecules interact through bonds (or forces): intramolecular and intermolecular. The forces are electrostatic as they arise from interactions (attractive or repulsive) between charged species (permanent, partial, or temporary charges) and exist with varying strengths between ions, polar, nonpolar, and neutral molecules. The different types of intermolecular forces are ion–dipole, dipole–dipole, hydrogen bonds, and dispersion; among these, dipole–dipole, hydrogen bonds, and dispersion forces exist between neutral atoms and molecules and are collectively known as van der Waals forces.
Dipole–dipole forces
Polar molecules have a partial positive charge on one end and a partial negative charge on the other end of the molecule—a separation of charge called a dipole. In a polar molecule like HCl, the more electronegative Cl atom bears the partial negative charge, whereas the less electronegative H atom bears the partial positive charge. An attractive force between HCl molecules results from the attraction between the positive end of one HCl molecule and the negative end of another. This attractive force is called a dipole–dipole attraction—the electrostatic force between the partially positive end of one polar molecule and the partially negative end of another.
Dispersion Forces
One of the three van der Waals forces is present in all condensed phases, regardless of the nature of the atoms or molecules composing the substance. This attractive force is called the London dispersion force in honor of German-born American physicist Fritz London who, in 1928, first explained it. This force is often referred to as simply the dispersion force. Because the electrons of an atom or molecule are in constant motion (or, alternatively, their location is subject to quantum-mechanical variability), at any moment in time, an atom or molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically. The presence of this dipole can, in turn, distort the electrons of a neighboring atom or molecule, producing an induced dipole. These two rapidly fluctuating, temporary dipoles thus result in a relatively weak electrostatic attraction between the species—a so-called dispersion force.
Dispersion forces that develop between atoms in different molecules can attract the two molecules to each other. The forces are relatively weak, however, and become significant only when the molecules are very close. Larger and heavier atoms and molecules exhibit stronger dispersion forces than do smaller and lighter atoms and molecules. F2 and Cl2 are gases at room temperature (reflecting weaker attractive forces); Br2 is a liquid, and I2 is a solid (reflecting stronger attractive forces).
Hydrogen Bonding
Nitrosyl fluoride (ONF, molecular mass 49 amu) is a gas at room temperature, whereas water (H2O, molecular mass 18 amu) is a liquid, even though it has a lower molecular mass. Both molecules have about the same shape, and ONF is the heavier and larger molecule. It is, therefore, expected to experience more significant dispersion forces. Additionally, this difference in boiling points cannot be the result of differences in the dipole moments of the molecules. Both molecules are polar and exhibit comparable dipole moments. The large difference between the boiling points is due to a particularly strong dipole–dipole attraction that may occur when a molecule contains a hydrogen atom bonded to a fluorine, oxygen, or nitrogen atom (the three most electronegative elements). The very large difference in electronegativity between the H atom (2.1) and the atom to which it is bonded (4.0 for a F atom, 3.5 for an O atom, or 3.0 for a N atom), combined with the very small size of a H atom and the relatively small sizes of F, O, or N atoms, leads to highly concentrated partial charges with these atoms. Molecules with F-H, O-H, or N-H moieties are very strongly attracted to similar moieties in nearby molecules, a particularly strong type of dipole–dipole attraction called hydrogen bonding. Examples of hydrogen bonds include HF⋯HF, H2O⋯HOH, and H3N⋯HNH2, in which the hydrogen bonds are denoted by dots.
Ion–dipole forces
An ion–dipole force is the electrostatic attraction between an ion and a dipole. These forces are common in solutions and play an important role in the dissolution of ionic compounds in water.
When an ionic compound like KCl is added to a polar solvent like water, the ions in the solid separate and disperse uniformly. Ion–dipole forces attract the positive (hydrogen) end of the polar water molecules to the negative chloride ions at the surface of the solid, and they attract the negative (oxygen) ends to the positive potassium ions. The water molecules surround individual K+ and Cl− ions, reducing the strong interionic forces that bind the ions together (in a solid) and letting them move off into solution as solvated ions. Overcoming the electrostatic attraction permits the independent motion of each hydrated ion in a dilute solution as the ions transition from fixed positions in the undissolved compound to widely dispersed, solvated ions in solution.
The strength of ion–dipole interactions is directly proportional to i) the charge on the ion and ii) the magnitude of the dipole of polar molecules.
This text is adapted from Openstax, Chemistry 2e, Chapter 10: Liquids and Solids.